NBS SPECIAL PUBLICATION 450

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Blood pH, Gases, and Electrolytes

Proceedings of the workshop on

pH and Blood Gases, held at the

National Bureau of Standards,

Gaithersburg, Maryland,

July 7-8, 1975

Richard A. Durst, Editor

Analytical Chemistry Division

Institute for Materials Research

National Bureau of Standards

Washington, D.C. 20234

Sponsored by:

American Association for Clinical Chemistry

American Society of Clinical Pathologists

International Federation of Clinical Chemistry

National Committee for Clinical Laboratory Standards

U.S. Department of Commerce, National Bureau of Standards

U.S. DEPARTMENT OF COMMERCE, Juanita M. Kreps, Secretary

Dr. Sidney Harman, Under Secretary

Jordan J. Baruch, Assistant Secretary for Science and Technology

issued June 1977
Library of Congress Cataloging in Publication Data

Workshop on pH and Blood Gases, National Bureau of Standards,

1975. Blood pH, gases and electrolytes. (NBS special publication; 450)

Supt. of Docs, no.: €13.10:450

1. Blood—Analysis and chemistry—Congresses. 2. Blood gases

Analysis—Congresses. 3. Hydrogen-ion concentration—Measure- ment—Congresses. 4. Electrolytes—Analysis—Congresses. I. Durst, Richard A., 1937-11. American Association of Clinical Chemists.

III. Title. IV. Series: United States. National Bureau of Standards.

Special publication; 450. QC100.U57 no. 450 [RB45] 602'.ls [616.07'561] 76-608179

National Bureau of Standards Special Publication 450

Nat. Bur. Stand. (U.S.), Spec. Publ. 450, 338 pages (June 1977) CODEN: XNBSAV

U.S. GOVERNMENT PRINTING OFFICE

WASHINGTON: 1976

For sale by the Superintendent of Documents, U.S. Government Printing OflSce, Washington, D.C. 20402 (Order by SD Catalog No. C13.10 :450) . Stock No. 003-003-01792^2 Price $4.75

(Add 25 percent additional for other than U.S. mailing)

.
FOREWORD

The Analytical Chemistry Division of the National Bureau of Standards (NBS), Institute

for Materials Research has as its goal the development of new techniques of analysis,

standard materials and instruments to improve the accuracy and precision with which analyses

are performed. To obtain as broad input as possible from the user community on the needs

and trends in analytical chemistry and to disseminate information outside NBS, groups of

experts are invited to attend workshops at the Bureau. These workshops often are cosponsored

by other scientific and professional organizations. These proceedings include the presentations and deliberations of one of a series of Analytical Chemistry Division sponsored

workshops and are concerned with the accurate measurement of blood pH and gases. Other

workshops in this series have included Secondary Ion Mass Spectrometry, Aerosol Measurements, Marine Pollution Monitoring (Petroleum), Monte Carlo Calculations in Ion Probe

Microanalysis and Scanning Electron Microscopy, Standards for Spectrophotometry and Lu- minescence, Standards Required in Marine Science, Standards Required for Offshore Oil

Drilling, and Standards Required for Environmental Monitoring of Oil Shale Processing.

The international character of the attendance at this workshop and of the interest

exhibited by the cosponsoring organizations: the American Association of Clinical Chemists,

the American Society of Clinical Pathologists, the International Federation of Clinical

Chemistry, and the National Committee for Clinical Laboratory Standards, are indicative of

the interest and importance of this subject throughout the clinical community. This publication should serve to characterize the current state of accuracy for blood pH and gases

measurements and to point out the standards required in this field.

P. D. LaFleur

Chief

Analytical Chemistry Division

iii
PREFACE

This publication is the formal report of the proceedings of the Workshop on pH and

Blood Gases sponsored by the American Association of Clinical Chemists, American Society of

Clinical Pathologists, International Federation of Clinical Chemistry, National Committee

for Clinical Laboratory Standards, and the National Bureau of Standards. The purpose of

this meeting was to discuss the status and needs in this very important area of clinical

measurement and to provide a starting point for future cooperative efforts on an international level toward the standardization of pH and blood gas measurements and the various

quantities and terms used in this field.

The importance of the acid-base status for clinical diagnosis is reflected in the fact

that the determination of pH and blood gases is one of the most widely and frequently

performed clinical tests. However, even though it has been more than a half century since

Henderson and Van Slyke described the chemical and physiological relationships and mechanisms

involved in acid-base balance, there still remain many uncertainties as to which indices are best measured, what parameters should be derived from the measured data and how to calculate

them, and finally how these data should be interpreted by the clinicians.

The first requirement is an internationally agreed-upon set of definitions, terms, and

symbols. The conflict of ideas that still exists is serious, not only because of the need

for uniformity in nomenclature, but also for the more pragmatic reason that, as highly

sophisticated blood pH/gas analyzers are developed, the derived parameters will be critically

dependent upon the assumptions and "standard values" used in the calculations. It is hoped

that this workshop and the resulting proceedings will provide a preliminary step toward

achieving international accord in this area. Of course, a two-day workshop allowed little

time for problem solving or the resolution of details. However, it did permit the exposure

of diverse viewpoints on the critical needs in this field so that subsequent efforts, such

as those of the newly established IFCC Expert Panel on pH and Blood Gases, can be directed

toward the most important topics.

Editorial changes have been made to achieve some measure of internal consistency,

especially with regard to the International System of Units (SI). However no changes are made in the nomenclature used by the authors since one of the purposes of this workshop was

to reveal the diversity of symbols, terms, definitions, etc., in order to arrive at some future uniformity in this subject. Also, certain fundamental discussions appear in more

than one paper but, again, in view of the diversity of approaches to this subject, this

duplication was intentionally preserved. Although all the papers have been reviewed by the

editor, the views expressed are entirely the responsibility of the individual authors.

In order to specify the procedures adequately, it has been necessary to identify

commercial materials and equipment in this report. In no case does such identification

imply recommendation or endorsement by the National Bureau of Standards, nor does it imply

that the material or equipment is necessarily the best available for the purpose.

The organization of the workshop and publication of these proceedings would not have

been possible without the cooperation and assistance of a large group of people within the

National Bureau of Standards. Particular thanks are given to Philip D. LaFleur, Chief of

the Analytical Chemistry Division, for his encouragement and support in this effort, and to Ronald B. Johnson and Robert F. Martin of the Institute for Materials Research for their

assistance in the business matters associated with this project. I also want to express my

appreciation to Ellen N. Ring and her staff in the IMR Text-Editing Facility. The NBS

Office of Information Activities, with special help from Sara R. Torrence, and also Rebecca

J. Morehouse and Miriam K. Gland of the Office of Technical Publications gave invaluable

assistance in many phases of the effort, varying from the initial meeting arrangements to the final publication of these proceedings. Special thanks are given to Gloria Burdick for

her help in preparing the original workshop correspondence and program, Rosemary Hormuth for

transcription of the recorded discussions, and Carolyn A. Shipley for her continuous effort

in typing the coded manuscripts.

Richard A. Durst

iv
ABSTRACT

On July 7-8, 1975, a workshop was held at the National Bureau of Standards to discuss

the status and needs of this very important area of clinical measurement. A major goal of

this workshop was the initiation of cooperative efforts on an international level toward the

standardization of pH and blood gas measurements and the various quantities and terms used

in this field.

To this end, the first technical session was concerned with the acid-base status of

blood and included the topics: Definitions of Quantities and Concepts; Recommendations of

Nomenclature, Physiological Terminology and Symbols; Reference Values; and the Evaluation of

Nomograms and Algorithms. The second session addressed itself to the more practical aspects

of this subject and included the topics: Blood Sampling, Handling, and Storage; Instrument

Specifications; Quality Control and Standards; and the Development of Reference Methods.

Finally, a brief session was held on the newer topic of the electrometric measurement of

blood electrolytes.

This volume contains all of the papers invited for presentation at the workshop by some of the leading clinical and medical authorities on this subject and also includes a transcription of the extensive discussion sessions.

Key words: acid-base status; blood electrolytes; blood gases; blood pH; calcium; carbon

dioxide; hydrogen ion concentration; nomograms; oxygen; P-^ ; pH; P^ ; potassium; sodium.

V

CONTENTS

PAGE

FOREWORD iii

PREFACE IV

Part I. Acid-Base Status

DEFINITIONS OF ACID-BASE QUANTITIES: TERMINOLOGY, SYMBOLS, AND SI UNITS

01 e Siggaard-Andersen 1

THE BUFFER VALUE OF PLASMA, ERYTHROCYTE FLUID AND WHOLE BLOOD

Ole Siggaard-Andersen, M. R0rth, and D. A. P. Strickland 11

ACID-BASE ALGORITHMS

Ole Siggaard-Andersen 21

DETERMINATION OF TOTAL CO2 CONCENTRATION IN BLOOD OR PLASMA

P. Rispens, E. J. van Kampen, and W. G. Zijlstra 27

THE APPARENT OVERALL FIRST DISSOCIATION CONSTANT OF CO2 IN PLASMA

P. Rispens and W. G. Zijlstra 33

QUANTITATIVE RELATIONSHIPS BETWEEN TOTAL CO2 CONCENTRATION IN BLOOD AND PLASMA,

PLASMA BICARBONATE CONCENTRATION, PLASMA pH AND CARBON DIOXIDE TENSION

BETWEEN 16-42 °C

P. Rispens, J. P. Zock, and W. G. Zijlstra 39

P^Q INDEPENDENT QUANTITIES, WHY OR WHY" NOT

P. Rispens, W. G. Zijlstra, and E. J. van Kampen 47

BASE EXCESS. WHY REOPEN THE ACID-BASE DEBATE?

P. J. N. Howorth 53

USE OF IN VIVO CO2 TITRATION CURVES IN THE PHYSIOLOGICAL ASSESSMENT

OF ACID-BASE BALANCE

P. J. N. Howorth 57

RELATIVE ACTIVITY AND SI UNITS

B. F. Visser and A. H. J. Maas 69

SEMI-EMPIRICAL ACID-BASE PROGRAM

B. F. Visser, A. J. Hoelen, J. A. Kreuger, and A. H. J. Maas 73

"ACID-BASE SEMANTICS"--A CENTURY OF THE TOWER OF BABEL

Harry F, Weisberg 75

vii
PAGE

TRI-SLIDETM CALCULATOR FOR HENDERSON-HASSELBALCH EQUATION AND C02RREC°t-02-SLIDETM

FOR TEMPERATURE CORRECTIONS OF pH, Pco^, AND P02

Harry F. Weisberg 91

AIDS FOR EVALUATION OF ACID-BASE IMBALANCE— DIAGRAMS, NOMOGRAMS,

AND SLIDE-RULES

Harry F. Weisberg 103

THE OVERALL FIRST IONIZATION EQUATION OF CARBONIC ACID AS RELATED

TO CO2 IN GAS PHASE: A NEW pK

A. H. J. Maas and B. F. Visser 119

TOWARDS A PHYSIOLOGIC NOMENCLATURE FOR IN VIVO DISTURBANCES

OF ACID-BASE BALANCE

Jordan J. Cohen 127

MINIMAL ACCEPTANCE CRITERIA FOR ACID-BASE NOMOGRAMS

Jordan J. Cohen 131

A PHYSIOLOGICAL APPROACH TO ACID-BASE DIAGNOSTICS

Poul Kildeberg and Knud Engel 133

CO2 SOLUBILITY, pK' AND RELATED FACTORS IN ACID-BASE BALANCE

William H. Austin 143

DEFINITION OF OXYGEN SATURATION AND CHARACTERIZATION ' OF OXYGEN-HEMOGLOBIN AFFINITY

Arthur L. Mai enfant 153

PROBLEMS ASSOCIATED WITH THE DEFINITION OF MEASURED AND CALCULATED

QUANTITIES IN BLOOD pH AND GAS ANALYSIS

t Robert W. Burnett 163

Part II. Instrumentation, Methodology, and Standards

ARTERIAL AND VENOUS BLOOD SAMPLES IN ACID-BASE BALANCE

William H. Austin . 167

BLOOD SAMPLING AND HANDLING IN THE DETERMINATION OF BLOOD pH

AND BLOOD GASES

Arthur H. Richards 171

NON-ANALYTICAL SOURCES OF LABORATORY ERROR IN pH AND BLOOD GAS ANALYSIS

Jack H. Ladenson 175

vi i i
PAGE

INSTRUMENT SPECIFICATIONS

S. Raymond Gambino 191

EFFECTS OF THE LIQUID JUNCTION ON pH MEASUREMENT IN BLOOD;

THE 0.160 MOL/L SODIUM CHLORIDE BRIDGE

A. H. J. Maas 195

THE CRITICAL CARE LABORATORY: A 10- YEAR PERSPECTIVE

Myron B. Laver and Domenic R. Misiano 201

A THEORETICAL AND PRACTICAL ANALYSIS OF P^ MICROELECTRODE BEHAVIOR:

THE THREE-SHELL MODEL

R. G. Buckles, H. Heitmann, and M. B. Laver 207

DYNAMIC RESPONSE OF A pCOz ELECTRODE

S. J. Pace and M. J. D. Brand 227

MONITORING OF OXYGEN PRESSURE IN HUMAN AND ANIMAL BLOOD

H. P. Kimmich and F. Kreuzer 235

ALTERNATIVE METHODS OF CO2 MEASUREMENT, WITH PARTICULAR REFERENCE

TO CONTINUOUS RECORDING

L. H. J. van Kempen and F. Kreuzer , 239

NBS STANDARDS FOR pH AND ION ACTIVITY MEASUREMENTS IN BIOLOGICAL FLUIDS

Richard A. Durst and Roger G. Bates 247

USE OF CARBON DIOXIDE- AND OXYGEN TONOMETERED PHOSPHATE-BICARBONATECHLORIDE-GLYCEROL-WATER MIXTURES FOR CALIBRATION AND CONTROL OF pH,

PCO2, AND PO2 ELECTRODE SYSTEMS

A. H. J. Maas, A, H. Veefkipd, R. A. M. Van den Camp, A. J. Teunissen,

A. B. T. J. Boink, and T. J. C. Ruigrok 257

CALIBRATION OF BLOOD GAS ANALYZERS

Alan H. Runck 267

QUALITY CONTROL

S. Raymond Gambino 273

QUALITY CONTROL AND STANDARDS

Daniel C. Noonan 275

DEVELOPMENT OF REFERENCE METHODS: BLOOD GAS ANALYSIS

Arthur L. Mai enfant and Kevin D. Fall an 279

QUALITY CONTROL AND STANDARDS

SfJren K. Sfirensen 285

ix
DEVELOPMENT OF REFERENCE METHODS

Sfiren K. S^rensen

PAGE

289

Part III. Electrolytes

STANDARDIZATION OF ION-SELECTIVE ELECTRODES FOR SERUM ANALYSIS

M. S. Mohan and Roger G. Bates , 293

ELECTROLYTE ACTIVITIES IN HUMAN BLOOD PLASMA

M. J. D. Brand and W. J. Scott 301

THE KING'S COLLEGE HOSPITAL ION-SELECTIVE ELECTRODE

SERUM ELECTROLYTE ANALYZER

A. D. Hirst, P. Gay, P. Richardson, and P. J. N. Howorth 311

Part IV. Discussion

WORKSHOP DISCUSSION 315

X
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DEFINITIONS OF ACID-BASE QUANTITIES: TERMINOLOGY, SYMBOLS,

AND SI UNITS

Ole Si ggaard-Andersen

Department of Clinical Chemistry

University of Copenhagen

Copenhagen County Hospital

Herlev, Denmark

In clinical medicine, patients are often encountered with disturbances characterized by

accumulation or loss of hydrogen ions and/or carbon dioxide. For routine clinical purposes,

such disturbances are evaluated by measuring various quantities, for example, the "blood pH"

,

the "hydrogen ion concentration", the "eucapnic pH", the "PCO2", the "total -C02"-, the "plasma

bicarbonate", the "CO2 combining power", the "CO2 capacity", the "standard bicarbonate", the

"buffer base", and the "base excess." Most of these designations are "clinical slang" and

they ought to be replaced by the more systematical chemical terminology which has been re- commended by various international organizations [1-5]^.

1. Names and Symbols for Chemical Quantities

The name of a physical or chemical quantity must include a specification of the kind of

quantity and a sufficiently detailed characterization of the physical system to which the

quantity refers. This often involves a specification of one or more components of the

system or processes or reactions in the system.

Table 1 shows the names of some of the acid-base quantities according to the international recommendations [2], The names follow the scheme: System-component, kind of

quantity. For example: P-hydrogen ion, substance concentration equals 53 nmol/1. In symbols this may be written [6]:

eH'^(P) = {fiH^(P)} • [cH^(P)] = 53 nmol/1 . (1)

The name as well as the symbol should (explicitly or implicitly) include the following

specifications:

(1) Kind of quantity . The different kinds of quantities are defined by the International Organization for Standardization in a publication series: ISO 31 [3]. The total

number of different kinds of quantities amounts to several hundred. In clinical chemistry,

the kinds of quantities listed in table 2 are of special interest. The symbol for the

kind of quantity should always be a single letter and should always be printed in italics

(sloping type), while all other symbols are printed in Roman type (upright) [1], An exception to this rule is pH which unfortunately has been defined as a special kind of quantity (see later), but nevertheless is written in upright type.

According to the recommendations of the International Federation of Clinical Chemistry

(IFCC) and the International Union of Pure and Applied Chemistry (lUPAC) [3,4], substance

oonoentration (or simply concentration) should be employed rather than mass concentration in

all cases where a formula unit can be defined for the component. For example, the hemoglobin concentration should be the substance concentration based on a molar mass, of

16114 g/mol referring to one quarter of a hemoglobin molecule with two a and two b chains.

Figures in brackets indicate literature references at the end of this paper.

1
Table 1. Names of various acid-base quantities according to the recommendations

of the International Federation of Clinical Chemistry (IFCC) and the International Union of Pure and Applied Chemistry (lUPAC) [2]. (The symbols are author's suggestions.)

Name Symbol

U - Acid (H), substance concentration (method)

aB - Base (H"'^-binding groups), substance concentration

(Singer and Hastings, 1948)

aB - Base (H^-binding groups), substance concentration

difference (method; Pt - Norm)

Gas (aB equil.) - Carbon dioxide, partial pressure

(method; 37.0 °C)

(aB) P - Carbon dioxide, substance concentration

P - Carbonate + carbon dioxide, substance con- centration

P - Hydrogen carbonate ion, substance concentration

(blood; aiOz) = 0.21 mmol/1; e(C02) = l-IQ mmol/1;

0 = 37 °C)

(aB)P - Hydrogen ion, substance concentration

aB - Plasma, pH (37 °C)

Pt - Urine, pH

ActH^(U)

eBB' (aB)

AeB' (aB)

pC02(aB)

cC02(Pc:aB)

ctC02(P)

cHC03(PcstB)

cH"^(PcaB)

paH"^(PcB)

paH'^(U)

Table 2. Various kinds of quantities generally employed in clinical chemistry.

Name Symbol Unit Definition (C = component, S = system)

Number N 1 (one)

Amount of substance n mol

number of particles or molecules

number of formula units:

6.023 • 1023 • nC/mol = BZ

Mass m kg

Volume V 1 (or m^)

Number fraction

Substance fraction

Mass fraction

X

X

w

1 Cone)

1 Cone)

1 Cone)

^Cj(S)

a;C.(S)

wC(S) =

= /vc.

= nC • J mC(S

;S)/EffC.(S)

i [S)/EnC.(S)

i )/mS

Volume fraction 1 Cone) <|)C(S) = 7C(S )/7S

Number concentration C

1-1 CC(S) = A/C(S' m

Substance concentration c mol/1 cC(S) = nC(S^ /ys

Mass concentration P kg/1 pC(S) - mC(S' /ys

Molality m mol /kg mC(S) = nc(s; /mSolvent(S)

(Partial) pressure P Pa pC(S) = xc(s;

2
(2) Component . The names of the components are specified in considerable detail according to a certain set of rules [2]. Several of the names may be expressed in alternative

forms, e.g., carbonate + carbon dioxide = total CO2. The symbols for the inorganic substances should be the familiar chemical symbols. For certain organic substances, symbols

are also almost international, e.g., hemoglobin = Hb, protein = Pr.

(3) System . The above mentioned recommendations include a list of symbols for the

systems, e.g. blood = B, arterial blood = aB, plasma = P, urine = U, etc.

(4) Number . . Braces {} around the symbol for the quantity indicate the pure number

value of the quantity [1].

(5) Unit . Brackets [] around the symbol for the quantity indicate the unit of the

quantity [1]. Whenever possible, it is recommended to use the SI units (m, kg, s, -system).

This means that the traditional unit for pressure, mmHg, should be replaced by Pa (= N/m^),

conversion factor: 1 mmHg = 0.133 kPa.

The names recommended by IFCC and lUPAC are especially designed for tables and laboratory records [2]. In a running text or in the spoken language, a different order is often

chosen, e.g., the concentration of hydrogen ions in the plasma, the concentration of plasma

hydrogen ion, the plasma hydrogen ion concentration, the plasma concentration of hydrogen

ions, the hydrogen ion concentration in the plasma. All combinations are seen in the

1 iterature.

2. Stock Components, Amount of Substance, and Chemical Potential

The fundamental basis for a description of acid-base physiology was the development of

physical chemistry and thermodynamics which led to an exact description of a chemical system

at the beginning of this century.

The organism is an open chemical system in a dynamic equilibrium with its surroundings.

A complete description of the system therefore requires not only a description of the momentary composition of the different body phases, including the aoid-base status of the blood,

but also requires a description of the rate of intake, production, conversion, and excretion

of acids and bases, i.e. , a description of the aoid-base balance of the organism.

A chemical system, e.g., blood plasma, can theoretically be prepared from a limited num- ber of stock components, for example, H2O, NasPOi^, HCl , CO2, O2, NaCl , etc. It is not the

purpose here to describe all the stock components of plasma. We are in the present connection especially interested in changes in the amount of hydrogen ions in the system.

Adding hydrogen ions to the system presents a special problem because they cannot be

added or removed alone, according to the law of electrical neutrality. With the present

choice of stock components, a hydrogen ion is always accompanied by a chloride ion. Alternatively, hydrogen ions couia be added or removed in exchange for sodium ions which virtually

means removing or adding NaOH. The first method is described as addition or removal of

strong acid, the second method as removal or addition of strong base, and the two methods

are considered equivalent, although the effects on the ionic strength will be slightly

different.

The stock components can react mutually giving rise to several derived components,

e.g.

,

CO2 + H2O ^ H2CO3 ^ H"^ + HCO3 • (2)

with the above-mentioned choice of stock components, HCO3 is considered a derived component.

However, the choice of stock components is arbitrary; one could also choose H2O, NaHCOs, and

HCl as stock components, in which case CO2 would be a derived component.

3
The first mentioned choice of stock components is based on physiological considerations,

according to which H2O, CO2, and strong acid or base are ingested or formed in the organism

and excreted in the lungs and kidneys.

The important equation which expresses the change in free energy of the system (dc) when

the amounts of stock components are changed (at constant temperature and pressure) is the

following:

dG = yH"*" • dntH"^

+ yC02 • dntC02

+ yCl" • dntCl"

+ yNa^ • dntNa"*"

+ y02 • dnt02 J

t = prefix designating total amount added, n = amount of substance (unit: mole),

y = electrochemical potential (unit: joule/mole),

= y + 2 • F • (j).

where y = chemical potential, z = charge number, F = Faraday constant, and ^

= inner electrical potential.

This important equation forms the basis for derivation of important relationships among

the various quantities. In the present connection, the equation only serves to define the

ahemieal potential as the partial molar free energy of the different components, and the

equation serves to emphasize that each stock component of the system can be characterized by

two quantities: (1) an intensive quantity, i.e., the chemical potential, and (2) an extensive

quantity, i.e., the amount of substance of added component.

The component of primary interest is (strong acid or base), but it is necessary to

deal with another component which greatly influences the chemical potential of H"*" in an aqueous solution, namely 002- In the following, the intensive and extensive quantities are described in more detail for these two components.

3. Chemical Potential of Hydrogen Ions, AyH^ or pH

From a physico-chemical point of view, the chemical potential of the hydrogen ions is the most important quantity for description of the "acidity" of a solution. The significance

of the chemical potential is that hydrogen ions tend to diffuse from a phase with a higher

chemical potential to a phase with a lower potential, unless an electrical potential difference exists between the phases. In that case, the equilibrium condition requires that

the electrochemical potential [i.e. , the sum of the chemical potential and the electrical

potential times s times the Faraday constant) should be identical in the two phases.

It must be emphasized that it is impossible (theoretically) to measure absolute chemical potentials. Using two ideal H'*' electrodes: E|^lS|^|KCl |Sp|Ej^, we can measure yH'''(Sp) -

yH"*'(S|^) + FEA£'(LJ). It is possible to minimize ea£'(LJ) by means of a saturated KCl bridge.

If S. is a reference solution (S®), we can therefore measure (approximately):

AyH^(sQ/SJ = Ao^H'^(Sr) = yH'^(SR) mH^S^) (4)
i.e.:, the excess chemical potential of hydrogen ions in solution Sp. The reference solution

is chosen so that aH'*'(S®) = 1 (see below).

Often the activity of the hydrogen ions is used as a measure of the acidity. The activity is related to the chemical potential as follows:

aV(S) = exp^i^ (5)

ra*H"'(S®/S) = r.a*H^(S) = ^^^^ e aH^(S) (6)

where a* is absolute activity and a is relative activity; r is a convenient mathematical

symbol for ratio (compare a: Inra = Alnx). The reference solution (S®) is defined by tl following equations which also define the activity coefficient y:

^H+(s) = ltL(Sl^_|L(Sl. ^ o(S) - 1:^yH^(S) - 1 . (7)

mH (S^)

S® may be said to be a hypothetical ideal solution with mW^{s'^) = 1 mol/kg and yH^(S®) = 1. Actually, aH"*" is about 1 in a hydrochloric acid solution with mW^ about 0.85 mol/kg, i.e.,

yW^ = 1.18. yH"*" varies with the ionic strength. For human plasma, yH''"(P) is often assumed

to be about 0.8.

Calling the mathematical operator -log for p, gives the familiar pH quantity:

-logaH"^ = paH* = pH . (8)

It will be seen that changes in pH (with opposite sign) are directly proportional to changes

in the chemical potential of the hydrogen ions.

A number of international standard buffers have been defined in order to enable re- producible pH measurements by means of electrodes, and the pH concept has thereby been given

an international "operational definition." It was unfortunate, however, that lUPAC defined

pH as a special derived kind of quantity [3].

Occasionally the hydrogen ion aonoentTation [oW^) is used as a measure of the acidity.

The relationship between cH+ and ihW^ {i.e., molality) is the following:

cH'^(S) = mH'^(S) • pH20(S) , (9)

where pH20(S) is the mass concentration of H2O in the system [e.g., for plasma: pH20(P) = 0.94 kg/1). In the medical literature, opinion is divided as to whether pH or eH+ should be

preferred as a measure of the acidity [6]. For "medical purposes," the cH+ of the plasma is

generally calculated without distinguishing between molality and concentration and without

5
taking the activity coefficient into account: cH'*' = anti1og(9 - pH) nmo1/l. The correct

equation would be:

. pH20(P)/(kg/1)

cH (P) = T • antilog(9 - pH) nmol/1 (10)

yH^P)

Now that SI units are strongly recommended, it seems logical to use the SI unit joule

also for the "pH quantity", i.e., to use excess chemical potential of hydrogen ions. For human arterial blood plasma, the reference range would be:

AyH^: -43.76 to -44.23 kJ/mol , and the extreme pathological range:

AyH"^: -40.37 to -46.31 kJ/mol , (corresponding to pH 7.37 to 7.45, and 6.80 to 7.80, respectively).

For conversion of pH values to AyH^ and vice versa, the following equation applies:

AvH"^ = -i?? • In 10 • pH , (n)

or for T = 310.15 K:

AyH^ = -5.937 . pH kJ/mol . (12)

4. Titratable Acid or Base; the Base Excess Concept, AetH^ or AcB'

The amount of added hydrogen ions in the system (ntH^) , i.e., added strong acid, is an arbitrary quantity which requires a definition of the system before the addition. If the

initial state of the system is defined, then the amount of added strong acid or base can be determined by back titration with strong base, respectively, strong acid. In other words,

one can define and determine the difference (AntH''"): the amount of total hydrogen ion in the system before titration minus the amount of total hydrogen ion in the system after titration. By dividing with the volume of the system, this extensive quantity can be expressed

as an "excess concentration" of total hydrogen ion:

AetH^(S) = AMtH"^(S)/yS . (13)

In the usual terminology, this quantity with opposite sign is designated the excess concen- tration of base, or simply the "base excess":

ActH'^(S) = -AeB'(S) . (14)

The base excess concentration can also be expressed as being the concentration of

titratable base minus the concentration of titratable acid. For plasma, the endpoint of titration is defined at pH = 7.40, pCOz = 40 mmHg (= 5.33 kPa) , and t = 37 °C. For whole

blood, the endpoint is the same, the pH value referring to the plasma phase. For urine, the

6
endpoint may be defined at pH = 7.40, pC02 = 0, pNHs = 0, and t = 37 °C; by this choice HCO3

is included as titratable base, while NH^ is included as titratable acid.

For urine, the titratable acid or base is generally determined directly by titration.

For blood or plasma, however, the quantity is generally determined indirectly by calculation

from the directly measured pH and pC02 values [6].

It appears that the concentration of total hydrogen ion in the plasma is witho t practical significance, contrary to, for example, the concentration of total calcium ion 1. the

plasma. The reason is that the solvent itself (H2O) contains large amounts of hydrogen luns

in bound form. Therefore, it is only relevant to determine changes or differences in the

concentration of strong acid or base in the solution. Apart from this, however, there is no principal difference between the description of H"*" and Ca''""'': both ions can be added to or removed from the system (together with an indifferent anion) and both are bound specifically

by binding groups, which in the case the hydrogen ions, are called base groups.

It is only in recent years that the base excess concentration has been used routinely

for description of the acid-base status of the blood. Originally, the bicarbonate concentration or the total CO2 concentration was used as indicators of the accumulation of non- carbonic acid or base, i.e., as measures of a metabolic acid-base disturbance. When dealing

with a pure bicarbonate solution, the change in the bicarbonate concentration equals the

change in the concentration of titratable acid or base, and the bicarbonate concentration

is unaffected by changes in PCO2; but in the presence of non-bicarbonate buffers, the bicarbonate concentration also changes during isolated changes in the pC02 even when the concen- tration of titratable acid or base is constant. In order to compensate for this effect, the

PCO2 was standardized to a normal value of 40 mmHg (= 5.33 kPa), and the concentration of

total CO2 or HCO3 of such standardized blood or plasma was designated "the plasma CO2 com- bining power," "the CO2 capacity of the blood," or "the standard bicarbonate." However,

none of these quantities accurately reflect the accumulation of noncarbonic acid or base

in the blood or plasma.

5. The Chemical Potential or the Partial Pressure of CO2,

AyC02, or PCO2

Generally, the partial pressure of carbon dioxide in a gas phase in equilibrium with

the blood is used as a measure of the chemical potential of CO2. Conversion of pC02 to yC02 or vice versa follows the equation

AVCO2 = ifT.lnCpCOj/lOl .3 kPa) , (15)

or at T = 310.15 K:

AyCOz =^5.937-log(pC02/101.3 kPa) kJ/mol . (16)

To be exact, pC02 should be multiplied by a fugacity coefficient which can be taken to be 1 for pressures below 100 kPa. The reference system is a system in equilibrium with a hypothetical ideal gas phase with pC02 = 101.3 kPa.

For chemical substances, lUPAC has recommended the use of molecular quantities, i.e.,

substance concentration rather than mass concentration. This recommendation might be extended to a recommendation of using the excess chemical potential rather than other quantities

(activity, partial pressure, concentration of free component, etc.) which are presently

employed as substitutes for the chemical potential.

For human arterial blood, the reference range for the excess chemical potential of CO2

is: AyC02: -8.1 to -7.4 kJ/mol , and the extreme pathological range is about: AyC02; -11.2

to -5.2 kJ/mol, (corresponding to pC02 4.3 to 5.7 kPa and 1.3 to 13.3 kPa, respectively).

7
The partial pressure of CO2 is related to the concentration of dissolved CO2 according

to Henry *"s law:

0CO2 = aC02 • PCO2 , (17)

where aC02 is the solubility coefficient. Generally, aCOz includes the very low concentration

of H2CO3. The solubility coefficient for normal plasma is 0.306 mmol • 1"^ • mmHg"^ at 37 °C. (or in SI units: 0.231 mmol/J) with a normal biological standard deviation of about 0.0004

mmol • 1"^ • mmHg"^. For lipemic plasma, however, the solubility coefficient may be considerably higher, up to about 0.33 mmol • 1"^ • mmHg-^ [6].

6. The Concentration of Total CO2, etC02

The extensive quantity expressing the amount of added CO2 in the system is converted to the concentration of total CO2 in the system by division with the volume of the system:

ntC02(S)/7S = etC02(S) . (18)

The quantity has been and still is much employed for description of the acid-base status

of the blood. The reason is, most of all, that the concentration of total CO2 is easy to measure gasometrical ly by means of the Van Slyke apparatus, or automatically by means of

Leonard Skegg's continuous flow analyzer. For clinical purposes, this quantity is generally

used as a measure of a metabolic acid-base disturbance, i.e., as an indirect measure of the excess concentration of titratable acid or base.

The biaarbonate oonoentration of the plasma is often calculated as the concentration of

total CO2 minus physically dissolved CO2:

eHC03(P) = atC02(P) - eC02(P) . (19)

Therefore, the "bicarbonate" concentration generally includes carbonate and carbamate. The

bicarbonate concentration of the plasma is one of the components of the electrolyte scheme

(Gamble diagram) and it is used for calculation of the so-called anion deficit (undetermined

anions)

:

cUA" = cNA"^ - (oCr + cHCOi) . (20)

The concentration of undeterm-jned anions is used as an indicator of an accumulation of organic anions (lactate, e-hydroxybuturate, eto.) in the plasma.

7. Conclusion

In the preceding sections, four fundamental acid-base variables have been described,

relating to the two components H"*" and CO2:

(1) the excess chemical potential of H"*", AyH'*';

(2) the excess concentration of total H"*", t\otW^

;

(3) the excess chemical potential of CO2, AyC02; and

(4) the concentration of total CO2, etC02.

8
Only two of the four variables are inde-pendent variables; the other two being dependent

variables

From a physiological point of view, AyC02 and AetH^ are the independent variables,

because the system is an open system in equilibrium with a gas phase (the alveolar air) with

a certain AyCOa value. In such an open system, it is possible independently to alter the

AyC02 of the gas phase or to add various amounts of strong acid or base (add or remove H+).

Clinically speaking, AyC02 is a parameter (i.e., indicator) of the alveolar ventilation, while

ActH'*' is a parameter of the non- respiratory acid-base balance.

References

[1] International Standard, ISO 31 Series, Technical Committee ISO/TC 12, Quantities, Units,

Symbols, Conversion Factors, and Conversion Tables, International Organization for

Standardization, Switzerland (available from the different National Standards Institutes) .

[2] List of Quantities in Clinical Chemistry, Recommendation 1973, lUPAC, Section on Clinical Chemistry, Commission on Quantities and Units, and IFCC, Committee on Standards, Expert Panel on Quantities and Units, Pure and Appl. Chem. 37, 547 (1974).

[3] Manual of Symbols and Terminology for Physicochemical Quantities and Units, lUPAC,

Division of Physical Chemistry, Commission on Symbols, Terminology, and Units, Pure

and Appl. Chem. 21, 1 (1970).

[4] Quantities and Units in Clinical Chemistry, Recommendation 1973, lUPAC, Section on Clinical Chemistry, Commission on Quantities and Units, and IFCC, Committee on Standards, Expert Panel on Quantities and Units, Pure and Appl. Chem. 37, 517 (1974).

[5] Recommendation for Use of SI in Clinical Laboratory Measurements, ICSH, IFCC, and

WARS, Z. Klin. Chem. n., 93 (1973).

[6] Siggaard-Andersen, 0., The Acid-Base Status of the Blood, 4th ed., 229 pp. (Williams

& Wilkins, Baltimore, and Munksgaard, Copenhagen, 1974).

9

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Gaithersburg, Maryland, July 7-8, 1975. Issued June 1977.

THE BUFFER VALUE OF PLASMA, ERYTHROCYTE FLUID AND WHOLE BLOOD

Ole Siggaard-Andersen, M. R^rth^ and D. A. P. Strickland^

Department of Clinical Chemistry

University of Copenhagen

Copenhagen County Hospital

Herlev, Denmark

The buffer value (6)^ for hydrogen ions in a solution (S) is generally defined as the

Slope of the titration curve at a given pH value when titrating the solution with strong

acid or base (B') [1]'+:

BH^(S)=M_|||. „)

The unit is mol/1, although in certain cases it may be advantageous to use molality {i^):

m'^iS) = 3;?2B'{S)/9pH(S), unit: mol/kg. (2)

An alternative definition of buffer value is:

= ^fj§[f|, (3)

where tC = total added component C. The SI unit is mol^ • 1"^ • J~^. C may be H^ or any

other component, e.g., Oz- Actually, the reciprocal value might be more relevant being the

first partial differential coefficient of the function:

AyC(S) = F(ctC(S), -, -, —). (4)

The relationship between the two buffer values of eqs. (1) and (3) is:

eVcs) • <E • T ' InlO = bh'^(s). (5)

The buffer value of H"*" in a solution can be expressed as the sum of the buffer values of

the solvent (H2O) and the solutes (C). The change in concentration of added base can be

expressed as:

dcB'(S) = deOH"(S) - deH'^(S) + zdcC'(S), (6)

^Present address: Surgery Department D, Rigshospitalet, Copenhagen, Denmark,

^Department of Clinical Measurement, Westminster Hospital, London, England.

^Concerning terminology and symbols, the reader is referred to the preceding paper, e = con- centration, y = chemical potential.

'Figures in brackets indicate the literature references at the end of this paper.

11
where C symbolizes a base group. Dividing with dpH gives:

6H^(S) = 3H"^{H20 c S) + Z6H"^(C c S). (7)

The buffer value of the water is approximately:

BH"^(H20 c S) « (cH^(S) + eOH"(S))-lnlO. (8)

Therefore the buffer value of the water can generally be ignored for 3 < pH < 11.

The buffer value for due to solute C(CH i:? H"^ + C-) is equal to:

BH^(C c S) = 8eC"(S) 3cCH(S)

(9)

8pH(S) 9pH(S)

The equality is only valid for cC = cC" + eCH = constant. If the system is an open system

where CH or C" disappears, the amount which has disappeared must be included in order to obtain equality. This especially applies to the CO2 and NH3 buffers.

The buffer value of the solutes may be expressed as the molar buffer value (g ) or as the specific buffer value (3,,):

where pC(S) is the mass concentration of C in S.

The buffer value is a partial differential and the other independent variables which

are maintained constant must be specified. The buffer value of a pure bicarbonate solution

[S] is widely different for a closed system with constant ctC02 and for an open system

with constant pC02. The following equations are obtained by differentiation of K . = aH • cHC03/(atC02 - cHCOs):

6^H'^(C g S) h 6H"^(C c S)/cC(S),

= S) 5 bH'^(C c S)/pC(S),

(10)

(11)

K

'aak 2.303 etC02 (12)

aak

2.303 • cHC03(S). (13)

12
The maximal molar buffer value for a closed system is therefore 0.576 for pH = pXacA- For

pK = 6.1, pH = 7.4, ctCOa = 25.7 mmol/l, and CHCO3 = 24.5 mmol/l, the values are: BH+(S, closed) = 2.7 mmol/l, and 3H'''(S, open) = 56.6 mmol/l. This illustrates the insignificance of the carbonic acid/bicarbonate buffer in a closed system at pH = 7.4, but the

great importance in an open system [2].

1. Plasma

The buffer value for H"*" in plasma can be expressed as the sum of the buffer values of

the bicarbonate buffer and the non-bicarbonate buffers. The change in the concentration of

added base can be expressed as the following sum (ignoring water, compare eq. (6)):

dcB (P) = dc?HC03(P) + dcX'(P) (14)

Dividing with dpH gives;

eH'^(P) = 3H"^(HC0: c p) + eH'^(X c P) 05)

In the following, X is used as a symbol for "nonbicarbonate buffer."

The buffer value of the nonbicarbonate buffers can be determined as the slope of the

titration curve, when titrating the plasma in the absence of CO2. The value for normal

plasma (pPr(P) = 70 g/1, wAlb(Pr) = 0.60, t = 310.15 K) determined by Siggaard-Andersen and

Rfirth (unpublished experiments), shows a maximum of 7.8 mmol/l at pH = 7.3, while the value

for pH = 7.40 is 7.7 mmol/l. The pH variation of 6 is shown in figure 1.

1 )8X(P) buffer value of non -bicarbonate buffers

I I I I II I I I I I I I I I I

I I I I

(mmol/l)

'80

Human plasma (P)

/)Pr(P) = 70 g/l

i^Alb(Pr) = 0 60

= 37 'C

65 J \ L J \ L

pH(P)

J \ L

Figure 1. The buffer value for H due to nonbicarbonate buffers in the plasma (symbol bH

(X c p) in the text) as a function of pH, in the absence of CO2 as well as at pC02'^45

mmHg [2]. p = mass concentration, w = mass fraction. The maximum at pH 7.3 is probably

due to the imidazole groups of the albumin. The decrease in buffer value for pH > 7 in the

presence of CO2 is due to a change in pz value of the terminal amino groups from about 7.8

(Pr-NHg) to about 5 (Pr-NH-COOH).

13
An alternative method for calculating the buffer value of the nonbicarbonate buffers

of plasma is based on CO2 titration. The reaction following CO2 titration is:

X' + + XH"^ + HCO3, 06)

where X' represents nonbicarbonate buffer base. To be more exact, HCO5 includes CO5 and

PrNHCOO" and is therefore designated tHCO^ = titratable bicarbonate. For varying pC02 we therefore have:

/9ctHC03(P)\ /3cX'(P)\

9PH(P) /,B.(p) \ 9PH(P)4'(P)

As a good approximation we have:

/ 3cX' (P)\ / 3cX' (P)

(17)

,3PH(P) ^.(p) V 3PH(P)/pC02 = 0

(18)

The first coefficient refers to varying pC02(P) and the second to varying AcB'(P). The latter

value should be slightly higher than the former in the pH range about pH = 8 (due to buffering

by terminal amino groups), while the former should be slightly higher in the pH range about

pH = 5 (due to buffering by the carbamate groups, PrNHCOO").

The coefficient (actHCOs/apH) p, can be calculated from experimental values of

(slog pC02/3pH)^g.: ^'^

dlog pCO^ , detHCO^

= - (1 - dpZ/dpH) + ^— ' ^ . (19)

dpH 2.3 ctHCOg dpH

This equation is obtained by differentiating K = cxH^ • ctHCOs/a • pC02, taking a to be

constant, (3log pC02(P)/3pH(P) )^gi read on the curve nomogram provided the data for the

buffer value of nonbicarbonate buffers which were plotted in figure 1 (dotted curve).

According to figure 1 the buffer value of the nonbicarbonate buffers of plasma varies

significantly with pH. As an approximation the value is generally taken to be constant in the physiological pH range, independent of pC02 and pH, varying with the protein concentration

only:

BH^(Xcp) = pPrCP) • 0.11 mol/kg. (20)

For pPr(P) = 70 g/1 the buffer value therefore is 7.7 mmol/1.

Inorganic phosphate accounts for about 5 percent of the nonbicarbonate buffer value

in normal plasma at pH = 7.4 (calculated from an equation anaolgous to eq. (12) with a total phosphate concentration of 1 mmol/1 and P^^^^ = 6.8).

Albumin is the major contributor to the buffer value of nonbicarbonate buffers of

plasma. The imidazoles (16 per albumin molecule) are the most important buffer groups in

the physiological pH range. Since the maximum buffer value in our experiments is at pH = 7.3, we assume that most of the imidazoles have dissociation constants of about lO-''-^. With an albumin concentration of c?Alb(P) = 0.62 mmol/1 {-^p Alb(P) = 42 g/1) and hence an

14
imidazole concentration of 16 • 0,62=>9.9 mmol/l, the maximal buffer value of the imidazoles

at pH = 7.3 is calculated to be 0.576 • 9.9=>5.7 mmol/l, i.e., they account for 73 percent of

the nonbicarbonate buffer value of plasma of 7.8 mmol/l at pH = 7.3. The remainder, i.e., 22

percent => 1.7 mmol/l may be ascribed to the globulins, and with pGloCP) = 28 g/1 , the specific

buffer value of the globulins is calculated to be 0.06 mol/kg, while the specific buffer

value of the albumin is 0.14 mol/kg.

2. Erythrocyte Fluid

The buffer value of the non-bicarbonate buffers of erythrocyte fluid determined by

titrating erythrocyte fluid with strong base in the absence of CO2 is shown in figure 2. ^ I I I I I /?(E)

(mmol/l)

I— 80

I I I I I I I I I I I I

I I I I "~|

14.2

ctDPG(E)

(mmol/l)

HbO

— 20

^> L

6.5

Erythrolysate (E) CtHb(E) = 21-0 mmol/l pCO; = 0 ^ = 37 °C

J I L

7.0 7.5 8.0 J \ \ \ \ I \ \ \ \ L

pHlE) -

J I I »l

Figure 2. The buffer value for H ^in erythrocyte fluid (symbol 3H (E) in^the text) at pCOa =

0, i.e., the buffer value for H by non-bicarbonate buffers (symbol 3H (XcE) in the text),

as a function of pH at 37 °C [2]. Hb = deoxyhemoglobin, Hb02 = oxyhemoglobin. The dotted

curve represents the buffer value of oxygenated erythrocyte fluid with a three-fold increase

in 2,3-diphosphoglycerate concentration (ctDPG). The peak in the buffer value at pH = 7.1 indicates the px value of some of the acid-base groups of the DPG molecule. The difference

between the buffer value of oxygenated and deoxygenated erythrocyte fluid is due to the

fall in pz value of the so-called oxygen-linked acid-base groups following oxygenation.

The value varies with pH, with the concentration of hemoglobin, with the oxygenation of the

hemoglobin, and with the concentration of 2,3-diphosphoglycerate. The value for normal

erythrocyte fluid is about 63 mmol/l for pH = 7.2, cHb = 21 mmol/l, and ctDPG = 4 mmol/l

[2].

2,3-DPG accounts for about 7 to 8 percent of this value. This is calculated fron] pz = 7.1 of two of the phosphate groups of DPG (by means of an equation analogous to eq. (12))

^6H+(DPGcE)«4.6 mmol/l.

The remainder is mainly due to hemoglobin. The molar buffer value of the hemoglobin

therefore is about 58/21 = 2.8. Assuming that the principal buffer groups in the physiological

pH range are the imidazole groups, a maximum of about 6 imidazoles out of the 9.5 imidazoles

of the hemoglobin chains (10 for a chains, 9 for B chains) are engaged in buffering at pH «

7.2. The remainder may be engaged in salt bridge formation.

The difference between the buffer value of oxyhemoglobin and deoxyhemoglobin may be

ascribed to the change in pis: of the so-called oxygen-linked acid-base groups, i.e., acid15
base groups which participate in salt bridges in deoxyhemoglobin but not in oxyhemoglobin.

These salt bridges are the following, including those obtained by interchanging ai with a2 and ei with 32 [3]:

1) a-carboxyl of arginine HC3(141ai) - a-amino of valine NA2(la2)»

2) guanidinium of arginine HC3(141ai) - B-carboxyl of aspartate H9(126a2),

3) e-amino of lysine C5{40ai) - a-carboxyl of histidine HC3(14662)»

4) imidazole of histidine HC3(1466i) - 6-carboxyl of aspartate FGI(946i).

The salt bridge formation causes a rise in dk of the C-terminal imidazole groups (HC3(1463))

and of the N-terminal amino groups (NA2(la)) so that the buffer value of deoxyhemoglobin

becomes higher than that of oxygemoglobin for pH > 7.2, but lower for pH < 7.2. The pK

values are apparently shifted from the range of 7.8 to about 6.3. The greatest difference

between the molar buffer value for oxyhemoglobin and deoxyhemoglobin (about 0.4 at pH 7.3

and 6.3) is less than the maximal difference (0.576) which should be expected if one buffer

group per heme is shifted. The explanation may be that the pK values of the imidazoles and

the valines differ somewhat so that the effect is smoothed.

As a rough approximation, the buffer value of the nonbicarbonate buffers of erythrocyte

fluid may be taken to be independent of pH varying with the total hemoglobin concentration

only [2]:

BH"^(XcE) = 3.0 • cHb(E). (21)

More accurate data for the buffer value of the nonbicarbonate buffers of plasma and

erythrocyte fluid are necessary in order to improve the accuracy of the algorithms for

calculating the relationship among the acid-base variables of blood and plasma (see subsequent paper on acid-base algorithms).

3. Whole Blood

The buffer value is a quantity which may refer to an equilibrium system of several

phases, e.g. whole blood, provided we define the buffer value as:

8etH^(B)

bVO) = —-r , (22)

3yH^(B)

where p is the electrochemical potential: y = m + zFi, where F is the faraday constant

{F = 96487 C/mol), and $ is the inner electrical potential of the plasma. If the electrical

potential of the continuous phase, i.e., plasma, is the reference electrical potential, the

buffer value of whole blood in terms of pH is:

3H''(B)

8eB'(B)

3pH(P)

(23)

The buffer value of the nonbicarbonate buffers of whole blood can be expressed in

terms of the values for plasma and erythrocyte fluid. We have:

cX'(B) = (^P(B) . eX'(P) + <(>E(B) • cX'(E), (24)

where <j)P(B) + (})E(B) = 1. Differentiation of this equation with respect to pH, and rearrangement (regarding (j)E(B) as constant) gives:

16
dcX'(B) dcX'(P) /deX'(E) dpH(E) deX'(P)'

= + (|,E(B) ; dpH(P) dpH(P) \dpH(E) dpH(P) dpH(P)

(25)

SH'^CX c B) = eH"^CX c P) + <^E(B) • CBH'^CX c E) • - eH'^(X CP)). (26)

The relationship between pH(P) and pH(E) is given by (see eq. (6) of the subsequent

paper) [4]:

pH(E) = 7.19 + 0.77 (pH(P) - 7.40), (27)

for oxygenated blood at 37 °C. Inserting dpH(E)/dpH(P) = 0.77 as well as BH'^(Xcp) = 7.7 mmol/1 and 3H+(XcE) = 63 mmol/1 (for cHb(E) = 21 mmol/1) into eq. (26) gives:

BH"^(XcB) = (7.7 + (^E(B) • 40.8) nimol/1, (28)

eH"^(XcB) = 1.94 • cHbCB) + 7.7 mmol/1. C29)

In analogy with eq. (17) we have:

/8ctHC0T(B)

eH (X^B) = - 9pH(P) /cB'(B)

(30)

where the concentration of bicarbonate refers to whole blood.

Often (3ctHC03(P)/3pH(P))^g, is used as an indicator of the buffer value of non- bicarbonate buffers in the blood, where the concentration of bicarbonate refers to the

plasma phase. We have:

etHC03(B) = ctHC03(P)-cl)P(B) + etHC03(E) • <j)E(B)

= etHC03(P) -d-d - rc) • <).E(B)), (31)

where <t>P{Q) + (f>E(B) = 1 , and reHC0q(P|E) = etHC03(E)/ctHC03(P) . Regarding ())E(B) as constant,

differentiation with respect to pH(P) and rearrangement gives:

(32)

3dtHC0"(B) drc

+ (l)E(B) • etHCO"(P)-

8etHC03(P) 9pH(P) ^ dpH(P)

9pHCP) (1 - (1 -re) • 4.E(B)

By application of l/(l-x) = ^ + x + x + ... (for |£|< 1), we obtain:

/9ctHC0"(P)\ / . dre \

"(~^1b.,B)

= ^ • Wi) • (33,

(1 + (1-rc) • <j,E(B) + (d-rc) • <}.E(B))^ +

17
Inserting 6H^(X c B) from eq. (26) we obtain:

/ 3ctHC0:^(P)\ , o

. 2 = 6H (Xc P) + ^ . (^E(B) +A (1-re) • {<^E{B)r

V 3pH(P) /cB'(B) , , , + A . UE(B))2

where:

^ = eH^X e E) . - eH"(X e P) . ra -H etHCO-CP) • ^)

(34)

(35)

We assume that the small diffusible ions are at equilibrium across the red cell membrane

and hence:

aH'^(P) aCr(E) aHCOZ(E)

ra = — = = ^ . (36)

aH (E) aCr(P) aHC03(P)

Conversion of raHCOs to rcHCO^ = cHC03(E)/eHC03(P) requires a knowledge of the ratio of the

activity coefficients (y) and the mass concentrations of H2O (p):

yHCO"(P) pHpO(E)

rcHCO::(P/E) = raHCO;(P/E) •

• — . (37)

yHC03(E) pH20(P)

Experimental data indicate that yC1"(E)/yC1"(P) = 0.93 [1,3]. Using this value for

YHC03(E)/YHC0i(P) and using pH20(P) = 0.94 kg/1, and pH20(E) =0.73 kg/1, we have:

rcHCO" = raHC03 • 0.835. (38)

From eq. (27) we obtain dpH(E)/dpH(P) = 0.77, and for pH = 7.40 we further obtain ra = 0.615 and dra/dpH(P) = -2.303 (1-0.77) ra = -0.326. Hence ra = 0.515 and dre/dpH(P) = 0.272. Inserting these values together with ctHC02(P) = 24.4 mmol/l into eq. (35), we

obtain: A = 37.9 mmol/l. As an approximation ignoring the variation with ctHCOatP) and

with pH(P) we therefore have:

/3cHC0Z(P)\ , , , , , , xv2

- JlJ-] / mmol/l = 7.7 + 37.9 <{.E B) + 18.4 (<J)E(B))^

\ 3pH(P /cB'(B) , , xx3 f',n\ ^ ^ +8.9 (<^E(B))-^ + .... (39)

As a further approximation fitting this equation for <))E(B) = 0 and (j)E(B) = 0.45 and for

cHB(E) = 21 mmol/l we get [2]:

/9eHC0;(P)\

- ^ /(mmol/l) = 7.7 + 48.4 (f>E(B) (40)

\ 3pH(P) /cB'(B)

'

= 7.7 + 2.3 eHb(B)/(mmol/l). (41)

This is the approximation which is often employed for the slope of the pH, ctHCOj equilibration curve of whole blood. The present derivation indicates where the principal approximations have been made: (1) assuming the buffer values of nonbicarbonate buffers of plasma and

erythrocyte fluid to be independent of pH, (2) ignoring the variation with ctHCOa at constant

18
pH(P), (3) ignoring the variation with pH(P) at constant ctHCOs (P), and (4) assuming a linear variation with 4)E(B).

References

[1] Van Slyke, D. D., On the measurement of buffer values and on the relationship of buffer

value to the dissociation constant of the buffer and reaction of the buffer solution, J. Biol. Chem. 52, 525 (1922).

[2] Siggaard-Andersen, 0., The Acid-Base Status of the Bloody 229 pp. (Williams & Wilkins,

Baltimore, and Munksgaard, Copenhagen, 1974).

[3] Perutz, M. G., Stereochemistry of cooperative effects in haemoglobin, haem-haem interaction and the problem of allostery, the Bohr effect and combination with organic

phosphates. Nature (Lond. ) 228, 726 (1970).

[4] Funder, J. and Wieth, J. 0., Chloride and hydrogen ion distribution between human red

cells and plasma. Acta Physiol. Scand. 68, 235 (1966).

19

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

ACID-BASE ALGORITHMS

Ole Siggaard-Andersen

Department of Clinical Chemistry

University of Copenhagen

Copenhagen County Hospital

Herlev, Denmark

A knowledge of the relationships between the acid-base variables of the blood in

vitro is of considerable importance, because it is often required to calculate the re- maining variables on the basis of measurement of only two of the variables. For these

calculations various nomograms have generally been employed, e.g., the Van Slyke and

Sendroy nomogram [1]^, McLean's Z-nomogram [2], the Singer and Hastings nomogram [3], and

others [4,5]. With the development of small programmable calculators, an arithmetic calculation will probably in the future be more suitable. The problem will often be that one wishes to calculate the base excess concentration of the blood, or perhaps of the average

extracellular fluid, on the basis of pH and pC02 values measured at 37 °C. Process diagrams

for two arithmetic algorithms for this purpose are shown in figures 1 and 2. The rectangular blocks are procedure symbols [6] and each of these will be described separately in the

following. The word algorithm (or algorism) is derived from al Kuwarizmi, a Persian

mathematician from the 8th century.

pC02(B) Henderson & Hasselbalch

equation

eq (1)

cHC03(P)

^

( \

PH(P) Van Slyke

equation for

whole blood

eq (4)

AcB'(B)

eHb(B)

I Figure 1. Procedure diagram for calculation of the bicarbonate concentration

I of the plasma, cHCOi(P) and the base excess concentration of blood or extraI cellular fluid AcB'(B) or AcB'(Ecf) after measuring the blood pC02, the plasma

pH, and the hemoglobin concentration of the blood cHb(B). The HendersonHasselbalch equation is employed for calculation of the plasma bicarbonate

concentration. The Van Slyke equation for whole blood is employed for

calculation of the base excess concentration. The various equations are described in the text. w

1 Figures in brackets indicate the literature references at the end of this paper.

21
pCOaCB) Henderson and Hasselbalch

r eHCOi(P)

equation

eq \ \

)

pH(P) Van Slyke AcB' (P)

equation for

plasma

eq (2)

pH(P) - pH{E)

equation

eq (6)

cHCOiCE)

equation

eq (7)

cHb(B)

Van Slyke

equation for

erythrocyte fluid

eq (3)

P-E-B equation

eq (5)

AcB'CB)

Figure 2. Alternative algorithm based on intermediary calculation of the pH

of the erythrocyte fluid (pH(E)) and the bicarbonate concentration of the

erythrocyte fluid (cHCOslE)). The Van Slyke equation for plasma respectively

erythrocyte fluid is used for calculation of the base excess concentration

of the plasma (acB'(B)) respectively erythrocyte fluid (acB'(E)). The

various equations are described in the text. Another almost identical

algorithm uses the Henderson-Hasselbalch equation for erythrocyte fluid

to calculate cHC03(E) from pH(E) and pCOzCB).

1. The Henderson-Hasselbalch Equation

The first equation of the algorithms is the Henderson-Hasselbalch equation, which can be expressed as follows for human plasma at 37 °C:^

cHCO:

pH = pz + log

aC02 •pC02

(1)

pK = 6.10.

aCOg = 0.0306 mmol/1 • mmHg (= 0.231 mmol/1 • kPa) .

Concerning terminology and symbols the reader is referred to the preceding paper on definitions of acid-base quantities. HCO3 includes COi and carbamate throughout.

22
The value for pK varies slightly with the composition of the plasma, with a normal biological standard deviation of about 0.0015. In pathological cases with large changes in

the ionic strength the variation is considerably greater. Due to the incorporation of

carbonate and carbamate in the "bicarbonate" concentration an apparent variation of pK

with pH appears (decreasing pK with increasing pH), but for most practical purposes the

variation of pK can be ignored. The Henderson-Hasselbalch equation can be derived by

expressing the equality of the sum of the chemical potentials of CO2 and H2O and the sum

of the chemical potentials of and HCO3.

2. The Van Slyke Equation

The CO2 equilibration curve of the blood can, as a good approximation, be represented

by a straight line in a pH, log pC02 coordinate system, and this is employed in the nomo- graphic algorithms. Another good approximation to a straight line is obtained by plotting

the CO2 equilibration line in a pH, CHCO3 coordinate system (Davenport diagram [7]). This

was originally utilized by Van Slyke [8], and I therefore find it appropriate to call the

equation for the pH, CHCO3 equilibration curve the Van Slyke equation. For plasma the Van Slyke equation can be written:

AcHC03(P) = -bH"^(XcP) . ApH(P) + AcB'(P), (2)

where

AcHC03(P) = AcHC03(P°lP) = eHC03(P) - 24.4 mmol/1,

ApH (P) - ApH(P°|P) = pH(P) - 7.4,

AeB'(P) = AcB'(P°|P) = cB'(P) - eB'(P°).

P° is plasma titrated with strong acid or base to pH(P) = 7.40 at pCOa = 5.33 kPa and T = 310.15 K ( ^cHCOg = 24.4 mmol/1).

3H'^(X c p) is the buffer value for H"*" due to non-bicarbonate buffers in the plasma.

According to eq (20) of the preceding paper it can be expressed as:

eH'^(XcP) = pPr(P) • 0.11 mol/kg.

i.e. J, for normal plasma with pPr(P) = 70 g/1 ^ bH'^(X c p) = 7,7 mmol/1.

The Van Slyke equation is derived from:

AcB'(P) = AeHC03(P) + AcX'(P),

and for pH(P) = 7.40 we have AeX'(P) = 0 (compare eq (14) of the preceding paper).

For erythrocyte fluid the Van Slyke equation is analogous:

AeHC03(E) = - 3H'^(X c E) x ApH(E) + AeB'(E), (3)

where E° is erythrocyte fluid in equilibrium with P°, i.e., pH(E°) = 7.19 for pC02 = 5.33

kPa and 37 °C => eHC03(E°) = 12.6 mmol/1. The buffer value for H"*" due to the non-bicarbonate

buffers is

3H"^(X c E) = 3.0 • cHb(E),

23
i.e., for normal erythrocyte fluid with cHB(E) = 21 mmo1/l eH'''(X c e) = 63 mmol/1 (see eq

(21) of the preceding paper).

For whole blood the Van Slyke equation expressed in terms of the plasma bicarbonate

concentration is:

AeHC03(P) = - B-ApH(P) + AeB'(B)/Z (4)

where B = (3cHC03(P)/3pH(P))^^g, ,gv can be expressed by the following approximation (eq

(41) of the preceding paper):

B = 2.3 • eHb(B) + 7.7 mmol/1.

z is derived from eq (31) of the preceding paper:

pH(P) = 7.40

AcB'(B) = AcHC03(B)

= AeHC03(P) • (1 - (1 - re) • (|)E(B)),

i.e. Z = (1 - (1 - rc) • (t)E(B),

where rc = eHC03(E)/eHC03(P) = 0.515 (for pH(P) = 7.40). For eHB(E) = 21 mmol/1 we therefore

get:

Z = 1 - 0.023 • eHb(B)/(mmol/l).

The relationship between the base excess concentration of plasma, erythrocyte fluid,

and whole blood is the following:

AeB'(B) = P(B) • AeB'(P) + <j)E(B) • AcB'(E) (5)

= (1 - (f>E(B) • AeB'(P) + <|)E(B) • AeB'(E).

For eHb(E) = 21 mmol/1 (^E(B3 = cHB(B)/(21 mmol/1).

3. pH(E) - pH(P) Equation

The relationship between pH(E) and pH(P) measured in normal blood can be expressed as follows (eq (27) of the preceding paper):

pH(E) = 7.19 + 0.77 • (pH(P) - 7.40). (6)

The coefficient 0.77 is due to a pH variation of the concentration of "non-diffusible"

ions {e.g., protein anions) in plasma and erythrocyte fluid. The coefficient can be

calculated approximately from the buffer value for H"*" due to non-bicarbonate buffers in

plasma and erythrocyte fluid, assuming the erythrocyte membrane is impermeable to cations

for short term changes.

24
The value 7.19 is dependent on the membrane potential across the erythrocyte membrane

(normally approximately -13.0 mV, inside negative) which is due to the active transport of

Na"*" out of the erythrocytes. Inhibition of the Na"*" pump causes a diminution in the membrane

potential and a swelling of the erythrocytes (hemolysis). Increased activity of the Na'^ pump causes an increase in the membrane potential and therefore a lower pH(E) for pH(P) = 7.40.

According to eq (38) of the preceding paper we have:

re = ra • 0.835,

where re = cHC0;(E)/eHC03(P) , and ra = aHC0;(E)/aHC03(P) = aH'^(P)/aH"^(E) . By combination of

this equation and eq (6) we get:

cHC03(E) = cHC03(P) • 0.835 • antilog (1.492 - 0.23 pH(P)) (7)

4. Comparison of the Van Slyke Equation and the Alignment Nomogram

The nomographic algorithm is based on a linear representation of the CO2 equilibration

curves of the blood in a pH, log pC02 coordinate system [3,4]. This algorithm is more difficult to express in a simple arithmetic form because the slope of the equilibration

lines varies with both the_hemoglobin concentration and the base excess concentration of

the sample. In a pH, CHCO3 coordinate system the variation of the slope with the base excess concentration (or with pH) is small. Both algorithms are approximations. A comparison

of the results obtained with the arithmetic algorithms and the nomographic algorithm

(Alignment Nomogram) [4] is shown in table 1. The table includes values for the calculated

Table 1. The base excess concentration of the blood and the extracellular

fluid calculated (A) by means of the alignment nomogram [4],

(B) by means of the Van Slyke equation for whole blood (eq (4)),

and (C) by means of the Van Slyke equations for plasma and erythrocyte

fluid (eq (2) and (3)) together with the relationship between pH(P)

and pH(E) (eq (6)). The calculations are based on a hemoglobin

concentration of: cHb(b) = 9.0 mmol/1, and eHb(Ecf) = 3.6 mmol/1.

Measured Calculated

pH(P) pCOp(B)

(mmflg)

AcB'(B)

(mmol/1)

AeB'(Ecf)

(mmol/1

)

Acute hypercapnia 7.09 100 A -4.6 ±0.0

B -2.6 +0.5

C -2.3 +1.0

Acute hypocapnia 7.70 16 A +0.5 -2.8

B +2.9 -0.1

C +3.0 -0.3

Chronic hypercapnia 7.30 80 A +8.2 +11.7

B +9.2 +11.7

C +9.4 +11.9

Chronic hypocapnia 7.44 25 A -5.2 -6.5

B -5.2 -6.4

C -5.1 -6.5

Acute base deficit 6.90 30 A -27.6 -25.0

B -26.0 -24.4

C -27.0 -24.3

(Chronic) base deficit 7.15 10 A -25.8 -24.6

B -22.3 -22.9

C -22.8 -22.9

(Chronic) base excess 7.60 55 A +26.0 +29.0

B +27.4 +29.4

C +26.7 +28.9

25
excess concentration of base in the average extracellular fluid (including blood). This

value is calculated on the assumption that the buffer value for H"*" due to non-bicarbonate

buffers in the extracellular fluid is mainly due to hemoglobin, and that the hemoglobin

concentration of the total extracellular fluid (including blood) is approximately 3.6

mmol/1 [4]. The advantage of this quantity is that it remains virtually constant during

acute changes in the pC02(B) in vivo, whereas the base excess concentration of blood and

plasma varies in opposite directions due to a redistribution of H"*" between phases of

different buffer value.

The results indicate that the base excess concentration of the blood or extracellular

fluid can be calculated with sufficient accuracy for all clinical purposes by means of the

arithmetic algorithms. In spite of the approximations made in the derivation of the Van Slyke equation for whole blood where the variation of B with cHC03(P) at constant pH(P) is ignored, there seems to be no reason to use the more complicated algorithm based on the

empirical relationship between pH(P) and pH(E). The major approximation undoubtedly lies

in the use of a constant buffer value for non-bicarbonate buffers in plasma and erythrocyte

fluid irrespective of a considerable pH variation.

When the pC02 value falls in the range between 25 and 80 mmHg, the alignment nomogram

is likely to provide the best approximation, because it was constructed from experimental

data in the pCOa range. When the pH is around 7.3 to 7.5 but the pC02 is very high or

low, the Van Slyke equation is likely to provide the best approximation, because the non- bicarbonate buffer values is most accurate for this pH range.

References

[1] Van Slyke, D. D. and Sendrby, J., Jr., Studies of gas and electrolyte equilibria in

blood. XV. Line charts for graphic calculations by the Henderson-Hasselbalch

equation, and for calculating plasma carbon dioxide content from whole blood con- tent, J. Biol. Chem. 79_, 781 (1928).

[2] McLean, F. C, Application of the law of chemical equilibrium (law of mass action) to biological problems, Physiol. Rev. 18, 495 (1938).

[3] Singer, R. B. and Hastings, A. B. , An improved method for the estimation of disturbances of the acid-base balance of human blood. Medicine (Baltimore) 2J_, 223 (1948).

[4] Siggaard-Ander^en, 0., The Acid-Base Status of the Blood, 229 pp. (Williams and Wilkins,

Baltimore and Munksgaard, Copenhagen, 1974).

[5] Thews, G. (ed.), Nomogramme zum S'dure-Basen-Status des Blutes und sum Atemgastransport,

134 pp. (Springer-Vorlag, Berlin, 1971).

[6] International Standard ISO 1028, Flowchart Symbols for Information Processing,

International Organization for Standardization, Geneva, Switzerland (available from

the different National Standards Institutes).

[7] Davenport, H. W., The ABC of Acid-Base Chemistry, 119 pp. (University of Chicago Press,

Chicago, 1969).

[8] Van Slyke, D. D., Studies of acidosis: XVII. The normal and abnormal variations in

the acid-base balance of the blood, J. Biol. chem. 48, 153 (1921).

26
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DETERMINATION OF TOTAL CO2 CONCENTRATION IN BLOOD OR PLASMA

P. Rispens, E. J. van Kampen, and W. G. Zijlstra

Laboratory of Chemical Physiology

University of Groningen

Groningen, The Netherlands

and

Laboratory of Clinical Chemistry

^ Diakonessenhuis

Groningen, The Netherlands

Carbon dioxide is present in blood as dissolved CO2, carbonic acid (H2CO3), bicarbonate

(HCO3), carbonate (C0|~) and carbamate, i.e., CO2 bound to free aminogroups of proteins

(RNHCOO"). The quantity total CO2 concentration (c^q ) is defined as the sum of the concen- trations of all forms in which CO2 is present. In most methods [1]^ for the determination of

in blood or plasma (c^q and c^q , respectively), blood or plasma is added to an acid

reagent for conversion of bound CO2 (HCO3, CO3" and RNHCOO") into free CO2 {H2CO3 and dissolved CO2). As a measure for c^^q^ may then be used

1) the total amount of CO2 which can be extracted from the mixture of sample and

reagent (extraction methods);

2) the rise in P^q in a given volume of gas which is equilibrated with the samplereagent mixture (equilibration methods);

3) the rise in P^q^ in the sample-reagent mixture when escape of any CO2 is prevented.

In the method developed in our laboratory (the cediometer), the equilibration principle is

applied [2,3].

1. Extraction Methods

Classic extraction methods are the volumetric [4] and the manometric Van Slyke techniques

[5,6]. More recently gas chromatographic [7,8] or mass spectrometric [9] techniques have

been applied to determine the total amount of CO2 extracted from the sample-reagent mixture.

The main problem in applying these techniques is to ascertain that all CO2 has indeed been

extracted from the sample-reagent mixture. Under favorable conditions, highly reproducible

results can be obtained (coefficient of variation s=l percent).

The absolute accuracy depends on the reliability of the calibration procedure. When

gas chromatographic or mass spectrometric techniques are used, Na2C03 or NaHC03 reference

^Figures in brackets indicate the literature references at the end of this paper.

27
solutions are used for calibration. The accuracy then obviously depends on the reliability

of the reference solutions.

The volumetric or manometric techniques are usually considered as absolute methods,

i.e., methods in which calibration with reference solutions is not required. The accuracy

depends on the accuracy of the volume or pressure measurements. In the volumetric technique,

CO2 is extracted by shaking and vacuum exposure. The volume of the CO2 set free is measured

at barometric pressure. Corrections are necessary for other gases (O2, N2) set free from the

sample-reagent mixture. Because of the uncertainty of these corrections the accuracy of the

rather simple volumetric technique is dubious.

The more laborious manometric technique of Van Slyke and Neill [5] is usually considered

as the most accurate method for determining c^q . In this method, CO2 is also extracted from

the sample-reagent mixture by shaking and vacuum exposure. After all gas has been extracted,

it is reduced to a known volume v and the pressure is determined. CO2 is then absorbed using

a concentrated NaOH solution and the pressure of the remaining gas is again measured at volume 7. The pressure difference before and after absorption of CO2 corresponds to the

pressure in the volume F of the CO2 set free from the sample, c^q of the sample is calculated by multiplying the pressure difference with a factor depending on F, the sample

volume, the acid reagent volume and the temperature. This factor, which can be obtained

from reference [10], contains a correction for the amount of CO2 redissolving in the samplereagent mixture while the gas is reduced to volume 7. The correction has been determined

using Na2C03 reference solutions. However, there is some doubt about the magnitude of this

correction, because the amount of CO2 that redissolves depends upon the apparatus used and

upon the skill and speed of the operator. With utmost care and skill, highly reproducible

c^Q determinations are possible (coefficient of variation 0.6 percent). However, because

of the need for Na2C03 reference solutions for finding the correction for CO2 redissolving in

the sample-reagent mixture, the manometric Van Slyke technique is no absolute method either.

2. Equilibration Methods

For the equilibration of a fixed gas volumq with a fixed volume of blood or plasma

after the addition to acid reagent, a closed system containing a circulating pump is usually

employed. For the determination of the rise in v^^ of the gas mixture, photometric [11],

potentiometric [2,3], gas chromatographic [12] and mass spectrometric techniques, as well as infrared gas analysis [13] can be used. Under favorable conditions highly reproducible

results can be obtained (coefficient of variation *1 percent). Na2C03 or NaHCOs reference

solutions are used for calibration.

In the cediometer, equilibration of the sample-reagent mixture with a fixed gas volume

is performed in a closed system consisting of a measuring chamber filled with NaHCOs-NaCl

solution, a sample chamber filled with acid reagent, and tubing connecting measuring chamber

28
Figure 1. Schematic diagram of the closed gas-liquid circuit of the

cediometer. E = measuring chamber; S = sample chamber; T = gas tubing;

A = CO2 absorber; C = cock; P = pump.

and sample chamber (fig. 1). The gas mixture is driven through the system by means of a pump. The rise in p^q in the gas mixture caused by adding a constant volume of blood or plasma to the acid reagent, is determined by the simultaneous equilibration of the gas

mixture with the HCOs-NaCl solution and measuring the ensuing change of pH (ApH^) by means

of a combined glass-Ag/AgCl electrode. ApH^ is a measure of the total CO2 concentration

of the sample injected into the acid reagent. The method is easy to perform; the time

necessary for a determination is about 3 minutes. The coefficient of variation is 0.8

percent.

3. Methods Based on -Rise in p^q of Sample-Reagent Mixture

For the determination of the rise in p^q of the sample-reagent mixture, a CO2 electrode

can be used. The method is quite simple and rather reproducible results have been reported,

the coefficient of variation being «1.2 percent [14]. Na2C03 or NaHCOs reference solutions

are used for calibration.

4. Discussion

In spite of the simplicity of many available methods for determining e practice in most laboratories is to measure pH and p^q and to calculate

using the Henderson-Hasselbalch equation. Yet the direct determination of e distinct advantages over that of P-^ . LU9

CO;

, common

has some

29
1) Calibration can easily be performed with NaaCOs or NaHCOs reference solutions, which

can be prepared from commercially available analytical grade reagents.

2) In contrast to P^q , c^q is inde-pendent of temperature. Therefore, mistakes due to incorrect thermos tating of the measuring system are excluded.

3) c^Q^ is faj> less influenoed by metabolism of anaerobical ly stored blood than P^q .

4) c^Q is less influenoed by escape of CO2 from the sample than P^q .

5) The instrumentation required for measuring c^^ is less vulnerable than membrane

covered CO2 electrodes are.

Because of these advantages, the determination of c^q should be recommended, especially

in smaller laboratories where some time passes between sampling and analysis or between the

determination of pH and p^q .

References

[I] Rispens, P. and Zijlstra, W. G., Blood gas transport and acid-base balance, in Clinical

Biochemistry, Principles and Methods, H. CH. Curtius, and M. Roth, eds., p. 1604 (Walter

de Gruyter, Berlin, New York, 1974).

[2] Rispens, P., Van Assendelft, 0. W., Brunsting, J. R., Zijlstra, W. G., and Van Kampen,

E. J., A direct method for the determination of the HCO3 concentration as total carbon

dioxide in blood and plasma, Clin. Chim. Acta, 1_4, 760 (1966).

[3] Rispens, P., Brunsting, J. R. , Zijlstra, W. G., and Van Kampen, E. J., Determination of

total carbon dioxide in blood and plasma by means of the cediometer. Theory and experimental verification, Clin. Chim. Acta, 22_, 261 (1968).

[4] Van Slyke, D. D. , Studies of acidosis. II. A method for the determination of carbon

dioxide and carbonates in solutions, J. Biol. Chem. 30, 347 (1917).

[5] Van Slyke, D. D. and Neill, J. M., Determination of gases in blood and other solutions

by vacuum extraction and manometric measurement, J. Biol. Chem. 6J_, 523 (1924).

[6] Natelson, S., Routine use of ultramicro methods in the clinical laboratory, Amer. J. Clin. Path. 21_, 1153 (1951 ).

[7] Ramsey, L. H., Analysis of gas in biological fluids by gas chromatography. Science,

129 , 900 (1959).

[8] Wilson, R. H., Jay, B. E., Doty, V., Pingree, H., and Higgens, E., Analysis of blood

gases with gas absorption chromatographic technique, J. Appl. Physiol. 1_6, 374 (1961 ).

[9] Lotz, P., Dahners, H., and Pichotka, J. P., Massenspektrometrische Bestimmung des O2

und C02-Gehaltes von Blut, Pflugers Arch. 3]5_, 86 (1970).

[10] Bartels, H., Biicherl , E. , Hertz, C. W., Rodewald, G., and Schwab, M.

,

Lungenfunktionsprufungen, Methoden und Beispiele klinischer Anwendung, p. 219

(Springer Verlag, Berlin, 1959).

[II] Zijlstra, W. G., Brunsting, J. R., Dorlas, J. C, and Rispens, P., Eine Methode zur fortlaufenden COa-Messung in der Atemluft und zur schnellen Bestimmung des

Gesamtkohlensauregehaltes von Blut und Plasma, Anaesthetist, 13^, 200 (1964).

30
[12] Gimeno, F. 0. and Tammeling, G. J., A recirculation system for the determination of

blood gases by chromatography, J. Appl. Physiol. 24, 119 (1968).

[13] Gimeno, F. 0., Orie, S. A. M. , and Tammeling, G. J., Determination of carbon dioxide

content by infrared analysis, J. Appl. Physiol. 21_, 1377 (1966).

[14] Reyes, R. J. and Neville, 0. R., A rapid electrochemical technique for measuring the

total carbon dioxide content of blood, Clin. Chem. 14^, 637 (1968).

31
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i

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National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

THE APPARENT OVERALL FIRST DISSOCIATION CONSTANT OF CO2 IN PLASMA

P. Rispens and W. G. Zijlstra

Laboratory of Chemical Physiology

University of Groningen

Groningen, The Netherlands

Since Hasselbalch [1]^ in 1916 put the equilibrium equation for the dissociation of

carbonic acid in the logarithmic form, many investigations have been devoted to the determination of the apparent overall first dissociation constant of CO2 in plasma (table 1).

''HCO" pH = pz] + log (1) ^ ^COa

p^i denotes the negative logarithm of the apparent overall first dissociation constant of

CO2, ^HCOs concentration of HCO3 (mmol-l'M, S the solubility coefficient^ of CO2

(mmol 'l"^' (mm Hg)"^) and P^^ the CO2 tension (mm Hg). Introduction of any new method for

UU2

, the determination of pH, cuCOg '^^ ^COa ''"variably resulted in redetermination of p^i. It proved that the value of pz{ depended on the analytical methods. To stress the experimental

nature of the constant, the symbol p^| was introduced, denoting the value calculated for

pK\ using eq. (1) when experimentally found values for pH, i^^^q-, P^q and S are used. In the p^i determinations of the twenties and early thirties [2-7], P^g was calculated from

Ppri of the gas mixture with which the plasma had been equilibrated. pH was measured using

LU2

hydrogen electrodes. <2^qq- was calculated from the total CO2 concentration c^q as measured

with the manometric Van Slyke technique, assuming c^^ to be equal to the sum of c^^qand S'P^Q 'S had been determined by experiments in which acidified plasma or serum was equilibrated with gas mixtures of known P^q and c^q measured manometrical ly . A mean value

of 6.10 was found for p^i at body temperature with no appreciable variation, even when

abnormal plasma was used.

However, with the introduction of newer methods for pH and Pm measurement, doubt arose

ix, UU2

about the constancy of pzj. The introduction of glass electrodes for the determination of

pH resulted in reports in which a decrease of the value of pzj with increasing pH was

stated, Api^l/ApH ranging from -0.03 to -0.06 [8-10,12]. The introduction of CO2 electrodes

for the direct measurement of p^q resulted in investigations [13,14] in which appreciable

variations of pxj in seriously ill patients were measured (from 5.96 to 6.27). The tem-

^Figures in brackets indicate literature references at the end of this paper.

^ ^^C02 dissolved ^H2C03^ ' ^^C02^ '

33
J2 CO

--I II c?^ II Q. n:

+-) a.

•(-> o\ II

CL CL 0 ^

-(-> c<1 0 01

'r-

-l-> n3 ro 0;

c:

'r-£

S.

<u «^

-t-> (U 0

II 0. Q-

</) 0•p->

(U

s-Q.

«4- 0

(U

1/) a>

4-> Q. 10 3

CO

O) q:

Q.E O

<U o

S4->3

CD

SZ3 -Q

SG

CU>> 0 CD CO CO CO CM 00 CM ==1- CM CM 0 CD CO 0 CD 0 • 0• 0 0 0

<x> tX) to to to to to to to to • to 0 >vl- CO LO 0 CO ^ to LO 0 0 0 0 0 0 0 0 01 1 0 01 1 01 0 (J T3 OJ Min CO CO CO CD CM 0 0 to CO CO CO 0 0 0 0 0 0 0 0 0 01 0 01 1 c 01 1 0 . _ <^ 00 ' <^ 0^ to CM f—i

CM CO CO to 1— CM 0 — !— CM 0

10 lO <x; >X) to tc to1 to to to to to to to 0 to to 1 1 0 CM 1 VX> 1 1 r— "^1- 1 1 1 1 1 1 1 CO M CO CO CM CD LO CM 00 CM 0 CM 0 t— 0 0 0 1— 0 0 0 0 r— 0 1— f— 0 —

00

to to to to to to to to to to to to

LO CO ^X) 0^ 1

0 CM to i

CD CM CO CO i— CM 0 C\l to 'Jd" c^ 0^ CO LO CTl 0 <X3 1— CM 0 1— * 0 0 f— 0 1

• to to to to to to to to to to to to to to

00 LO LO CO CO 00 to LO CO to CO to to CM 0 CM CM CM r— CM CM 1— 1— 1— to >^i- CM

in CO , p— CO 0 0 LO to LO 0 0 1X1 to to Ln LO LO r- 0 0 to to 00 00

06 00 1^ 00 CO CO 1 1 1 1 1 1 LO 1 1 1 1 1 1 I 1 CO p~ CM to ,— LO CM LO (—i

"d- CO CO 0 CM CO CM 0 0 CO '^f CM

IX) IX) 1^ r-- r-- 1^ to to to to to

LO

CO C_) CO rr\ 00 f—^

CO 00 CO LO CO 00 CO LO

CO CM CO CD CO CM CO CO CO CO CM CO CM CO CO CO CM • ' ' _ 1—

fo 3

to ' , ro i_ 0 ^ }

<u -E m (/) 4-* ^ _ _ ' ^

ro

<o 0 0 -Q

O- tn

C/) E 'J LJ ~1 ro ^ E ^ 'O Si X E ^ O- ro ro 0 E rt3 E <

<

cz J-: cz

t.^ 3 o> ro (/) "O S- S- B ^ E S- E E S- 0 T5 ro B ro E <u <u rtS 3 O) -0 E 1

(/> ro to in S- 1— I/) S- to ro CL ro Q.

<D (0 I— <U <U OJ <u lO 0 10 E ro to E E E E Q. E

10 to to E E ro ro ro ro ro

s- S- S- CD CD s- E S- CD E E E E CD E 0 0 0 0 0 0 3 0 0 3 3 3 3 0 3

3: n:

Q r— z: -o ^ in 0 3Z

4^ 3O 1 1 1 1 E0 1—1

to CO s_

a 1 1 CU

(/) TD 3 E

Q) ro 1—1

=1 « ^ CD

to E CDi 1 -0

CD 0 E i- E to ro E Z. ro

4-> CDl

—1

to > cno

ro 0 Q 00 00 1—1

<i)

CU L—1 =1 1—1 1 1 «E

•r— -P

to3

cC

>> J3E

<u>

to

ro

U)

0)3

to ro E > 0

+J

roE CD

•r— E E •1— S- to <U 3

+J

0) T3

<U +J

ro 0 r—3

s-<u r—

XI ro E 0 3 0) E sII

II zc E Q. a.

00to

II

>>^ o

II II

CL o. Q. 3 «u 4-

34
perature dependence ApKi/^T, usually inferred from measurements of but two temperatures, was reported to vary between -0.0032 and -0.0063 pZ] unit-(°C)"^.

After having developed a new method for the determination of c^q in blood or plasma

[15] and a tonometry set-up [16] allowing blood or plasma to be equilibrated with gas

mixtures of known composition, we decided to redetermine pK'i over the temperature range of

16-42 °C and the pH range of 6.8-7.8. Plasma of healthy volunteers was equilibrated in

horizontally rotating tonometers at 16, 20, 26, 30, 32.5, 35, 37.5 and 42.5 °C with CO2/O2

gas mixtures from cylinders, F^q ranging from 2-16 percent. F^q was measured with a Haldane gas analysis apparatus, the standard deviation calculated from duplicate determinations being 0.04 percent. After equilibrium, two 5 ml samples were taken from each

tonometer. Plasma pH was measured with capillary microelectrodes (Radiometer E5021a)

connected to a direct reading pH meter (Radiometer PHM27) or a balancing pH meter (Radiometer PHM4). Level and slope of the pH measuring system were controlled with two NBS

buffers, the slope daily, the level before and after each pH determination. In each sample

pH was measured in duplicate; the standard deviation was 0.006 pH unit. Plasma c^q was measured in each sample in duplicate using a cediometer; the coefficient of variation was

0.8 percent, pi^l was calculated as

^ ''CO2 " '^'^C02 mi = pH - log

^ . (2) ^ ^C02

For S the values given by Austin et al. [17] were used.

A total of 555 duplicate determinations of p^{ in plasma equilibrated at known tem- perature with gas mixtures of known F^^ were performed (table 2). The standard deviation

as calculated from the duplicate determinations was 0.008. The best approximation of the

relationship between p^{, T and pH was found to be [18].

p^i = .4.7416 . MMil- + 0.015906 T - log (l + ^^?^) • (3)

The mean difference between the experimental p^{ values and those from eq. (3) was 0.0000 ± 0.0138 (SD). Figure 1 shows a curve relating p^i to temperature as calculated with

eq. (3) at pH = 7.4. The curve is in good agreement with the experimental data presented,

as well as with those reported by other investigators. Figure 1 also demonstrates that the

estimation of the relationship between p^i and temperature on the basis of p^j determinations

at but two temperatures is subject to substantial error, as p^| does not vary linearly with

temperature.

The decrease of p^i with increasing pH reported by other investigators has been confirmed

in the present investigation and is represented by the term -log(l + 0. 020682/ lO'P^"*"^) in eq.

(3). The decrease of p^{ with increasing pH has been ascribed to a pH dependent error in the

determination of Cur-n.- [10], to an alkaline error of glass electrodes [9] or to both.

35
Table 2. Experimental and calculated p^l at different temperatures and pH.

Temp.

°C

pH range pH mean P^i (exp.) mean

PK{ (calc.) Difference SD of

difference

6. 8-7.3 7. 167 6.096 -0.,005 0.,0068

7 3-7 ^ 7 fi 0Q7 6. 089 +n , uuo nU . m nR

7 5-7 8 7, 542 6.083 6. 079 n nn'?n 37 5 6. 8-7.3 -77. 181 fiU a 113 1 1 0 6. 110 +0.,003 0,,0124

fn = 13fi^ 7 o— / . o 7 d?2 U a 1 U 1 6. 100 +n n

7 1-7 ft 7 6.090 6. 088 , uu^ n m 7ft , u 1 / 0

3R n 6. 8-7.3 7. 162 fi nn 6. 119 -0.,009 0,,0102

fn = 33^ 7 0— / • 3 7 HDD fi 10"^ 6. 107 -0.,004 u

,

, u 1 1

D

7/ a R-7 ft 7 6.087 6. 097 -0,,010 n , mU 1 Oo

3? R 6. 8-7. 3 7. 191 fi HQ 6. 127 -0.,008 0,.0166

fn = 3'i') 7 3-7 5 7 457 fi 113 6. 116 -0,,003 n , nnfiQ uuuy

7 5-7 8 7 fi33 6.103 6. 104 -0.,001 nu

.

. U 1 Cc. 30 n 6. O *7 O 8-7.3 "77. 149 fi 1d2 6. 138 +0.004 0,.0131

fn = 74^ 7 3-7 ^ 7 fi 1?Q 6. 128 +0.,001 n . U 1 OH

7 5-7 ft 7 6.118 6. 111 +0.,007 n

6, 8-7.3 7. 131 fi 1 Rfi 6. 156 0,.0114

fn = 7 0— / • 0 7 fi Ml 6. 145 -0,,004 nu

.

nnoft

7 "1-7 ft 7 6.128 6. 132 -0.,004 n nn 2

U • o / » o 7 1 C.\J fi IQl 6. 188 +0,.003 n

fn = Q?^ 7. 3-7.5 7. 387 fi 17Q 6. 178 +0,,001 0,.0126

7. 5-7.8 7. 616 6.162 6. 164 -0.,002 0,.0127

16.0 6. 8-7.3 7. 050 6.212 6. 212 0,.0167

(n = 85) 7. 3-7.5 7. 382 6.202 6. 200 +0.,002 0,.0139

7. 5-7.8 7. 561 6.185 6. 190 -0,,005 0,.0164

Obviously, as o^qq- is calculated as (c^q - S-P^^ ) or determined by titration, C0|- and

RNHCOO" present in plasma are included. The value used for Cj^^q- in calculating p^j will be too high and pk'i consequently lower than pK[. As the fraction of bound CO2 present as COl"

or as RNHCOO" increases with increasing pH, the error increases with increasing pH. However,

g^q2- and ^RNHCOO- ai^e too small even at higher pH to completely explain the observed

decrease of pk'i with increasing pH. An alkaline error of the glass electrode might in part

be responsible.

The scatter of the pK'i values of table 1 and of the present investigation is much

smaller than that reported for pK'i in disease when measuring Pp-, and pH in whole blood,

using membrane covered electrodes for the determination of Pm • Some changes of pZi

(and thus of pKi) may be expected to result from changes in the ionic strength of plasma.

However, the changes of the ionic strength of plasma in disease are too small to be

responsible. Because of the unequal distribution of dissolved CO2 and HCO3 over the water

and non-water phase of plasma, the value of pK'i found in plasma differs slightly from that

in simple electrolyte solutions of the same ionic strength. Changes of the distribution

ratio of CO2 and HCO3 during disease might cause a change of pK'i. However, these changes

have been calculated to be small [19]. Variations of pK[ in disease must therefore largely

be ascribed to unknown errors in the determination of pH, Pp^, or Curn' abnormal plasma.

36
_J I I

' ^0 35 30 25 20

Figure 1. Relationship between the experimentally determined apparent first dissociation

constant of carbonic acid in human plasma and temperature. The curve represents eq. (3)

at pH = 7.4

• present investigation o Warburg [2]

e Cullen et at. [5] x Severinghaus et at. [9] + Siggaard-Andersen [10] a Maas [12]

The pifj values taken from the literature have been recalculated, using for s the

values reported by Austin et al. [17].

References

[1] Hasselbalch, K. A., Die Berechnung der Wasserstoffzahl des Blutes aus der freien und

gebundenen Kohlensaure desselben, und die Sauerstoffbindung des Blutes als Funktion

der Wasserstoffzahl, Biochem. Z. 78, 112 (1916).

[2] Warburg, E. J., Carbonic acid compounds and hydrogen ion activities in blood and salt

solutions, Bioohem. J. 16, 153 (1922).

[3] Cullen, G. E., Studies of acidosis. XIX. The colorimetric determination of the

hydrogen ion concentration of blood plasma, J. Biol. Chem. 52, 501 (1922).

[4] Van Slyke, D. D., Hastings, A. B., Murray, C. D., and Sendroy, J., Studies of gas and

electrolyte equilibria in blood. VIII. The distribution of hydrogen, chloride and'

bicarbonate ions in oxygenated and reduced blood, J. Biol. Chem. 65, 301 (1925).

[5] Cullen, G. E., Keeler, H. R., and Robinson, H. W., The pZ' of the Henderson-Hassel balch

equation for hydrogen concentration of serum, J. Biol. Chem. 66, 301 (1925).

[6] Hastings, A. B., Sendroy, J., and Van Slyke, D. D., Studies of gas and electrolyte

equilibria in blood. XII. The value of \>k[ in the Henderson-Hassel balch equation

for blood serum, J. Biol. Chem. 79, 183 (1928).

[7] Robinson, H. W., Price, J. W., and Cullen, G. E., The value of pZ' in the HendersonHasselbalch equation for human and dog sera, determined with the Simms electrode, J. Biol. Chem. 106, 7 (1934).

37
8] Dill, D. B., Daly, C, and Forbes, W. B., The pK{ of serum and red cells, J. Biol.

Chem. m, 569 (1937).

9] Severinghaus, J. W., Stupfel , M., and Bradley, A. F., Variations of serum carbonic

acid pz' with pH and temperature, J. Appl. Physiol. 9^, 197 (1956).

10] Siggaard-Andersen, 0., The first dissociation exponent of carbonic acid as a function

of pH, Sound. J. Clin. Lab. Invest. T4, 587 (1962).

11] Albers, C, Kappey, F., and Thiele, P., Die Elektrolytbrilcke als Fehlerquelle bei der

pH-Messung, Klin. Wsahr. 41_, 1095 (1965).

12] Maas, A. H. J., Van Heyst, A. N. P., and Visser, B. F., The determination of the true

equilibrium constant (p^ig) and the practical equilibrium coefficient (pi^jg) for the

first ionization of carbonic acid in solutions of sodium bicarbonate, cereorospinal

fluid, plasma and serum at 25 and 38 °C, Clin. Chim. Acta, 33, 325 (1971).

13] Trenchard, D., Noble, M. I. M,, and Guz, A., Serum carbonic acid pii:{ abnormalities in patients with acid-base disturbances, Clin. Soi. 32^, 189 (1967).

14] Sinclair, M. J., Hart, R. A., Pope, H. M., and Campbell, E. J, M., The use of the

Henderson-Hassel balch equation in routine medical practice, Clin. Chim. Aota, 19, 63

(1968).

15] Rispens, P., Brunsting, J. R., Zijlstra, W. G., and Van Kampen, E. J., Determination of

total carbon dioxide in blood and plasma by means of the cediometer. Theory and

experimental verification, Clin. Chim. Acta, 22^, 261 (1968).

,16] Rispens, P., Signifiaanae of plasma bicarbonate for the evaluation of homeostasis,

p. 97 (Van Gorcum and Comp. , Assen, The Netherlands, 1970).

17] Austin, H. W., Lacombe, E., Rand, P. W., and Chatterjee, M., Solubility of carbon

dioxide in serum from 15-38 °C, J. Appl. Physiol. ]S_, 301 (1963).

:i8] Rispens, P., Dellebarre, C. W., Eleveld, D., Helder, W., and Zijlstra, W. G., The

apparent first dissociation constant of carbonic acid in plasma between 16 and 42.5 °C,

Clin. Chim. Acta, 22, 627 (1968).

19] Austin, H. W., Ferrante, V., and Anderson, C, Evaluation of whole blood pz' in the

acutely ill patient, J. Lab. Clin. Med. 11, 129 (1968).

38
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975; Issued June 1977.

QUANTITATIVE RELATIONSHIPS BETWEEN TOTAL CO2 CONCENTRATION

IN BLOOD AND PLASMA, PLASMA BICARBONATE CONCENTRATION,

PLASMA pH AND CARBON DIOXIDE TENSION BETWEEN 16-42 °C

P. Rispens, J. P. Zock, and W. G. Zijistra

Laboratory of Chemical Physiology

University of Groningen

Groningen, The Netherlands

In our laboratory the acid-base status of patients or experimental animals is evaluated

by determining total CO2 concentration in blood {oqq^), blood pH at 37 °C (pH^(37)),

haemoglobin concentration and oxygen saturation (502)- Fi^om these quantities plasma

pH at 37 °C (pHp(37)), total CO2 concentration in plasma i^Q^) > plasma bicarbonate concentration (^QQ") si^d blood pH and carbon dioxide tension at body temperature (pH^(r^)) and

^CO ^^Zj^^

calculated [1]^.

pH (37) is calculated from pH,(37) using eq. (1)

pHp = pH^ + 0.01 . (1)

c^Q and (^QQ- ai^e practically independent of temperature in anaerobical ly stored blood and

can be calculated from the measured quantities without considering body temperature T^. c^Q is calculated from c^q using eq. (2)

("2.244 - 0.422 5q )

• (8.74 - pH^(37))

aP = - - • (2) CO2 CO2

^2.244 - 0.422 S^J • (8.74 - pH^(37)) -0.02880^^^

e^^Q- can be calculated from c^q^ and pHp(37) using the Henderson-Hasselbalch equation,

which for this purpose may be written as

402

^C03 =

1 , ^oP^l - pHp(37)

The value to be used for p^j at 37 °C and pH (37) is calculated using eq. (4)

P

p^l = -4.7416 + IM0J4i ^ 0.015906 T - log (1 + ^^_pH°(gf+ 7 )(^ °K). (4)

I ^Figures in brackets indicate the literature references at the end of this paper.

I

! 39
In contrast to c^q^ and

-^^QOg' ^CO

s^'"°"9^.y depend upon temperature. PH^C^^) is

calculated using eq. (5)

P%iT^) = pH^C37) + C37 - T^] -(o.OlSl + 0.0058 (pHp(37) - 7.4)) (r^in °C). (5)

Pqq^C^^) is calculated from PH^C^^) and c^q using the Henderson-Hasselbalch equation,

which for this purpose is written as

"XOz

Ppn (r, ) = . (6)

S

(

pH (r ) - p^l \

'

. \10 P ^ +1/

The value to be used for S at is calculated from eq. (7), which has been derived from

data of Austin et at. [2]

S = - 3.24615 + ^-^y^^ + 0.0049854 T [t in °K) . (7)

The value to be used for pH^Cr^) is calculated from pH^(r^) using eq. (1), the value to be

used for p^i at and pH^Cr^) is calculated using eq. (4).

Using the eqs. (1) through (7), a computer program can be made for the rapid calculation

of tJ^QQ- and pH^Cr^) and P^q (r^) at any temperature between 16 and 42 °C. When no computer

is available, the nomogram presented in figure 1 may be used [3]. This nomogram is an extension of the nomogram described by Brunsting [4] which is, in turn, a modification of the

Singer-Hastings nomogram [5]. The nomogram is valid for oxygenated blood only. Correction

for desaturation is possible with the aid of the data of table 1.

Obviously, eqs. (1) through (7) can also be used to calculate pH,{2',), ^cOo^^;,)' ^HCOq»

eQ02 ^""^ ^COa PH^(37), Pco2(37), o^^ and Sq^ have been measured or to calculate pH^(37),

^h^h^^ ^Z^M^' ^feo; ^?0p' w^^" 4op' ^C0.(^7)' ^Hb ^02 ^^^^ '^^^^ measured.

1 . Calculation of pH . P

It has been found that there is a slight difference between pH^ and pH^ [6]. This

difference is probably due to the effect of haemolysis at the interface between saturated KCl and blood on the liquid junction potential during the measurement of blood pH [7-9].

2. Calculation of Cqq^-

Within the pH range occurring in blood in vivo, c^q in red cells (cqo2^ ^'^ smaller than

c^Q . The difference is mainly due to the impermeability of the red cell membrane for

haemoglobin, organic phosphate and plasma protein ions, causing a Gibbs-Donnan distribution

40
c"^.. (mmol/l)

^ 30

n40

PHb(37)

shift in point of

ratation in relaP^Q^(mm Hg)

-10

20T(^{°C)

cPHC03-(meq/1)

20

-30

^0

-50

60

70

H80

90

100

110

1 20

(g/lOO ml)

Figure 1 Nomogram to derive c^q^> ^COs'

'CO2 ^Hb h- To derive

) XO2 temperature {p^^ {T-j^)) and pH^ at body temperature (pH^(r^)) from pH^ at 37 °C (pHj^(37)),

^Q^and

of pH,

^j,Q- a line is drawn through the points corresponding to the measured values

(37), c^Q^ an^l a^^. a^^^ is then read at the point where the c^q scale for

= 0 is intersected, and <2^qq- at the point where the <^qq- scale is intersected.

(37)) is read at the point where the p^q scale is intersected. To

find P^Q^ {T^) when differs from 37 °C, a line has to be drawn through the point

(at the e^Q^ scale for a^^ = 0) and another

''Hb ^C02

37

(r, ) when T. corresponding to the value found for c?^^ (at the e; point at the central Hj^ (37) scale. The latter point may be found using the lines

representing the shift of this point with temperature for a given value of (37) in the

right hand grid. In the right hand grid the line through the measured value of pH^(37).

is followed up (or down) to the point of intersection with the line corresponding to the

actual value of T^. The point of intersection is then projected on the central pHj(37)

scale. The lines of the left hand grid represent the change of pH^ with temperature

for a given value of pH^(37).

41
Table 1. Correction factors for cP^q- and p^q^ for deoxygenated blood.

"^Hb g/1

b ^

' 6.8 7.0 7.2 7.4 7.6 7.8

20 0.3 0.3 0.4 0.5 0.5 0.7

40 0.6 0.7 0.8 0.9 1.1 1.3

60 1.0 1.1 1.2 1.4 1.7 2.1 80 1.3 1.5 1.7 1.9 2.3 2.8

100 1.6 1.8 2.1 2.5 2.9 3.7

120 2.0 2.2 2.6 3.0 3.6 4.5

140 2.3 2.7 3.1 3.6 4.3 5.5

160 2.7 3.1 3.6 4.2 5.1 6.5

180 3.1 3.5 4.1 4.8 5.9 7.5

200 3.5 4.0 4.6 5.5 6.7 8.7

The figures given represent the percentage to be subtracted from the value read

for ^QQ- and for p^q in the nomogram of figure 1, when the blood is completely

deoxygenated. When the blood is partly oxygenated the correction factor is found

by multiplication of the figures given with (1 - S^. ).

of the diffusible ions, amongst them HCO3. a given haemoglobin concentration within the

red cells (MCHC = a'^^) the Gibbs-Donnan ratio mainly depends upon pH^ and 5q [10], whereas

at given a^^^ and ^qq^> c^q^ depends upon the haematocrit H,

Thus, the ratio c^q

/'^c02 ^^P®"'^^ ^P°" ^ ^""^ '^Hb furthermore upon pH^ and .

Therefore, it seems preferable to base empirical equations for calculating o^q from c^q

upon c^j^ (= a'^^ • h) , pH^ and 5q . Equation (2) was derived by McHardy [11] from an equation

given by Visser [12] who derived his equation from a nomogram given by Van Slyke and Sendroy

[13]. The validity of eq, (2) was tested [9] using 317 blood samples in which o^q , c^q , pH^(37), c^^ and Sg were measured. Of the 317 blood samples, 83 were centrifuged at 37 °C,

the remaining 234 at room temperature. Results are shown in table 2. The mean difference

between c^q /c^q calculated using eq. (2) and c^q /c^q measured directly is but small. The

Table 2. Difference between o^q /c^q derived from pH^(37), and a^^

(or H or O2 capacity), and '^q^/(^^q^ calculated directly.

Centrifuged Mean , at (°C) difference SD° Remarks

37

19-25

19-25 and 37

83

234

317

number of experiments

standard deviation

+0.007

+0.014

+0.012

0.027

0.034

0.032

equation (2)

equation (2)

equation (2)

42
mean difference between calculated and measured values of c^q /c^q of blood samples centrifuged

at room temperature is slightly greater than that found for samples centrifuged at 37 °C.

h

However, eq. (2) may be used to calculate c^q from c^^ at lower as well as at = 37 °C

without introducing an appreciable systematic error. Figure 2 illustrates that eq. (2) is

valid over the entire c^q range encountered in blood in vivo.

Figure 2. Relationship between ct^q^ calculated from c^q^. pHj,(37), o^^ and s^^ with

eq. (2) ^c?^Q^(calc)j and c^q^ measured directly (^q^

(exp)^

3. Calculation of "^^qqIn contrast to eq. (6) (y.i. ) an overall check of the validity of eq. (3) is not possible, because c^COs ^^"^"^^^ measured independently of c^q or pH^. However, because

the denominator of the quotient in eq. (3) differs but slightly from unity (in the pH range

7.8-7.0, it ranges from 1.020 to 1.126), the inaccuracy in p!^-] and pH^ has a negligible

effect on c^^q- and the accuracy with which c^rn: determined depends upon that of

only.

HCOc 'CO.

4. Calculation of pW^{Tj^)

.

pH^ of anaerobically stored blood increases with decreasing temperature, because the

acid reaction of both carbonic acid, haemoglobin and plasma decreases at decreasing temper43
ature. Theoretically [9], ApH^/AT depends upon pH^(37), c^q and a^^ • 5q has only a minor

influence. Equation (5) was derived from experiments with bovine blood, in which the influ- ence of c^q and o^^ proved negligible. The validity of eq. (5) for human and canine blood

was checked by measuring pH^(37) and pH^(30) of 336 human and 829 canine blood samples. In addition, pH^(37) and pH^(20) was measured in 52 human blood samples and pH^(37) and pH^(16)

in 68 human blood samples. Results are shown in table 3. An excellent agreement between

Table 3. Mean difference and standard deviation between pH^(30), pH^(20) and pH^(16)

calculated from pH,(37) and pH,(30), pH,(20) and pH,(16) measured directly.

Number of Temperature

Blood measurements (°C) pH, (meas)-pH, (calc) SD

canine 829 30 +0.001 0.016

human 336 30 -0.003 0.015

human 58 20 -0.031 0.026

human 68 16 -0.040 0.015

calculated and measured pH^(30) was found both in human and canine blood. At 20 °C and 16 °C

the calculated values are somewhat lower than the measured ones. No explanation for this

discrepancy can be given. It is in contrast with the good agreement of P^^ at 16 °C calcub

2 lated from pH^(37), a^^ , e^|^ and 5q with P^^ in the gas mixture with which the blood had

been equilibrated at 16 °C {v.i.).

5. Calculation of p^q^{Tj^) .

It is clear from eq. (6), that the accuracy with which P-^ (^t,) can be calculated from

b

2 " r, pH^(37), Cj,Q , a^^ and 5q , depends upon the reliability with which c^q is calculated from

e^Q (eq. (2)) and P^^iTj^) is calculated from pH^(37) (eqs. (1) and (5)), as well as on the

validity of eqs. (4) and (7) for calculating pk^ and 5. The overall accuracy of the determination of P^Q (t^) was evaluated by equilibrating human blood with CO2/O2/N2 mixtures of known

Table 4, Mean difference and standard deviation of Pm calculated at tonometer

temperature from pHj(37), c^q^, c^^ and , and P^q of tonometry gas mixture.

^C02

""^"9^ (mm Hg) 37 "C 30 °C 16 °C

10-30 +1.2 ± 0.9 +0.,6 + 1.2 +0.,6 + 1.1

n = 57 n 38 n 15 30-50 +0.8 ± 1.7 +0..1 + 2.5 +0.,8 + 2.3

n = 48 n 17 n 15

50-80 -0.3 ± 2.0 -0,,9 + 2.3 +0.,6 + 4.3

n = 61 n 36 n 14

80-120 -1.9 ± 3.7 -4,.4 + 4.4

n = 27 n 13

44
1- composition, measuring c^q , pH^(37), a^^ and , calculating p^q^{t^) from the measured

values and comparing the calculated P^q (t^) with P^q in the gas mixture used. One hundred

ninety three blood samples were equilibrated at 37 °C, 104 at 30 °C and 44 at 16 °C. Results

are shown in table 4. At P-^ = 10-80 mm Hg a fair agreement is found. The lower values

for blood with Pm > 80 mm Hg, most evident in samples equilibrated at 30 °C, may be due to

loss of CO2 in the sampling or analysis procedure.

References

[1] Rispens, P. and Zijlstra, W. G., Blood gas transport and acid-base balance, in

Clinical Biochemistry^ Principles and Methods , H. CH. Curtius and M. Roth, eds., p. 1604 (Walter de Gruyter, Berlin, New York, 1974).

[2] Austin, H. W., Lacombe, E., Rand, P. W., and Chatterjee, M., Solubility of carbon

dioxide in serum from 15-38 °C, J. Appl. Physiol. 18, 301 (1963).

[3] Rispens, P., Brunsting, J. R., Zock, J. P., and Zijlstra, W. G., A modified SingerHastings nomogram, J. Appl. Physiol. 34, 377 (1973).

[4] Brunsting, J. R., Clinical Caj>hoxymetry, p. 95 (Van Gorcum and Comp., Assen, The Netherlands, 1962).

[5] Singer, R. B. and Hastings, A. B., An improved clinical method for the estimation of

the acid-base balance of human blood. Medicine, 2J_, 223 (1948).

[6] Severinghaus, J. W., Stupfel , M., and Bradley, A. F., Accuracy of blood pH and PCO2

determinations, J. Appl. Physiol. 9^, 189 (1956).

[7] Siggaard-Andersen, 0., Factors affecting the liquid-junction potential in electrometric

blood pH measurement, Scand. J. Clin. Lab. Invest. J_3, 205 (1961 ).

[8] Maas, A. H. J., pH determination of body fluids with a micro glass electrode and a saturated KCl bridge in the cell, Clin. Chim. Acta, 28, 373 (1970).

[9] Rispens, P., Significance of plasma bicarbonate for the evaluation of homeostasis^

p. 116 (Van Gorcum and Comp., Assen, The Netherlands, 1970).

[10] Hastings, A. B., Sendroy, J., Mclntosch, J. F., and Van ,Slyke, D. D., Studies of gas

and electrolyte equilibria in blood. XIII. The distribution of chloride and bicarbonate

in the blood of normal and pathological human subjects, J. Biol. Chem. 79^, 193 (1928).

[n] McHardy, G. J. R., The relationship between the differences in pressure and content of

carbon dioxide in arterial and venous blood, Clin. Sai. 32, 299 (1967).

[12] Visser, B. F., Pulmonary diffusion of carbon dioxide, Phys. Med. Biol. b_, 155 (1961 ).

[13] Van Slyke, D. D. and Sendroy, J., Line charts for graphic calculations by the HendersonHasselbalch equation and for calculating carbon dioxide content from whole blood content,

J. Biol. Chem. 79, 781 (1928).

45
I

I

I

i
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

Pf.f. INDEPENDENT QUANTITIES, WHY OR WHY NOT

P. Rispens, W. G. Zijistra, and E. J. van Kampen

Laboratory of Chemical Physiology

University of Groningen

Groningen, The Netherlands

and

Laboratory of Clinical Chemistry

Diakonessenhui s Groningen, The Netherlands

In our laboratories, disturbances of acid-base balance are assessed by plotting plasma

pH at 37 °C (pH(37)) against plasma bicarbonate concentration (cSrr>-) "in a pH -cf,rnp nLUq p nUU3

diagram (fig. 1). Such a diagram, first introduced by Van Slyke [Ip and popularized by

comp. comp. prim.

n.r.oc. rolk? Figure 1. A pH (37)-t?2pr,- diagram with normal area and the area in which pH and may

P nl^U3 r p HCO3

be expected to fall during primary and compensated respiratory and non-respiratory disturbances and during mixed disturbances of acid-base balance; r.ac. = respiratory acidosis;

r.alk. = respiratory alkalosis; n.r.ac. = non-respiratory acidosis; n.r.alk. = non-respiratory

al kalosis.

^Figures in brackets indicate literature references at the end of this paper.

47
Davenport [2] is denoted a Van Slyke-Davenport diagram. In this diagram, the areas in

which pHp and c^qq- may be expected to fall in normal conditions, primary respiratory and

non-respiratory disturbances, compensated respiratory and non-respiratory disturbances and

mixed disturbances are indicated. In addition, the diagram contains CO2 isobars and CO2

buffer lines. A CO2 isobar is a set of points with equal P^q ; a CO2 buffer line is a graph showing the change of pH^ and <^^qq- upon changes of P^q . The basic ideas underlying

the use of the diagram can be found elsewhere [1-3].

The diagram demonstrates that the diagnosis of an acid-base disturbance cannot be

based on either pH^ or c^qq- (or Pqq^) alone. As Van Slyke put in 1921 [1]:

In order to determine which one of the possible variations

exists in the blood in vivo, it is necessary to ascertain two of the involved variables, such as the pH, [BHCO3], and

[H2CO3]. With any of two of them a point can be located in

its proper area on a diagram but with any one of them

alone it cannot be done.

The diagram obviates the use of so-called P^q independent quantities (table 1) for

the diagnosis of acid-base disturbances. For the following reason, these quantities are nevertheless widely in use. A respiratory disturbance of acid-base balance or a ventilatory

compensation of a non-respiratory disturbance manifests itself in a change of P^q . A non-respiratory disturbance of acid-base balance or a metabolic or renal compensation of a respiratory disturbance manifests itself in a change of c^qq-- However, whereas a res- piratory disturbance and a ventilatory compensation of a non-respiratory disturbance can be

assessed by comparing the actual P^q with normal, non-respiratory disturbances, metabolic

and renal compensations of respiratory disturbances cannot be assessed by simply comparing

the actual ^qq- with normal. This is because changes of P^q also cause changes of

c^^Q-. Hence, a deviation of <?^qq- from normal may be due either to a non-respiratory

disturbance or metabolic or renal compensation or to a respiratory disturbance or ven- tilatory compensation. The basic idea underlying the so-called P^q -independent quantities

is that the respiratory deviation of ^qq- is ruled out by restoring a normal P^q^ in blood

or plasma in vitro. Figure 2 illustrates that this basic idea is not true. The course of

pH with cSrn- when P-^, is restored to normal in blood or plasma in vitro differs from

^ p HCO3 CO2

that in vivo, i.e., the course of the buffer line of isolated plasma and of blood in vitro

both differ from that of blood in vivo. In the case of the primary respiratory acidosis of

figure 2, standard bicarbonate is lower than normal. This would suggest a non-respiratory

acidosis (complicating a prevailing respiratory acidosis). The same conclusion would be

drawn if base excess or buffer base values were calculated from pH^, c^^Qg ^C02*

Alkali reserve is higher than normal in the case of figure 2. This would suggest a metabolic or renal compensation of the respiratory acidosis.

It should be emphasized that neither the actual ^j^qq- nor the alleged P^Q^-independent

quantities are an unequivocal measure for non-respiratory disturbances from which the

amount of acid or base to be administered for restoring pH can be calculated [4,5]. This

y

48
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49
Fc02 (fn^Hg) KO 120 100 80

Figure 2. Alkali reserve and standard hicarhonate during a severe primary respiratory

acidosis. The alkali reserve is seen to be higher than normal, the standard bicarbonate

to be lower than normal.

is because changes of ^^^qq- (and thus, changes of P^^ -independent quantities which are in some way related to or derived from c^qq- (table 1)) are but a distant reflection of the

non-volatile acids or bases actually gained or lost. In contrast to a respiratory disturbance where the change of P^^ is a direct measure of the acid gained- or lost, a change

of t?^QQ- may indicate gain or loss of organic acids in metabolic disorders, gain or loss of

inorganic acids or bicarbonate in renal disorders and gain or loss of bicarbonate or hydrochloric acid in gastro-intestinal disorders. It has been found, that the same pH^ and

c^QQ- may occur while different amounts of non-volatile acids or bases have been gained or lost [6].

It should be also emphasized that neither the use of a pH -c^ro- diagram nor the use of

-independent quantities releases the physician who has to treat a patient with an acid-base disturbance, from the obligation to go deeply into the pathophysiology of this

disturbance. The following case taken from reference 7 serves to illustrate this. It concerns a 45-year old female patient suffering from chronic obstructive pulmonary disease,

admitted to the hospital during an acute exacerbation. On analysis of an arterial blood

sample, pH^ = 7.17, -^^^Qg ^ meq-r\ P^q^ = 92 mm Hg and S^^ = 59 percent were measured.

Treatment consisted of mechanical ventilation to bring p'^q^ back to normal. A second

arterial blood sample then yielded pH = 7.59, c^rr,- = 34 meq-l'^, P-^. = 36 mm Hg and

Sf. = 96 percent. From acidemic, the patient had become alkalemic; the treatment thus

50

ll
resulted in a serious condition leading to severe neurologic disfunction. Such mismanagement may easily result from the uncritical application of either a PH^-^^^qq- diagram or of

P^Q -independent quantities. Considering the patient's history, it may be assumed that

before the acute exacerbation the patient had a compensated respiratory acidosis with high

Pqq , high c^QQ- and a nearly normal pH^, and with an arterial Pq still high enough to ensure an arterial S~, of about 90 percent. During the acute exacerbation P-^ increased

U2 LU2

and Pq decreased. The increase of Pqq would in itself cause an increase of <^qq- (and a decrease of pH^). However, Pq^ then decreased to values corresponding to the steep part of

the O2 dissociation curve and 5q decreased considerably. 'Hence, the patient became hypoxic,

lactic acid accumulated, pH^ decreased further and c^qq- decreased to a value usually found

during a primary respiratory acidosis. However, in this patient with chronic respiratory

disease a much higher (^qq- was to be expected if her condition had not been complicated by

a non-respiratory disturbance. When hypoxia was removed by mechanical ventilation, the

accumulated lactic acid was metabolized and (^qq- increased. As P^q decreased far below

the value existing before the acute exacerbation, pH^ increased to a high value.

If the patient's history had been considered with enough insight in the pathophysiology

of acid-base disturbances, the mechanical ventilation would have been so regulated that a nearly normal pH^ would have been obtained instead of a normal Pqq^. For such an insight,

the alleged Pqq^- independent quantities are not required.

References

[1] Van Slyke, D. D., Studies of acidosis. XVII. The normal and abnormal variations in

the acid-base balance of the blood, J. Biol. Chem. , 48, 153 (1921).

[2] Davenport, H. W., The ABC of Aaid-Base Chemistry (The University of Chicago Press,

Chicago, 111., 1958).

[3] Rispens, P., Zijlstra, W. G., and Van Kampen, E. J., Significance of bicarbonate for

the evaluation of non-respiratory disturbance of acid-base balance, Clin. Chim. Aata,

54, 335 (1974).

[4] Van Slyke, D. D., Discussion remark, in Current concept of acid-base measurement, Ann.

N.Y. Aaad. Soi. , 133, 107 (1966).

[5] Garella, S., Dana, C. L., and Chazan, J. A., Severity of metabolic acidosis as a determinant of bicarbonate requirements. New Eng. J. Med., 289, 121 (1973).

[6] Lemann, J., Lennon, E. J., Goodman, A. D., Litzow, J. R., and Relman, A. S., The net

balance of acids in subjects given large loads of acid or alkali, J. Clin. Invest.,

44, 507 (1965).

[7] Rotherann, E. B., Jr., Safar, P., and Robin, E. D., CNS disorder during mechanical

ventilation in chronic pulmonary disease, JAMA, 189 , 992 (1964).

[8] Van Slyke, D. D. and Cullen, G. E., Studies of acidosis. I. The bicarbonate con- centration of the blood plasma; its significance and its determination as a measure of

acidosis, J. Biol. Chem., 30, 289 (1917).

[9] Singer, R. B. and Hastings, A. B., An improved clinical method for the estimation of

disturbances of the acid-base balance of human blood. Medicine, 27, 223 (1948).

51
[10] J0rgensen, K. and Astrup, P., Standard bicarbonate, its clinical significance and a new method for its determination, Scand. J. Clin. Lab. Invest., 9^, 122 (1957).

[11] Siggaard-Andersen, 0. and Engel , K. , An improved method for the calculation of the

relevant blood acid-base data, Soand. J. Clin. Lab. Invest., 1_2, 177 (1960).

[12] Brunsting, J. R., Clinical Cajcboxymetry, p. 95 (Van Gorcum and Comp. , Assen, The

Netherlands, 1962).

52
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

BASE EXCESS. WHY REOPEN THE ACID-BASE DEBATE?

P. J. N. Howorth

Department of Chemical Pathology

King's College Hospital Medical School

London, U.K. SE5 8RX

Base excess (BE) [1]^ was introduced as a formal mathematical concept with the Astrup

system (AS) [2] in 1960, although the descriptive terms base excess and base deficit were

used by Peters and Van Slyke in the 1930' s [3]. By 1963, many criticisms had been leveled

against it. Astrup himself [4] proposed a conference to discuss the many problems in acidbase understanding. The outcome was the 1964 New York Academy of Science Conference whose

contents were published in 1966 [5]. We must take this conference as a baseline to see what

progress has been made in the last decade. The Scandinavians conceded that BE was not a completely objective description of the whole body's acid-base state. Astrup indeed never

claimed that it was; he held that it was an excellent way of describing changes in the

blood, not the whole body [4]. However, some of his supporters still tried to apply the

system to the whole body. In 1966, came the "corrected base excess" which was a formula to modify the BE in the light of knowledge about the way the in vivo CO2 titration curve differed

from the in vitro one [6]. This corrected base excess did not achieve wide popularity. As the decade passed the old arguments became half forgotten and peace reigned once more. In 1971, BE was modified to allow for buffering in EC fluid (SBE) and this was incorporated into

the Sfggaard-Andersen chart nomogram without too much fuss [7]. Then more recently, rumbles

of the old acid-base debate have been heard again. Why is this and what do the rumblers

hope to achieve?

Table 1. Recent milestones in acid-base history.

1954 Mark I Astrup equipment

1957 Standard bicarbonate

1960 Astrup trolley & system: base excess

1962 RIpH Moran Campbell

1963 Schwarz & Relman critique of Astrup system

1964 NYAS Conference: peace for a decade

1975/5 SI units. New analysers with old parameters

Table 2. Measurements of the non-respiratory component.

Plasma HCOi : Van Slyke 1917, 1931

Non-respiratory pH: Hasselbalch 1916

Whitehead 1964

Stoker 1972

Buffer base: Singer & Hastings 1948

Standard HCO3 : Astrup 1957

Base excess: Siggaard-Andersen 1960

Corrected BE: Prys Roberts 1966

Standard BE: Siggaard-Andersen 1971

^Figures in brackets indicate the literature references at the end of this paper.
The first and foremost point is that the instrument manufacturers have themselves

brought various skeletons out of the closet on their current blood gas analysers. They

appear increasingly reluctant to make simple machines that measure only pH, PO2 and PCO2.

What they like to sell at greater profitability are larger machines with accessory electronic circuits to calculate and print-out or display many of the old derived parameters

discussed at the 1964 Conference. The only modern parameter seems to the SBE on the Radiometer ABLl and of course this has been attacked on the familiar arguments that there is no way of knowing what the ECF buffers are in any particular patient, especially one with an acute on chronic acid-base disturbance.

Name:

Table 3. Parameter measured on some current instruments.

Blood Gas Analysers

Radiometer ABL Corning Eel 165 IL 413 Sandoz AVL 937

Measures

:

Calculates;

Hb cone.

P-HCOi

Std HCO3

total CO2

base excess

Std base excess

O2 saturation

All measure pH, PCO2 & PO2

P-HCO;

total CO2

base excess

P-HCO;

Std HCO3

total CO2

base excess

P-HCO;

total CO2

base excess

buffer base

O2 saturation

The second major reason for the re-opening of the debate is the growing use of SI units

in chemical pathology. This process is gaining momentum in Europe; for example, SI units

will be in widespread use in the UK by the end of 1975 [8]. With the introduction of the SI,

it is convenient to ask ourselves whether H+ concentration can be more conveniently reported

in an arithmetic SI unit, i.e., nanomole/1 iter , as compared with the logarithmic pH notation.

Of course these arguments are anything but new. It is worth pointing out that the pH notation was introduced by S0rensen as a shorthand for the otherwise unwieldy H+' concentration

scale [9]. Arguments about what electrodes actually measure, and activity versus concentration arose soon after. For many years H"*" activity was reported by the special symbol paH

and it was only after Peters and Van Slyke, in their classic work published in 1931, pronounced in favor of the use of pH to cover both H"^ activity and concentration that this

special symbol was finally abandoned [3]. In more recent years, many individual workers have

made use of the simpler arithmetic scale. One classic plea was Moran Campbell's RIpH in 1962 [10]. The upheaval caused by the introduction of the SI provides us with a very con- venient opportunity to make this change. The arguments will remain, but at least those who

wish to use the arithmetic SI unit will find it easier to get their colleagues to accept the

change. Already some gas analysers have partial pressures in kilopascals and no doubt soon we will be able to get an arithemtic [H"*"] output by means of an antilog circuit board as used with ion-selective electrodes for ions other than the H'''.

Further reasons for the fresh debate are a little more nonspecific. It is partly due

to a rejection of the rather Teutonic heavy-handed approach of the Astrup system with its

formulae and nomograms by a younger generation of clinicians. They are often unaware of the

emotions generated in the past by the debates and they prefer a more flexible approach to acid base assessment. Some points are:

A. A growing suspicion of nomograms because the historic need for them has largely

gone. They were needed in the past because laboratory instrumentation lagged behind theory

of acid-base understanding for many years. Thus, nomograms were introduced to extrapolate

the scanty laboratory data, e.g., to predict pH from CO2 measurements. With the availability

of ion-selective electrodes the need for nomograms has gone.

54
Table 4. Why there used to be a need for nomograms.

(1) Need to estimate blood H"*" concentration, e.g.,

from plasma CO2 and HCO3

(2) Assumed that physio-chemical constants apply to biological systems:

CO2 solubility coefficient = 0.510

PKl = 6.10 ± 0.01.

B. The modern approach is thus that laboratories should only report what they actually measure and leave it to the individual clinician to carry around the Radiometer blood gas calculator [11] if he wishes to derive parameters from the primary data. At a higher level

of sophistication we are of course still in trouble. A CO2 electrode does not really measure PCO2 but a pH change related, we hope, to it. What do osmometers really measure?

How are they calibrated?

Table 5. Philosophy of the modern approach.

(1) Clinical chemists should report only:

what is actually measured

what is usefully measured, e.g.,

blood [H+], PgCOa, actual HCO3.

(2) They should not report derived parameters,

e.g., PCO2 at 33° when measured at 38°.

(3) Use of nomograms should be restrained.

C. Thirdly, there is a greater interest in the interrelationships between O2 delivery

to the tissues and H"^ concentration. Hence, more measurements are needed in acid-base work

than the big three (pH, PO2, PCO2). When we talk of the P50 we enter another nomogram

area because it is still a relatively new topic. Since the O2 dissociation curve is pHdependent the P50 would be best measured at the in vivo pH. One group at NIH calculates

the so-called in vivo P50 from the standard pH 7.40 P50 using the MCHC, DP6, BE, pH and

temperature [12]. So far we have restricted ourselves to measure red cell DPG as ymol/g Hb

(not per liter red cells because the cells alter their shape in acidosis and alkalosis).

Table 6. Summary of current acid-base tests.

1974

pH

PCO2 mmHg

P-HCO3 meq/1

Std P-HCO3 meq/1

Base Excess Mkl-III

PO2 mmHg

1975

H^ cone, nmol/1

PCO2 kPa

P-HCO3 mmol/1

PO2 & P50 kPa

2:3-DPG

References

[1] Siggaard-Andersen, 0. and Engel , K. , A new acid-base nomogram. An improved method for

the calculation of the relevant blood acid-base data, Saand. J. Clin. Lab. invest.

12, 177 (1960).

55
[2] Astrup, P., J0rgensen, K. , Siggaard-Andersen, 0. and Engel , K. , The acid-base metabolism.

A new approach. Lancet, j_, 1035 (1960).

[3] Peters, J. P. and Van Slyke, D. D., Quantitative Clinical Chemistry, I. Interpretations,

(Bailliere, Tindall and Cox, London, 1931).

[4] Astrup, P., Acid-base disorders, New. Eng. J. Med. 269 , 817 (1963).

[5] Nahas, F. F., ed.. Current concepts of acid-base measurements, Ann. N.Y. Acad. Sci.

133 , 274 (1966).

[6] Prys-Roberts, C, Kelman, G. R. , and Nunn, J. F. , Determination of the in vivo CO2 titration curve of anesthetised man, Brit. J. Anaesth. 38, 500 (1966).

[7] Siggaard-Andersen, 0., An acid-base chart for arterial blood with normal and pathophysiological reference areas, Soand. J. Clin. Lab. Invest. 2J_, 239 (1971).

[8] Brit. Med. J. Editorial., S. I. units, Brit. Med. J. 4, 490 (1974).

[9] Lancet Committee report. Acid-base terminology. Lancet 11^, 1010 (1965).

[10] Campbell, E. J. M. , RIpH, Lancet L , 681 (1962).

[11] Severinghaus, J. W., Blood gas calculator, J. Appl. Fhys. 21_, 1108 (1966).

[12] Bellingham, A. J., Detter, J. C, and Lenfant, C, Regulatory mechanisms of haemoglobin

oxygen affinity in acidosis and aldalosis, J. Clin, invest. 50, 700 (1971).

56
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

USE OF m VIVO CO2 TITRATION CURVES IN THE PHYSIOLOGICAL ASSESSMENT

OF ACID-BASE BALANCE

P. J. N. Howorth

Department of Chemical Pathology

King's College Hospital Medical School

Denmark Hill

London, SE5, 8RX, United Kingdom

In their 1963 critique of the Astrup system, Schwartz and Relman [1]^ recommended

instead the use of a physiologic approach to the clinical assessment of acid-base balance in

individual patients. This approach was based upon the interpretation of clinical and laboratory data in the light of the known behavior patterns in respiratory and metabolic disease. This approach can be updated by making use of the in vivo CO2 titration curves in man. Assessment can then be subdivided into three broad areas:

1. Primary acid-base data

a. The respiratory component - arterial blood PCO2.

b. The acidity of blood--H^ nmol/1 (or pH) which gives direct information about the

accumulation of non-volatile acid or base and indirectly of changes in PCO2.

2. Other cl inicopathological information

a. Clinical data such as the duration and severity of the patient's present illness,

its antecedent pathological conditions (if any), and treatment which might modify

acid-base status.

b. Laboratory data such as plasma actual HCO3 and other appropriate measurement, e.g.,

arterial blood P02, P50 (partial pressure of O2 when Hb is 50 percent saturated

with O2), red cell DPG, K"*" and concentration of other electrolytes.

3. Knowledge of the following:

The usual compensatory changes to altered arterial blood H"*" and PCO2:

i. The ventilatory response to metabolic acidosis and alkalosis.

ii. The renal response to respiratory acidosis and alkalosis.

In the physiological assessment of acid-base balance, it is important to know how the ^ body reacts to acute and chronic changes in the two primary measurements, arterial blood [H ] and PCO2. The relative changes in [H^] and PCO2 are different in metabolic and respiratory

disturbances since the ventilatory response to metabolic acidosis and alkalosis is rapid

whereas the renal response to respiratory failure (acidosis) is slower, taking 3-4 days at least to become maximal.

1. The Ventilatory Response to Changes in [H^] in Metabolic

Acidosis and Alkalosis

When the [H'^] in the blood changes abruptly the PCO2 may not alter, but within a few

hours there is a ventilatory response to changes in blood [H ] which may take 24 hours or

Figures in brackets indicate the literature references at the end of this paper.

57
more for the compensatory change in PCO2 to be complete [2]. In compensated metabolic

acidosis, it is rare for the arterial blood Pcoa to fall below 2.0 kPa (15 mmHg) and the PCO2 never falls below 1.33 kPa (10 mmHg) as the work of breathing limits the maximum ventilation

volume.

In metabolic alkalosis, the suppression of normal ventilation is similarly a compensating mechanism. Observations of patients with arterial blood [H+] varying from 30 to 130 nmol/1 (pH 7.55-6.88) show that there is a hyperbolic relationship between the [H"*"] and

PCO2 [3,4].

2. The Renal Response to Changes in PCO2; Respiratory

Acidosis and Alkalosis

When the arterial blood PCO2 is acutely altered either by changes in ventilation or breathing artificially high CO2 mixtures, the resulting acid-base values are known as the

whole body C02-titration curves. They have been used for a number of years in the assessment

of acute respiratory acidosis [5]. The blood [H"*"] values normally found as the blood Pcoa changes from 2.0 to 12.0 kPa (15-90 mmHg) are shown in figure 1 [6,7]. The body reacts to a

• 90 12 80 1 1

. PCO2 / /

iO (l)Melabolic alkalosis

• 70

9 My (2)Melabolic acidosis

- 60 8

(3)Respiralory acidosi s 50 7

(4)Re5piratory alkalos: 6

s

• 40 5 30 4 20 32 (2)

- 10 1

1 1 1 1 1

t, 1 1 1 1 mmHg kPa 30 50 70 90 110 I30

Blood H^concenlralion nanomol / litre

Figure 1. Blood H concentration and PCO2 chart showing the range of observed values

during primary compensated metabolic alkalosis and acidosis and acute uncompensated

primary respiratory alkalosis and acidosis. During primary metabolic changes,

there is a superimposed respiratory response. (Data from Bone, et at. [3] and

Fulop, et aZ. [4] for the metabolic changes, and from Arbus, et al. 17} and

Brackett, et al. [6] for the respiratory changes.)

raised PCO2 by the renal retention of HCO3 ^""^ excretion of H"*". This process is significant

during the first 24-48 hours and becomes maximal after 3-4 days [8]. If the arterial blood

PCO2 remains below about 8 kPa (60 mmHg), the blood [H'''] may remain near normal, but renal

compensation begins to fail when the PCO2 rises above 9.5-10 kPa (71-75 mmHg) [9,10]. An acute rise in arterial blood PCO2 is invariably accompanied by a fall in arterial blood P02

which may be severe enough to cause lactic acidosis. In respiratory failure without assisted

ventilation, it is rare for the PCO2 to rise above 10.5-12 kPa (79-90 mmHg) as the associated

hypoxemia (P02 4-2.6 kPa) (30-20 mmHg) endangers life [11].

Hyperventilation may cause an acute fall in arterial blood PCO2, the plasma HCO3 "^sy

not change significantly, but the [H"*"] may become dangerously low. The changes found during

inhalation of low O2 gas mixtures and during adaptation to high altitudes are similar; the

renal adaptation to a low arterial blood PCO2 is slow and may take several weeks for com- pletion [12].

58
If a high arterial blood PCO2 is reduced by artificial ventilation, the restoration of

normal acid-base balance is slow because of the time taken for the renal compensating process

to work in the opposite direction. The kidneys may continue to excrete an acid urine for

some time, thereby producing a dangerous acid-base state indistinguishable from primary

uncompensated metabolic alkalosis [13].

3. The Whole-Body in vivo COa-Titration Curves

When patients with different initial levels of [H"*"] (pre-existing metabolic acidosis or alkalosis) are ventilated, the resulting in vivo C02-titration curves are almost parallel to the normal response. Hence, it has been suggested that the in vivo whole-body C02-titration

curves be used as an aid to the interpretation of acid-base problems [14-17]. Figure 2 shows

whole-body in vivo COa-titration figures from the work of Stoker and his colleagues [14]

recalculated and plotted out in SI units [18].

(2) (1) 90

• 80

li -

10 • PC02(l)kPa / /

/

- 70 9

• (2)mmHg / V T /

60 87 50

6

- 40 5

- 30 4 -

3

- 20

2

-

- 10

1 -

10 20 30 40 50 60 70 BO 90 lOO

Blood H concenlralion nanomol/litre

Figure 2, Whole-body CO2 titration lines

in vivo in man at different values of

nonrespiratory H Concentration (blood

H+ at PC02 5.3 kPa (40 mmHg)). (Data

from Stoker, et al. [14].)

20 40 60 80

Blood H'^nmol /I

Figure 3. Whole body CO2 titration lines

in vivo in man as shown in figure 2 with

points omitted to show the effects on PCO2 of acute changes in H+ concentration (rise and fall in H"*") and on H"*" concentration of acute changes in PCO2

(rise or fal 1 in PCO2)

.

When a patient is ventilated or retains CO2 acutely, the points representing [H"*"] and

PCO2 move along a straight line parallel to those in figure 3 in which the individual points

have been omitted for clarity. An acute metabolic rise or fall in blood [H"*"] causes a horizontal movement to the right or left. Mixed and compensated acid-base disturbances show

movement resulting from both vectors.

4. Examples of the Use of the in vivo Chart

The use of the in vivo chart will be illustrated by some examples. It should be noted

that the chart is not a nomogram in that it is not used to calculate derived parameters. The

purpose of the chart is to simplify analysis of acid-base data by comparing the behavior of

the individual patient with the known behavior in metabolic and respiratory change.

59
A. Acute metabolic acidosis during chronic respiratory acidosis

Even apparently simple acid-base disturbances can have a "mixed" component. Figure 4 shows how an acute exacerbation of hypoxemia in chronic respiratory failure during which

period the blood PCO2 did not alter caused a "type A" lactic acidosis [19]. There was a horizontal "shift to the right" of H"*" and the HCOi fell from 38 to 30 mmol/1. This severe

20 40 60 80

Blood H'^nmol/l

Figure 4. Acute metabolic acidosis due to worsening of hypoxia during chronic

respiratory failure. The arterial blood P02 fell from 4.7 kPa (35 mmHg) at point (1) to 2.7 kPa (20 mmHg at point (2) and the bicarbonate fell from 38

mmol/1, point (1) to 30 mmol/1 at point (2) without change in PCO2 (9.3 kPa

(70 mmHg)). The primary acid-base results at point (2) could equally have

arisen from an acute respiratiry acidosis with early renal compensation in

a previously normal individual (point la to 2), or from an exacerbation of

CO2 retention in a patient with moderate chronic respiratory acidosis (point

lb to 2).

chronic respiratory acidosis, with maximal renal compensation, and the additional acute

metabolic acidosis cannot be distinguished from inspection of the blood acid-base results

from an acute severe respiratory acidosis with early renal compensation in a previously

normal individual (point la) or an acute exacerbation of CO2 retention in a patient with

moderate chronic respiratory acidosis (point lb).

B. Acute on chronic respiratory failure

Figure 5 shows the changes over 48 hours in a patient with an acute exacerbation of

chronic bronchitis. On admission (1), the arterial blood PCO2 was acutely elevated. Despite

antibiotics and controlled O2 administration, there was further CO2 retention. Intermittent

positive pressure respiration for 24 hours reduced the arterial blood PCO2 to 8.0 kPa (60

mmHg) which was similar to the patient's normal PCO2.

C. Acute respiratory acidosis with concealed metabolic alkalosis

A woman aged 60 years was admitted in acute respiratory failure; she was confused,

making a history impossible, restless and cyanosed. Intermittent positive pressure ven- tilation was instituted and after a few hours, ventilation results were at point (2) (see

fig. 6). At this point, it is now clear that the patient probably had a mixed disturbance

since, with the correction of the respiratory failure by intermittent positive pressure

60
J 1 L 20 40 60 80

Blood H'^nmol/I

Figure 5. Acute exacerbation of chronic bronchitis in a man aged 65 years. On

admission (1), the patient was hypoxemic, P02 4.0 kPa (30 mmHg), plasma HCO3

32 mmol/1; (2), further increase in CO2 retention during early treatment, plasma

HCO3 36 mmol/1. Intermittent positive-pressure ventilation was instituted

which rapidly lowered the PCO2 (3), plasma HCO3 38 mmol/1, but the patient

remained in compensated respiratory failure.

20 40 60 80

Blood H'*'nmol/l

Figure 6. Acute respiratory failure with concealed metabolic alkalosis in a woman aged 60 years. On admission, the patient was cyanosed and comatose;

P02 6.4 kPa (48 mmHg), plasma Na+ 137, r 5.8, CI" 83, HCO3 34, urea 3.3 mmol/1

(20 mg/dl). Intermittent positive pressure respiration quickly reduced the PCO2

and revealed the metabolic alkalosis (point 2) which was subsequently found to be

associated with long-term diuretic therapy for heart failure due to chronic

hypoxemia with bronchitis. At point (3), electrolyte results were plasma

Na+ 140, K+ 2.7, Ca2+ 1.85 (7.4 mg/dl), HCO3 45 mmol/1 and the patient began

to have tetanic convulsions due to the untreated metabolic alkalosis from which

she died 48 hours later.

61
respiration, her results at point (2) lie in the area of metabolic alkalosis shown previously. One can thus easily predict that more ventilation would only lower still further

the already dangerously low [H+], Unfortunately, ventilation was continued and the patient

died after 48 hours from tetanic convulsions. The full history later obtained from relatives

revealed that the patient was on long-term diuretic therapy for congestive cardiac failure

due to chronic hypoxemia with air way obstruction. The typical hypokalemia was concealed

initially by the respiratory acidemia. Thus, the patient had two separate reasons to develop a metabolic alkalosis: (a) urinary K+ loss due to diuretic therapy, and (b)

paradoxical aciduria after lowering the high PCO2 (see above).

In retrospect, it is evident that this patient's only chance was for energetic treat- ment of the metabolic alkalosis at point (2). Intravenous HCl is direct and logical [20].

If given stronger than 0.09 mol , it may cause hemolysis and intravenous sodium phosphate,

ammonium chloride, arginine hydrochloride, and even dialysis have been advocated.

There has been current interest in metabolic alkalosis because its incidence has been

reported to be rising and the body tolerates alkalosis much less well than acidosis [21].

Whereas patients have survived with blood [H+] 170 nmol/1 (pH 6.78), the mortality is high

in metabolic alkalosis with blood [H'''] 28 nmol/1 (pH 7.55) and survival is rare below blood

[H+] 23 nmol/1 (pH 7.65). Alkalosis may occur not only from potassium loss and in hepatic

failure, but also after surgery, e.g., after bypass operations, and in trauma, especially

associated with sepsis and peritonitis. The causes are varied and may include: loss of H"*" through gastric suction; over-use of alkaline liquids in early therapy, or urinary potassium

loss through adrenal cortical hormone activity [21].

Table 1. Metabolic alkalosis.

Mortality rises with fall in [H^]:

moderate < 28 nmol/1 pH 7.55

high H < 23 nmol .1 7.65

Occurrence:

Medical: drugs; e.g., diuretics

acute liver failure

chronic bronchitics on ventilation

Surgical: major trauma, by-pass operations

sepsis, hemorrhage, peritonitis

Aetiological Factors:

loss of gastric - suction

alkaline fluids - ACD blood, NaHCOg

paradoxical aciduria after fall in PCO2

K"*" loss in urine - adrenal hormones

Treatment

0.09 mol HCl in NaCl

iv NH4CI

iv Na phosphate, pH 7.0

iv arginine, HCl

In view of the large iatrogenic component and the graver risk to life in metabolic

alkalosis as compared with metabolic acidosis, it is clear that physicians and anesthetists

working in intensive care units need to exercise special care to prevent the unnecessary

development of metabolic alkalosis.

62
D. Compensated metabolic acidosis

Figure 7 shows the acid-base changes over 12 hours during the resuscitation of a patient with acute diabetic ketoacidosis. On admission (1), there was a marked metabolic

acidosis with compensatory low PCO2. Therapy with insulin, intravenous NaCl and potassium

supplements was enough to bring the patient to near normal acid-base balance. Intravenous

NaHC03 was not given.

20 40 60 80

Blood H^'ntnol/I

Figure 7. Acute diabetic ketoacidosis in a man aged 30. On admission, the plasma

glucose was 20 mmol/1 (360 mg/dl ) , plasma HCO3 less than 5 mmol/1. The

patient was resuscitated with intravenous NaCl with KCl supplements and insulin.

After 6 hours (2), the plasma HCO3 was 8 mmol/1. The patient had recovered

by 12 hours (3), although the acid-base state was not completely normal, plasma

HCO3 18 mmol/1.

E. Compensated metabolic alkalosis and acute respiratory acidosis

A further example of a complex acid-base disturbance was provided recently (figs. 8

and 9). A 45-year-old workman (stone mason) was admitted in a semi-comatose state, peripheral cyanosis, vomiting and tetany. From his past history of suspected alcoholic

excesses, he was thought initially to have possibly Wernicke or hepatic pre-coma. His

LFTs, blood EtOH and red cell TK did not support this view so he was urgently re-assessed.

Date 14 15 16 17 21 22 April 1975

H+

Na+

K+HCO3

Urea

PCO2

PO2

DPG

132

2.6

15.5

23

140

3.1 >50

27.4

10.0

36137

3.9

>40

36145

3.1

35 25.0 15.6

6.7' 6.1

37

138

3.4

354.5

6.9

35134

2.7

22

4.2

5.1

nmol/1

mmol/1

mmol/1

mmol/1

mmol/1

kPa

14th

15th

16th

A(l)

(2)

5.9 10.3 8.1 12.6 9.8 kPa

31.9 25.3 13.9 14.1 15.6 ymol/g Hb

semi-coma, cyanosis, ?li ver/EtOH

intubated and ventilated until 21st

pneumothorax:chest drained

chronic DU. Alkalosis from oral NaHCOs

RLLpneumonia

Figure 8. CI inico-pathological data (table)

on a man aged 45 with combined chronic

metabolic alkalosis due to ingestion of

NaHCOs and acute respiratory failure due

pneumonia and pneumothorax.

63
Figure 9. Primary acid-base results show ' days after admission in brackets. i

i

I 20 40 60 80

[H"^] nanomole/liire i

I

! A gross metabolic alkalosis, right lobar collapse and pneumothorax were found and he was i intubated and ventilated. The alkalosis was due to excess ingestion of NaHCOs for the pain

j of a chronic DU. The initial blood gas results are thus consistent with a profound com- ' pensated metabolic alkalosis and the high red cell DPG results support this. '

The acid-base disturbance was fairly rapidly restored to near-normality, but then his

acid-base point moved up and down the in vivo line a bit as is often found in patients on ventilation. There was thus a chronic condition with an acute respiratory episode which

led to the patient's hospitalization.

F. Combined metabolic and respiratory acidosis

In cardiac arrest, there is a combination of acute respiratory acidosis and hypoxic

metabolic acidosis. Figure 10 shows the sequence of events in the attempted resuscitation

of a man aged 55 who arrested after a myocardial infarct. Blood taken 15 minutes after the

arrest showed a profound combined acidosis. The PCO2 was reduced by ventilation (2) and

the metabolic component was improved by 100 mmol of intravenous NaHCOs (3). At this point,

the airway became obstructed and after fruitless attempts to improve ventilation the

patient died (5) about 2 hours after the arrest.

It is reasonable to consider reducing a high metabolic [H^] in life-threatening

situations by administering NaHCOs, although it is likely that death in severe acidosis is due to hypoxic damage rather that to H''' ions -per se. The arguments against NaHCOs therapy

are:

1. Most acute metabolic acidosis is due to oxidizable acids:

a. Lactic/Pyruvic acid in acute hypoxia.

b. Free fatty, acetoacetic and e-hydroxybutyric acids in acute diabetic

ketoacidosis.

2. O2 release to the tissues is partly controlled by red cell 2,3-diphosphoglycerate

(DPG) [22,23] which modifies the shape of the Oa-dissociation curve. For blood

with normal O2 affinity (P50 = 3.6 kPa (27 mmHg)), a "shift to the right" of 0.4

kPa (3 mmHg) may deliver over 20 percent more O2 to the tissues. Hypoxia and

[H'''] are the two major physiological factors which, with plasma inorganic phosphorus, determine the red cell content of DPG.

Acidosis both improves O2 release in the tissues (Bohr effect) [24] and depletes red

cell DPG. Restoration of red cell DPG is a slow process (T1/2 about 11 hours) and too hasty correction of acidosis allows the DPG deficiency to predominate by shifting the O2-

64
20 40 60 80

Blood H'''nmol/l

Figure 10. Attempted resuscitation of a man aged 55 years with cardiac arrest

due to cardiac arrhythmia after a myocardial infarct. Results at point (1)

taken on blood 15 minutes after the arrest showed a profound combined

respiratory and metabolic acidosis, plasma HCO3 18 mmol/1. The PCO2 was reduced

by ventilation (2) and the metabolic acidosis was improved after infusion of

100 mmol NaHCOs intravenously (3), plasma HCO3 22 mmol/1. Further acute CO2

retention supervened (4) and the patient died within 2 hours (5).

dissociation curve to the left, which impairs O2 delivery to the tissues. Some workers

have suggested correction of acidosis only to about [H"*"] 55 nmol/1 (pH 7.26) [25] while

others have advocated intravenous sodium phosphate so as to increase the availability of

phosphorus for DPG production [26]. Equally, methods for the enzymatic determination of

red cell DPG are improving [27] so that it is now possible to measure DPG routinely in

patients with acute acid-base disturbance.

I am obliged to the editor of the British Journal of Chest Diseases for permission to make

use of material published previously [17].

References

[1] Schwartz, W. B. and Relman, A. S., A critique of the parameters used in the evaluation

of acid-base disorders. New Eng. J. Med. 268, 1382 (1963).

[2] Pierce, N. F., Fedson, D. S., Brigham, K. L., Mitra, R. C, Sack, R. B., and Mondel

,

A., The ventilatory response to acute acid-base deficit in humans, Ann. Intern. Med.

72, 633 (1970).

[3] Bone, J. M., Cowie, J., Lambie, A. T., and Robson, J. S., The relationship between

arterial PCO2 and hydrogen ion concentration in chronic metabolic acidosis and

alkalosis, Clin. Sci. Moleo. Med. 46, 113 (1974).

[4] Fulop, M., Dreyer, N., and Tannenbaum, M., The ventilatory response in diabetic

ketoacidosis, Clin. Sai. Molea. Med. 46, 539 (1974).

65
[5] Van Ypersele de Strihou, C. and Frans, A., Les desordres acido-basique au cours de

1

' insuff icance respiratoire: Un nouveau diagramme destine a leur interpretation,

Presse Med. 75, 1797 (1967).

[6] Brackett, N, C, Cohen, J. J., and Schwartz, W. B., Carbon dioxide titration curve of

normal man: Effect of increasing degrees of acute hypercapnia on acid-base equilibrium.

New Eng. J. Med. Zn, 6 (1965).

[7] Arbus, G. S,, Hebert, L. A., Levesque, P. R., Etsten, B. E., and Schwartz, W. B.,

Characterization and clinical application of the 'significance band' for acute respiratory

alkalosis. New Eng. J. Med. 280 , 117 (1969).

[8] Sullivan, W. J. and Dorman, P. J., The renal response to chronic respiratory acidosis,

J. Clin. Invest. 34, 268 (1955).

[9] Refsum, H. E., Arterial blood gases in respiratory insufficiency. Acta Chir. Sound.,

Suppl. 253, 175 (1960).

[10] Refsum, H. E., Acid-base status in patients with chronic hypercapnia and hypoxaemia,

Clin. Sai. 27, 407 (1964).

[11] Refsum, H. E., Relationship between state of consciousness and arterial hypoxaemia 1n patients with pulmonary insufficiency, breathing air, Clin. Sai. 25, 361 0963).

[12] Chiodi, H., Respiratory adaptation to high altitude, in The Regulation of Human

Respiration, D. J. C. Cunningham and B. B. Lloyd, eds., p. 363 (Blackwell, Oxford,

1963).

[13] Refsum, H. E., Hypokalaemic alkalosis with paradoxical aciduria during artificial

ventilation of patients with pulmonary insufficiency and high plasma bicarbonate

concentration, Sound. J. Clin. Lab. Invest. 481 (1961 ).

[14] Stoker, J. B., Kappagoda, C. T., Grimshaw, V. A., and Linden, R. J., A new method for

assessing states of acute acidaemia in man, Clin. Soi. 42, 455 (1972).

[15] Stoker, J. B., Kappagoda, C. T., Snow, H. M., and Linden, R. J., The assessment of

acid-base disturbance in man by the use of carbon dioxide titration curves, Clin. Soi.

and Moleo. Med. 48, 133 (1975).

[16] Acids, bases and nomograms, Lunoet (editorial), 814 (1974).

[17] Howorth, P. J. N., The physiological assessment of .acid-base balance, Brit. J. Dis.

Chest. 69, 75 (1975).

[18] Howorth, P. J. N., RIpH revisited, Lunoet, U 253 (1974).

[19] Lactic acidosis, Lanoet (editorial), n_' 27 (1973).

[20] Hydrochloric acid for metabolic alkalosis, Lunoet (editorial), l_, 720 (1974).

[21] Wilson, R. E., Gibson, D., Percinel, A. K., Ali, M. A., Baker, G., LeBlanc, L. P., and

Lucas, C, Severe alkalosis in critically ill surgical patient, Arohs. Surg., Chiougo,

105 , 197 (1972).

[22] Benesch, R. and Benesch, R. E., The effect of organic phosphates from the human erythrocyte on the allosteric properties of haemoglobin, Bioohem. Biophys. Res. Comm. 26^, 162 (1967).

[23] Chanutin, A. and Cornish, R. R., Effect of organic and inorganic phosphates on the

oxygen equilibrium of human erythrocytes, Arohs. Bioohem. Biophys. 1 21 , 96 (1967).

66
[24] Bohr, C, Hasselbalch, K., and Krogh, A., Uber einer in Biologischer Beziehung Wichtigen

Einfiuss, den die Kohlensaurespannung des Blutes auf dessen Sauerstoff Bindung, Ubt,

Soand. Arch. Physiol. 16, 402 (1904).

[25] Zimmet, P. Z., Taft, P., Ennis, G. C, and Sheath, J., Acid production in diabetic

acidosis; a more rational approach to alkali replacement, Br>. Med. J. 3^, 610 (1970).

[26] Alberti, K. G. M. M., Darley, J. H., Emerson, P. M., and Hockaday, T. D. R., 2,3-

diphosphoglycerate and tissue oxygenation in uncontrolled diabetes Mellitus, Lancet,

Jl, 391 (1972).

[27] Teunissen, A. J., De Leeuw, R. J. M., Boink, A. B. T. J., Hamelink, M. L., and Maas, A.

H. J., Comparison of five methods for determination of 2,3-diphosphoglycerate in blood,

Clin. Chem. 20, 649 (1974).

67

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

RELATIVE ACTIVITY AND SI UNITS

B. F. Visser

Department of Anaesthesiology

University Hospital

Nijmegen, The Netherlands

A. H. J. Maas

Department of Cardiology

University Hospital

Utrecht, The Netherlands

1. Sl-Units

The name International System (Systeme International, SI) of Units was adopted by the

11th Conference G^nerale des Poids et Mesures (CGPM) in 1960 [1]^- The system is now based

on seven base units for the corresponding basic kinds of quantities (table 1).

Table 1. The seven base units of the International System of Units.

(Symbols for quantities are printed in italics , symbols for

units in Roman type).

Quantity Symbol Dimension Unit Symbol of

length I L metre m

mass m M kilogram kg

time t T second s

electric current I I ampere A

thermodynamic temperature T 0 kelvin K

luminous intensity ^VJ candela cd

amount of substance n N mole mol

All other SI units are derived from base units by multiplication or division without

the introduction of numerical factors; for example the derived, coherent unit of volume is

the cubic metre (m^) and the derived, coherent unit of (substance) concentration is the

mole per cubic metre (mol/m^ or mol'm"^). Several derived SI units have been given special names, some of them are to be found in table 2.

The litre (1) is a non-coherent unit of volume. The litre has been redefined [2] and

is now exactly equal to the cubic decimetre (dm^). All other units containing the litre in

either the numerator or denominator are non-coherent, e.g.^ mol /I as the unit of substance

concentration.

Figures in brackets indicate the literature references at the end of this paper.

69
Table 2. Some derived SI units with special names.

(A full stop (•) indicates multiplication

of units. A solidus (/) indicates division

of units; alternatively the denominator

may be expressed with a negative exponent.)

Quantity Symbol Unit Name Symbol

force F I<g«m/s2 newton N

pressure V N/m2 pascal Pa

work, energy W N-m joule J

Pa.m3

kPa-1

An extremely important feature of the International System is the fact that there is only one unit of energy, the joule (1 J = 1 N-m = 1 kg^m^/s^). Work is mechanical energy

and, therefore, the unit of work is also the joule. The product of pressure (p) and volume

(f) has the dimension of work; coherent units are pascal (Pa = N/m^), cubic metre (m^) and

Pa-m^ = N-m = J. When the litre is chosen as the non-coherent unit of volume, the kilopascal (kPa) must be the noncoherent unit of pressure [3].

1 kPa-1 = 103 Pa. 10-3 ni3 = 1 Pa.m3 = 1 N-m = 1 J (1)

2. Reference Quantities

The International Organization for Standardization defines a reference quantity in the

following way [4]:

"Physical quantities are concepts used for qualitative and quantitative descriptions

of physical phenomena. Such quantities may be classified into categories, each category

containing only quantities which are mutually comparable. If one of the quantities in such

a category is chosen as a reference quantity, called the unit, any other quantity in this

category can be expressed as a product of this unit and a number, called the numerical value

of the quantity. For a quantity symbolized by A, this relationship may be expressed in the

form

4 = U) X U] (2)

where U] is here used to symbolize the unit chosen for the quantity A, and [a) to symbolize

the numerical value of the quantity A when expressed in the unit U]."

The simplest case is present, when the unit is unity (proposed symbol I, not yet accepted

by CQUCC and EPQU [5]). Examples are: volume fraction {^) and m.ole fraction {x). The SI units for plane angle and solid angle, the radian (rad) and the steradian (sr), respectively,

are called supplementary units in the International System of Units.

When a quantity has to be made dimensionless , two ways are followed.

a) The quantity A is divided by the unit U], thus giving the numerical value {a)

70
This procedure is recommended by Siggaard-Andersen [6,7].

b) The quantity A is divided by a quantity .4^=1, representative for the standard state

^0

U)- 'A (!)• A_ = (A) . (4)

This convention is advocated by Bates [8].

In both cases the result of the operation is a numerical value depending on the unit

chosen. This is in contrast with true dimensionless quantities, where the value is independent of the system of units, the unit being unity.

c) Theoretical ly--and why not practical ly--the solution is to refer to a standard

quantity

In this way a ratio of numerical values is obtained, this ratio being independent of the

system of units.

3. Relative Activities

There are three different activity scales:

a) on a concentration (c) basis

a = yo lim a /e = 1 (6) ^ c=o

^

b) on a molality (m) basis

a = ym lim a /m = 1 (7]

c) on a mole fraction {x) basis

a = f'x Mm a /x = ] (8) X X

Activities on the molality and concentration scales are related by the equation [8]:

a /a = p„ (9)

where is the density of the solvent. When the solvent is water, the value of "is very

close to unity, and for this reason no distinction is made between a and a " cm [61.

71
The value of Pq is close to unity only if the litre is chosen as the non-coherent unit

of volume. The coherent SI unit of density is kg/m^, the value of the density of water becomes

approximately 1000. Therefore, the simplification is not justified.

When activities have to be made dimensionless , the suggestion of section 2.c. may be

followed. Foreach scale there will be one standard quantity, but the relationship between

these quantities will be known.

References

[1 ] Comptes rendus des seanaes de la Heme Conferenoe Generate des Poids et Mesures, Paris,

1960 (Gauthier-Vallers, Paris, 1961).

[2] Comptes rendus des seanaes de la 12eme Conference denerale des Voids et Mesures, Paris,

1964 (Gauthier-Villars, Paris, 1964).

[3] Visser, B. F., Quantities and units in respiratory physiology. Bull. Physiopath. Hesp.

9, 513 (1973).

[4] International Standard ISO 21 series. Part 0. General principles concerning quantities,

units and symbols, 13 pp. (Internationpil Organization for Standardization, Geneva,

1974). Part 8. Physical chemistry and molecular physics, 13 pp. (International Organization for Standardization, Geneva, 1973).

[5] Minutes of the meeting of CQUCC (Commission on Quantities and Units in Clinical Chemistry,

lUPAC section on Clinical Chemistry) and EPQU (Expert Panel on Quantities and Units,

IFCC Committee on Standards), Wageningen (April 18-21, 1975).

[6] Siggaard-Andersen, 0., The Aaid-Base Status of the Blood, 4th revised ed., 229 pp. (Munksgaard, Copenhagen, 1974).

[7] Siggaard-Andersen, 0., Clin. Chem. 20, 727 (1974).

[8] Bates, R. G. , Determinations ofpH. Theory and Practice, 2nd ed. , 435 pp. (John

Wiley & Sons, New York, 1973).

72
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

SEMI-EMPIRICAL ACID-BASE PROGRAM

B, F. Visser

Department of Anaesthesiology

University Hospital

Nijmegen, The Netherlands

A. J, Hoelen, J. A. Kreuger, and A. H. J. Maas

Department of Cardiology

University Hospital

Utrecht, The Netherlands

In spite of many efforts {e.g., [1-3]^), mathematical models of the acid-base behavior

of human blood have not yet led to useful programs for calculating derived acid-base data

from measured primary data. The calculations of Scheffner and Martin [4] are restricted to the interpolation method for pC02 and the determination of actual bicarbonate, total CO2 and

standard bicarbonate.

On the other hand, empirical approaches have been quite successful. Siggaard-Andersen

[5] distinguishes between programs utilizing linear pH, log pC02 equilibration curves and

those utilizing linear pH, ctHCOa equilibration curves. We confine ourselves to the first

category.

Dell and Winters [6] store the coordinates of the base excess and buffer base curves in

the computer. Hardt [7] uses three different empirical equations for base excess, each

working in its own pH area. Recently the same system was followed by Knoll et al. [8].

Vallbona et al. [9] apply a simplified equation for the calculation of the base excess concentration. Englesson et al. [10] make use of an assumed "mathematical center," the

point of convergence for the pH-log pCOa lines.

We call our program a semi -empirical one, because the basic equations have a sound

theoretical background, but are empirically adjusted to measured data by the introduction of

second-order correction terms or factors.

The slopes of the equilibration lines of separated human blood plasma in the pH-log

PCO2 diagram could be predicted [11]. The same type of equation was used to describe the

slopes of the equilibration lines of whole blood as a function of plasma protein, hemoglobin

(Hb) and standard bicarbonate (SB) concentrations. Adding a term with Hb^ led to the

necessary accuracy for the calculation of Hb from a given slope M

M = - 1.0285 + 0.0051 Hb - (3.5242 + 0.6105 Hb + 0.009 Hb2)/SB

Base excess (BE) is calculated from Hb content and Van Slyke standard bicarbonate (VB)

BE = (1 + D X Hb)-(VB - 24) (1)

Figures in brackets indicate the literature references at the end of this paper.

73
with

D = - 0.0233(BE + 20)/(34.232 + BE) (2)

instead of taking a constant value of D [5].

For further details of the program refer to [12].

The semi -empirical approach combines high accuracy with low computer power and could be

used as a means of averaging measured data.

References

[I] Visser, B. F., Zuur-base eigenschappen van bloed bij constante temperatuur en constante

verzadiging, Utrecht (1966) 26 pp. Cited by [3] and [13].

[2] Lloyd, B. B. and Michel, C. C, A theoretical treatment of the carbon dioxide dissociation curve of true plasma in vitro, Resp. Physiol. 107 (1966).

[3] Brodda, K., On the theory of base excess curve in the Siggaard-Andersen nomogram.

Respiration 32, 378 (1975).

[4] Scheffner, D. and Martin, H., Vergleich zwischen elektronisch errechneten und aus dem

Nomogramm nach Singer und Hastings bestimmten Werten des Saurebasenhaushalts, Clin.

Chim. Acta, J_6, 297 (1967).

[5] Siggaard-Andersen, 0., The Aoid-Base Status of the Blood, 4th revised ed., 299 pp. (Munksgaard, Copenhagen, 1974).

[6] Dell, R. B. and Winters, R. W., A computer program for the blood pH-log pCOa nomogram,

Scand. J. Clin. Lab. Invest. 19, 29 (1967).

[7] Hardt, J., A computer program for calculating blood acid-base parameters on an Olivetti

"Programma 101" desk computer, Clin. Chem. 18, 658 (1972).

[8] Knoll, E., Wisser, H. and Dettmer, K., Berechung der Parameter des Saure^Basen-Haushaltes

mit Hilfe eines Kleincomputers, Z. Klin. Chem. Klin. Bioohem. 13, 37 (1975).

[9] Vallbona, C, Pevny, E. and McMath, F., Computer analysis of blood gases and of acidbase status, Comp. Biomed. Res. 4^, 623 (1971 ).

[10] Englesson, S., Grevsten, S. and Olin, A., Some numerical methods of estimating acidbase variables in normal human blood with a haemoglobin concentration of 5 g/100 cm^,

Soand. J. Clin. Lab. Invest. 32, 289 (1973).

[II] Visser, B. F. and Maas, A. H. J., The pH-log pC02 diagram of separated human blood

plasma, Clin. Chim. Acta, S_, 850 (1960).

[12] Maas, A. H. J., Kreuger, A. J., Hoelen, A. J. and Visser, B. F., A computer program for

calculating the acid-base parameters in samples of blood using a mini-computer, (Pflugers)

Arah. Ges. Physiol. 334, 264 (1972).

[13] Siggaard-Andersen, 0., The Aoid-Base Status of the Blood, 2nd ed., 134 pp. (Munksgaard,

Copenhagen, 1964).

74
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

"ACID-BASE SEMANTICS"—A CENTURY OF THE TOWER OF BABEL^

Harry F. Weisberg

Division of Biochemistry

Mount Sinai Medical Center

Milwaukee, Wisconsin 53201, USA

Almost 100 years have passed since 1877 when Walter [3]^ originated the concept of

"alkaline reserve" which was estimated by the oarhon dioxide oontent of the plasma. It was assumed that the alkali of the blood existed mainly in the form of carbonates (or carbonates

and bi carbonates as the terms are used today) and that the quantity of carbon dioxide was essentially proportional to the quantity of alkali contained in the blood. Walter [3]

stated that "...the carbon dioxide content of the blood will .. .permit of conclusions as to the quantity of alkali that has been withdrawn from the body as a whole."

Subsequently, alkali (ne) reserve has been used not only to designate the total carbon

dioxide content of blood but also the carbon dioxide capacity, the plasma bicarbonate

(determined by the carbon dioxide combining capacity or carbon dioxide combining power),

the amount of sodium combined with the plasma bicarbonate, and the total cation ("base")

concentration. In 1931, Cummer [4] wrote that the alkali reserve "...by means of which

acids are neutralized, is made up of the bicarbonates , small quantities of phosphates, and

alkaline protein compounds."

The four "CO2" determinations mentioned above are compared to each other and to the

AutoAnalyzerf^ technic in table 1. Skeggs placed the AutoAnalyzer trays in a box to equilibrate the plasma [5] or serum [6] specimens with carbon dioxide to obtain the CO2 combining

power since it was difficult to achieve the necessary anaerobic conditions during collection

and analysis required for the CO2 content. A modification was described in which each cup

was equilibrated three times just prior to aspiration into the apparatus [7]. The present

technic for the AutoAnalyzer CO2 is termed a "content" since no equilibration is utilized

but, in reality, the result is an approximation of a bicarbonate concentration.

If the PCO2 is 40 mm Hg, the relationship among the determinations is shown in eq. (1)

and (2) for arterial and venous blood, respectively.

CT,^^° = [CY^°°- 0.6] = [CC + 1.2] = [CP + 1.2] (1)

a

CTy = CCY^" - 0.4] = [CC + 1.4] = [CP + 1.4] (2)

If the PCO2 is above 40 mm Hg, the relationship is as in eq. (3);

CT > CC > CY < CP (3)

^Based upon and expanded from Weisberg [1,2] with permission of publishers, Williams and

Wilkins [1] and ASCP [2].

^Figures in brackets indicate the literature references at the end of this paper.

75
1 Q QJN

_

QJ

+-> >^ c \ CM Q CO

CM CM Q. CO o O 03 0) o O o

<_> o (/I >> c e c

o>

~r~

<;,.. E

_Q o CO E

(_j 1 ^ ^ Q O Q~ • 

^ 5 (J

C_j Q_ _Q

00 VII VII

Csl C\| CL CO Q o_ D_ O 03 CO >> c CM 1 0

"O

I/) CO Q

>^ " * •M CO CM C\J 1 1 CD CO C\J dJ 1 CM O Q • 

^

CD 00 0 CD CNl 'N CO

CU

Q o LO

C/) C\J 1 CM CM

_ CO Oif) >^

4_j _^ ^ 00 CTi CM OJ CM

(_) fO (_) 1 1 LO 4_) Q_ o _Q CM CM (tJ 4— Q Q (_) CM cu CD QJ 00

c~ CM CM CO CO

d o j_ cu o Q

^_;> _j_> CM CM Q. Q (

csl Q_ _ Q Q

s_ c^ ^ ^ 0^

p- r— CM CM t/) ^0 (_) 1 4-' CO LO Q _Q CM CM CO d) Q 4_} j Q ~ 4_) '—^ (J i_

QJ cu (U CM 0 CM CO CO c to o o C <D QJ Q Q LD

>^ >^ CM CM

fC 4J

(/) 2^ d> 0^

•r- • ,

"O CM <D O CO

<-> zr. CO cn 0^ E CO E •JZ CO CD

&.

CO CM E Q d)

OJ

Q E Q •!-> *+- E 4_) CO

OJ 0

CO

Q 13 „ + cr CU ^ |cn Q Q 4_)

Q

(/I (N (J SI o (J 0 Q_ d) 4_> -M o o 0 CM J_ d>

CO CM

Q t|H Q. d) i- E

TO CM

U

Q

+-> £Z 0. ^ i-o CU E CM _ fO (J C;

CO dJ ' * f ^ S- :3 0.

_

CU 0 C/) CO

CM

cu • r— i- QJ +-> c dJ d)

Q o o Q i- CU

Q S- cu ry. >> o o cr O 0 S- CO 1— (jU +-> <_) 4- -Q 0

76
and if the PCO2 is below 40 mm Hg, the relationship is as in eq. (4).

CY > CC > CT < CP . (4)

In 1883, Stadelmann [8] had formulated the concept of "acidosis" which resulted from excess acid production Which could then result in alkali deficit. Naunyn [9] expounded on the use of the term acidosis due to endogenous production of hydtroxyhutyric acid in diabetes,

which was then followed by the secondary alkali deficit or hypoalkalinity or alkalipenia.

Acidosis was used to designate qualitative and/or quantitative changes of acid in the blood

[10,11] but not the "reaction" of the blood [10].

From 1909, however, the consensus [12-15] seemed to revert to the position of Walter

and of Stadelmann, the term "acidosis" being used to indicate any type of acid which could

alter the acid-base balance, resulting in a lowering of the alkaline reserve. If the

respiratory mechanism could decrease the carbonic acid concentration to maintain the

"normal" bi carbonate rcarboni c acid ratio, it was called compensated acidosis^ but if the

carbonic acid could not be reduced sufficiently, with an accompanying increase of hydrogen ion (pH decreased), Hasselbalch and Gammeltoft [16] called it uncompensated acidosis.

In 1931, Peters and Van Slyke [17] used the terms 1° alkali deficit and 1° alkali

excess to denote changes in the blood. These can be traced back to 1877 (Walter [3]), and

to 1883 (Stadelmann [8]). Changes in body content were called "base" deficit and "base" excess [17]. In 1947 (first edition), Davenport [18] modified the terms to describe blood

changes as extra'fixed acid (or acid excess) and extra fixed base (or base excess). Singer

and Hastings [19,20] used combinations of these terms

fixed acid excess (or base deficit)

and fixed acid deficit (or base excess); Roos and Thomas [21] also use the terms fixed acid

excess and fixed acid deficit for metabolic acidosis and metabolic alkalosis, respectively.

In 1948, Singer and Hastings [19,20] introduced the determination of buffer base which

describes the buffer system of whole blood. It is the cation equivalent to the sum of

buffer (or "labile") anions ( "base")— bicarbonate, proteinate, and phosphate in the plasma,

and the oxyhemoglobin and other erythrocyte buffer anions [e.g., phosphate) [22,23]. The

term "buffer base" was used by Stadie, et al. [23a] in 1925 to represent the bicarbonate and

protein concentration.

The "non-bicarbonate" buffer base was designated as [Buf"]; the non-buffer (or "fixed")

anions consist of chloride, sulfate, and the organic acids. The pH and total carbon dioxide

content are determined in whole blood, which is corrected for the hematocrit or hemoglobin

content, and the partial pressure of carbon dioxide and "buffer base" calculated from a nomogram [20] or diagram [19].

The normal buffer base of whole blood of normal hematocrit is 49 mEq per liter (range

46 to 52) and, for plasma, 42 mEq per liter (range 39 to 45). The difference is due to the

much higher buffer (oxyhemoglobin) content in the erythrocytes. Note the similarity of the

definition of alkaline reserve of plasma by Cummer [4] and the buffer base of plasma (omitting

hemoglobin factor for whole blood buffer base). The buffer base is decreased in metabolic

acidosis and increased in metabolic alkalosis. Recently [24], buffer base has been equated

with total bicarbonate.

Buffer base, originally represented as B5+, was modified by Elkinton and Danowski [25]

to BB to signify the body buffers which were equal to the total buffer cations (value is

equivalent to the buffer anions, also termed Buf). The buffer base was the difference

between the fixed (total) cations and the "fixed" anions, necessary to maintain electroneutrality [26].

The symbol for buffer base, as used by Owen, et al. [27], is B~; Kintner and Gambino

[28,29] used B for "base". Roos and Thomas [21] use BBB to represent blood buffer base and

suggested the term standard blood buffer base or SBBB to designate the buffer base which

would be obtained if mixed venous blood were brought to PCO2 40 mm Hg in vivo. They utilize

the Singer-Hastings nomogram [20] to derive the conventional BBB and the SBBB; both BBB and

SBBB include the carbamino-COa [21].

77
At a symposium on pH and blood gas measurement, Astrup [72] introduced the term A acidbase to represent the total amount of surplus acid or surplus base in one liter of blood at

a given concentration of standard bicarbonate [50,51] {vide infra). One could calculate the

amount of sodium bicarbonate or ammonium chloride needed to correct a surplus acid or base,

respectively, by multiplying the a acid-base per liter by the body weight in kilograms times

0.3 (the extracellular space) [31,72].

In 1960, Astrup and his colleagues [30-32] changed acid excess and base excess as descriptive terms for the nonrespiratory factors to base excess ±. Base excess is defined

as the deviation of the buffer base (aBB) from the "normal" buffer base. Instead of the

actual titration of the blood to pH 7.40 at 37 °C at a PCO2 of 40 mm Hg, it is much easier

to calculate BE [33] or to derive it from the curve nomogram [34], alignment nomogram [35], or slide rules [36,37].

To avoid the conceptual ambiguity of "negative base excess" Lyons and Moore [38] use buffer base deviation; Kintner and Gambino [28,29] use base deviation. Suero and Woolf [39] use the term nonrespiratory acid excess.

Davenport [18] uses the terms base excess (as previously) and base deficit (for his

previous acid excess) for blood changes; these were the terms used by Peters and Van Slyke

[17] in 1931 for changes in body content. Davenport [18] utilized the bicarbonate-pH

diagram for graphic representation of his terms extra fixed acid (or acid excess or base

deficit)--point below the buffer slope of carbon dioxide--and extra fixed base (or base

excess)--point above the line. I have used similar graphic representations [40] on a modified CO2 content-pH diagram first described by Cullen and Jonas [41] and Hastings et al.

[42]. The descriptive term for the (metabolic) acid-base alteration is delta CO content

(ACT), eq. (5).

^

A CT = Actual CO2 CT - Theoretical CO2 CT (5)

The theoretical CO2 content can be calculated for arterial and venous blood, utilizing a buffer slope or molar buffer value of 28, from the intercept of actual pH with the CO2

buffer slope or from eqs. (6) and (7).

Theoretical CO2 CT, =25+28 (7.40 - pH) (6)

a

Theoretical CO2 CT^ =27+28 (7.37 - pH) (7)

The values for arterial and venous theoretical CO2 content also may be obtained from

the table on the Weisberg TRI-SLIDETM calculator for the Henderson-Hasselbalch equation.

A "negative" delta content (ACT) signifies a bicarbonate (combined CO2) deficit or metabolic

acidosis and is equivalent to the "base excess negative" or "base deficit" whereas a positive ACT signifies bicarbonate (combined CO2) excess or metabolic alkalosis and is equivalent to the "base excess positive" or "base excess". Changes in the bicarbonate

concentration may be substituted for the CO2 content changes but only when the PCO2 value is normal

.

Since the non-bicarbonate buffers [Buf~] are not directly measurable, Filley [43]

utilized the corrected a[HC03] as the base excess. However, the actual A[HC03]--the difference between the actual plasma bicarbonate and 24 mEq per liter— seldom differs by more

than 3 mEq per liter from the "base excess" [43]. Collier et al. [44] have applied the

concept of base excess to the extracellular fluid (BE^rp) rather than the blood. It can be

calculated from their bicarbonate-pH diagram (with buffer line of extracellular fluid rather

than that of plasma) or from their formula;

BE^(,p = [HC03]p + e^(,p (pH - 7.4) - 24 (8)

78
in which [HCOslp is the actual plasma bicarbonate and 3ecf "^^ ^"^^^ apparent buffer

value (approximately 12). One can transform this concept to that of "ACT" by altering eq.

(5) to eq. (9).

ECF^^^ = Actual CO2 Content - Theoretical ECF(,q (9)

Similarly, eqs. (6) and (7) are transformed to eqs. (10) and (11).

Theoretical ECF^q = 25 + 12 (7.40 - pH) (10)

Theoretical ECF^q^ = 27 + 12 (7.37 - pH) (11)

Use of eqs. (9), (10), and (11) is much easier than the necessary plotting on diagrams and

finding the vertical displacement, etc.; special slide rules [37,45,46] are easier to use than the mathematical calculations with formulas.

In 1930, Henderson, et at. [47] introduced the term, T^q? to represent the millimoles

of total carhonic acid per liter of blood when the PCO2 equals 40 mm Hg. The carbon dioxide

content had also been standardized to a pH of 7.40. The formula of Peters and Van Slyke

[17] for the standardized carbon dioxide content is

[C02CT]^u = [CO2CT] + (pH - 7.40) (8.2 + 2.6 Hb) - 0.36 Hb02

P"7.40

(12)

in which the carbon dioxide content, hemoglobin (as oxygen capacity), and oxygen content are expressed as millimoles per liter. Peters and Van Slyke [17] also had a formula for the

standardized bicarbonate, adjusted to pH 7.40.

[HC03]„M = [HCO^] + (pH - 7.40) (8.2 + 2.3 Hb) (13)

The hemoglobin was expressed in terms of millimoles per liter which can be obtained by

dividing grams percent by 1.67. This equation can be rewritten in terms of grams percent

hemoglobin.

[HC03]p,^^ = [HCO3] + (pH - 7.40) (8.2 + [1.37 x g% Hb]) (14)

It was assumed that the total plasma protein concentration was constant at 7 g per 100 ml.

In 1952, Bunker et al. [48] reported a modified formula for the corrected bicarbonate.

[HC03]^n = tHCO;] + (pH - 7.40) (8.6 + [1.4 x g% Hb]) (15)

Eichenholz, et al. [49] determined the standardized bicarbonate using the intersection at pH

7.40 of a line parallel to the normal carbon dioxide absorption slope.

79
In 1957, Astrup and his associates [50,51] introduced the term standard bicarbonate,

which is defined as the concentration of bicarbonate in plasma separated from the cells with

the hemoglobin completely oxygenated at a PCO2 of 40 mm Hg and at a temperature of 38 °C. They use the standard bicarbonate to characterize nonrespiratory ("metabolic") disturbances

because the carbon dioxide content will vary with the PCO2 and P02 present in the blood

sample. The standard bicarbonate is approximately 1.2 mmol per liter less than the carbon

dioxide capacity determined at a PCO2 of 40 mm Hg but on arterial ("completely" oxygenated)

blood at 38 °C (instead of the usual procedure utilizing venous blood at 20 °C). It is not

as useful as changes in buffer base or of base excess [30]. The original [52] and modified

[34] Siggaard-Andersen curve nomogram yields values for standard bicarbonate, etc. In 1967, however, the standard bicarbonate on the nomogram (distributed by Radiometer A/S) was changed to plasma bicarbonate (at PCO2 40 mm Hg) with the supporting descriptive statement

that this "plasma bicarbonate" is the standard bicarbonate.

Armstrong, et al. [53] prefer T^q bicarbonate (named after L. J. Henderson [47] who

based many of his studies on blood at a CO2 tension (PCO2') of 40 mm Hg)--also equilibrated

to PCO2 40 mm Hg, but as if it occurred in the subject [in vivo) rather than after the

sample is removed {in vitro). The T^g bicarbonate differs from the standard bicarbonate by

the amount of dilution by the interstitial fluid of the bicarbonate generated during hypercapnia. As explained above, the standard bicarbonate can be described as a "modification"

of the carbon dioxide capacity; similarly the T^g bicarbonate is similar to the carbon

dioxide combining capacity, the bicarbonate concentration at a PCO2 of 40 mm Hg (determined

on separated plasma without the hemoglobin buffering capacity). It can be calculated from

the Singer-Hastings nomogram [20] (using an "effective" (extracellular) hematocrit of 0.09)

or the Astrup log PCO2/PH diagram [30,52] (using an "effective" hemoglobin of 3 g per 100

ml); such calculations utilize a "dilution factor" of five to obtain the effective hematocrit or hemoglobin (based on ratio of blood volume to extracellular fluid and ratio of

buffering capacity of blood and interstitial fluid). In addition, the Singer-Hastings

nomogram and the Astrup diagram have been modified [44], allowing direct estimation of the

T40 bicarbonate without hematocrit or hemoglobin values, respectively. The Ti+q bicarbonate

can be estimated mentally if the PCO2 is known. For "hypercarbia" , eq. (16) is used [53],

whereas for "hypocarbia" , eq. (17) is used [54].

PCO2 - 40

[HCO3]

-oivo l-^^'-'sl vitro 15

^"'^

PCO2 - 40

[HCO^] . . = [HCO;] . ^ in ^n vt.vo ^ xn v%tro ID

The in vivo buffer curve for carbon dioxide (expressed as [HCO3]) is curvilinear

rather than rectilinear [18,55]. Davenport [18] describes the relationship of bicarbonate

for specific values of PCO2 as eq. (18).

PCO2

[HCO;] = 31.39 X (18)

PCO2 + 12.95

The same curvilinear relationship holds for the carbon dioxide titration curve when expressed

as carbon dioxide content. It is best therefore to estimate ACT or base excess/deficit from

the in vivo buffer data when the pH is below 7.40 ("acidosis") and from the in vitro buffer

data when the pH is above 7.40 ("alkalosis").

In similar fashion, Severinghaus and Bradley [56] state that the blood in vivo behaves

as if the hemoglobin concentration were one-third (e.g., 5 g per 100 ml) of the actual value

since the extracellular fluid is about twice the blood volume and the bicarbonate, generated

from carbon dioxide, diffuses from the blood into the extracellular fluid [sic). With the

Siggaard-Andersen alignment nomogram [35] they use the one-third hemoglobin value on the

80
original line as a fulcrum to connect with PCO2 40 mm Hg (rather than the "original" base excess with PCO2 40) and thus derive the in vivo standard (plasma) bicarbonate and base

excess (latter read at the actual hemoglobin concentration).

An in vivo carbon dioxide titration curve, based on a buffer capacity of about 5 g per

100 ml hemoglobin was reported by Siggaard-Andersen in 1967 [57] and expanded upon in 1971

[58], giving values for base excess and base deficit, and for bicarbonate at PCO2 40 mm Hg.

Roos and Thomas [21] discuss the "error'" of the conventional standard bicarbonate (SB

in vitro minus SB in vivo) which should be subtracted from the in vitro value to obtain the

estimated in vivo value of standard bicarbonate; the "error" varies with the pH, hematocrit,

and PCO2.

Owen et at. [27] proposed the term appropriate hicarhonate , that bicarbonate appropriate

to the actual Pao2 present in the patient. The difference between the appropriate and the

actual (observed) bicarbonate values is a measure of the change in plasma buffers.

A Plasma buffers = [HCOs]^^^^^^ - [HCOi] . . (19)

In 1916, Hasselbalch [11] saturated blood with carbon dioxide at 37 °C under 40 mm

tension and determined the "reduced hydrogen ion concentration" (or reduced pH or nonrespiratory pH). Under such conditions the carbonic acid concentration is fixed and the hydrogen

ion concentration must vary inversely as the concentration of (sodium) bicarbonate; thus the

"reduced hydrogen ion concentration" is a measure of the blood bicarbonate [14,14a].

In 1959, Peirce and his colleagues [59,60] determined the metabolic and respiratory pH

factors by a "double pH" method--measuring the actual pH and the euoapnic pH (blood equilibrated to PCO2 40 mm Hg). If the actual and eucapnic pH values were the "same", a pure

"metabolic" imbalance could be present; a low pH indicated metabolic acidosis and high pH

indicated metabolic alkalosis. If the eucapnic (but not the actual) pH equalled "7.40", a pure "respiratory" imbalance was present. If the actual pH was lower than the eucapnic pH,

respiratory acidosis was present; conversely, if the actual pH was higher than the eucapnic

pH, respiratory alkalosis was present.

In 1965, Whitehead [61,62] described changes in acid-base balance in terms of hydrogen

ion units, expressed as nanoequi valents (or nanomoles) per liter. The pH values must be

converted into hydrogen ion concentration; the Astrup diagram is used to determine the

hydrogen ion concentration at PCO2 40 mm Hg (with hemoglobin fully saturated). The total

change is the sum of the changes in respiratory and nonrespiratory factors, a[H''"] = R + NR.

The complete equation is.

in which [H ]tot is the actual hydrogen ion concentration for the determined pH at the

actual PCO2 in the patient and [H"*"] PCO2 40 is the hydrogen ion concentration for the pH at PCO2 of 40 mm Hg. Each term has a "normal" range of ± 4 nmol ; a negative sum means a deficit

of hydrogen ions (alkalosis) and a positive sum, an excess of hydrogen ions or acidosis.

Kintner [63,64] describes acid-base changes in terms of A and B, respectively, the

deviation of PCO2 ("acid") and carbon dioxide content ("base") from their respective normal

values. Subsequently, the "base" portion was equated with changes in buffer base (base excess and base deficit) and the Siggaard-Andersen alignment nomogram was used to determine

the A and B factors; in addition, the term delta [H"*"] was used to quantitate the acid-base

imbalance [28,29].

81
The molar buffer value was used by Van Slyke to describe buffering capacity (or buffer

slope [43])--the number of moles of acid needed to decrease pH by one unit in one liter of a molar solution. In 1965, Woodbury [65] introduced the term Van Slyke or slyke (abbreviated

si) for the molar buffer value, defined as the ratio of the change in base [65] or bicarbonate [43] to the change in pH as the measure of "buffer capacity" of blood due to the pre- sence of hemoglobin and plasma proteins. Hemoglobin has a buffer value of 3 sl/mmol and

plasma protein has a value of 0.1 sl/g. Instead of converting hemoglobin as grams per

deciliter to millimoles per liter by dividing g/dl by 1.67 or multiplying g/dl by 0.621, the

entire relationship can be described in eq. (21),

sl/1 = (1.86 X Hb g/dl) + [TP g/dl x (1 - Hct)] (21)

in which the Hb is hemoglobin, TP total protein, and Hct hematocrit.

Substitution of "average" values of 14 g/dl and 6 g/dl for hemoglobin and total protein,

respectively, with a hematocrit of 0.45 gives 29.3 sl/1 of blood. Filley [43] gives 29

sl/1 as a normal value for blood of pH 7.3 to 7.5. For "in vivo" conditions as exemplified

in a 70 kg individual with a blood volume of 7.1 percent of body weight ("average" of 7.6

percent for average male and 6.6 percent for average female), there are 146.5 si in the 5 1 of blood. The "interstitial fluid" volume in both sexes is about 12 percent of body weight

and contains an average of 0.66 g/dl protein (with a hematocrit of "zero"); therefore 0.66

g/dl X 10 X 0.1 x 8.4 1 yields 5.5 si. The grand total of 152 si divided by the 13.4 1 of

extracellular fluid yields a concentration of about 11.3 sl/1. Factors of 28 si and 12 si are utilized in eqs. (6) and (7), and (10) and (11), respectively.

Data for PCO2, [HCO3], and [H"*"] changes have also been presented as index values (sample

value/median normal value), permitting calculation of the compensation (and buffering)

present as a percentage of the (calculated) "uncompensated" increase of hydrogen ion [66].

Table 2 summarizes some schemes for determining and describing acid-base alterations of

the blood. The "old" system consisted only of measuring the alkaline reserve (carbon dioxide

content), a decreased value establishing a diagnosis of acidosis [4]; the possibility that

alkalosis could exist was denied [13].

There are no significant differences among the scheme that I use, or Siggaard-Andersen's

[33] or the report (not recommendations, as originally intended) from the New York Academy

of Sciences [67]. It was agreed that the preferential order for describing an acid-base

alteration is to (a) use the specific numerical data for the test result, (b) describe the

test values as "high, low or normal", or (c) use the general "descriptive" terms.

The ad hoc committee [67] suggested that acidemia and alkalemia be used to describe

alterations of blood pH and to utilize acidosis and alkalosis to describe the overall physiological process or condition which tends to cause a deviation in pH without being dependent upon deviation of the pH per se.

My criticism of acidemia-alkalemia is that these terms refer to blood and leave the pH

of the interstitial fluid (in equilibrium with the intravascular fluid) in limbo. When pH

changes of the intracellular compartment are reported (primarily research), it is so designated since the normal values differ from that of the "extracellular" fluid. These comments

are even more relevant today with the emphasis on the in vivo titration curve, etc. [44,56-

58]. A possible compromise may be to use eupUemia and eupHuria to represent a normal pH of

blood and urine, respectively, since these terms "are compact, easily spelled, and easily

pronounced" [68]. In a similar vein one could use the terms hypopHemia and hyperpUemia to designate decreased pH (acidosis and/or acidemia) and increased pH (alkalosis and/or alkalemia), respectively.

All agree that the respiratory changes are determined by measuring PCO2. I have

designated the test as PCO2 or [H2CO3]; the latter (used by Peters and Van Slyke) is equal

to 0.03 x PCO2 and emphasizes the carbonic acid term, similar to that of the HendersonHasselbalch equation, avoiding confusion with the various "CO2" determinations. The descrip82
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83
[HCO;]

Combined CO2 [HCO3]

CO2 combining capacity

CO2 combining power

Alkaline resemje^

available anion

available "acid"

gaseous "acid" (anion)

volatile "acid" (anion)

"Base" bicarbonate

+ [Prot"] + [HPO"] + [cr] + [so^] +

[Org Acids"]

T

Y

[HCOi]

Buffer base

diffusible buffer

V

Ibuf"^

Non-bicarbonate buffer base

Non-diffusible buffer base (P")'

Fixed anions

fixed (non-carbonic) "acids"

non-gaseous "acids" (anions)

non-volatile "acids" (anions)

I y

base

V

Alkaline resevve (Cummer)^

Labile anions

Buffer anions (also Buf~)

Buffer base (Bg+) (B") (B) (BBB)*^

Equivalent to: buffer cations

body buffers (BB)

V

Fixed anions

non-buffer anions

V

TOTAL ANIONS^

Total "acids"

Additional "definitions" of alkaline reserve:

Total CO2 content: combined CO2 [HCO3] + free CO2 [H2CO3]

CO2 capacity: (combined CO2 + free CO2) at PCO2 40 mm Hg and 20 °C

Portion of [Na"*"]: "base" bicarbonate or bicarbonate-bound "base" or available alkali

or available "base" or available cation.

. Total cations or total "base" or fixed cations or fixed "base."

A Non-diffusible buffer base also includes hemoglobin proteins.

Buffer base of plasma also defined as bicarbonate plus protein. Buffer base of whole

, blood also includes buffer activity of hemoglobin, etc. of erythrocytes.

Total anions minus those anions which have been determined equals RESIDUAL ANIONS or Residual "acids."

Figure 1. Terms applied to various fraction of total plasma anions modified from Weisberg [2];

duplicated terms in italics.

84
tive terms of the committee for respiratory alterations do not meet with my support. For conformity with acidemia and alkalemia, they should use hyper- and hypo-capnemia. There is easy confusion between the words -capnia (Greek root kapnos meaning smoke or vapor) and

-capnea (Greek root pnoe meaning breathing). Some authors have used the terms hyper- and

hypo-carbia (Latin root aarbo meaning coal), which are "barbarisms"— combinations of Greek

prefixes and Latin nouns [69].

It is obvious that the major disagreements are in relation to the diagnosis of the

"metabolic" al terations--especiany which test to use. Utilizing the Br0nsted concept, all

anions are conjugate bases and could fit into the descriptive terms (hyper- or hypo-basemia)

of the committee report, but only the bicarbonate anion and organic acid anions have a primary effect on the hydrogen ion concentration. Bicarbonate is considered to be part of

the labile "buffer" anions (equivalent to buffer base) whereas the organic acids are part of

the fixed "non-buffer" anions.

Although not intended to confuse the reader, figure 1 shows the interrelationships of some of the various terms applied to different components of the total anions; some duplications are self-evident. Though all the anions are "conjugate bases" and can act as a base

by accepting a proton, the "buffer" anions are the major buffers since their pK' values are closer to the pH of blood (the pK' values of the "fixed" anion systems are lower).

Figure 2 is a simple flow chart for the differential diagnosis of acid-balance imbalance.

Of potential thirteen diagnoses, seven can be distinguished by use of pH and PCO2. To distinguish between a pure respiratory or a mixed respiratory and metabolic imbalance,

the actual bicarbonate concentration or CO2 content is utilized; and to distinguish between

an acute or chronic respiratory imbalance, the base excess/deficit or delta CO2 content is necessary.

pH

PCO2 and/or [H2CO3]

[CO2 CT] or [HCO^]

mixed vs respiratory

ACT or BE/D

acute vs chronic

respiratory

~\ r

I N

3 4

I 1

8 (9 or 10)

10

I

n (12 or 13)

N12 13

Figure 2. Flow chart for differential

Key to numbers in chart:

1) normal acid-base

2) resp acid, comp

(met alk, comp)

3) resp alk, comp

(met acid, comp)

4) met acid, acute

5) met acid, partial comp

6) met alk, acute

7) met alk, partial comp

diagnosis of acid-base imbalance.

8) mixed resp acid and

met acid

9) resp acid, acute

10) resp acid, chronic or

partial comp

11 ) mixed resp alk and

met alk

12) resp alk, acute

13) resp alk, chronic or

partial comp

It is ironic that we have come full circle in the course of a century. As stated by

Filley [43], "Ambiguity in the meaning of words, a major curse of specialization, arises

85
both because words often outlive the concepts they originally stood for and because the same word is used by different experts to mean different things, i.e., abused." Additional

commentaries are by two poets--ex-Senator Eugene McCarthy [70] said, "If the language is debased or misused, if the meaning of words is obscured, the basis for common judgment is undermined, if not destroyed"; and W. H. Auden [71] said "... there is only one political

duty, and that is to defend one's language from corruption. When it's corrupted, people

lose faith in what they hear, and this leads to violence."

If one can communicate without misunderstanding, it does not make much difference which

"terms" are usedl One cannot, however, take the position described by Lewis Carroll in

Through the Looking Glass— "When I use a word, Humpty Dumpty said, in a rather scornful

tone, it means what I choose it to mean. Neither more nor less."

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[59] Peirce, E. C, II, Further development of a simplified method for determining metabolic

and respiratory pH factors. Trans. Amer. Soc. Artif. Intern. Organs, 6^, 240 (I960).

[60] Peirce, E. C, II, Effects of 2-Amino-2-hydroxymethyl-l ,3-propanediol (Tris) during

cardiac bypass procedures, Ann. N.I. Acad. Sci. 92^, 765 (1961).

[61] Whitehead, T. P., Ph.D. Thesis, U. of Birmingham (1964).

[62] Whitehead, T. P., Acid-base status, pH and PCO2 (letter to editor). Lancet, 2_, 1015

(1965).

[63] Kintner, E. P., The A/B ratio: a new approach to acid-base balance, Amer. J. Clin.

Path. 47, 614 (1967).

[64] Kintner, E. P., Calculation of the A/B ratio: a new alignment nomogram. Tech. Bull.

Registry Med. Teohnol. 38, 254 (1968).

[65] Woodbury, J. W., Regulation of pH, in Physiology and Biophysics, T. C. Ruch, and

H. D. Patton, eds., 19th ed., pp. 899-934 (W. B. Saunders, Philadelphia, PA, 1965).

[66] Quintero-Atencio, J., Index values in acid-base characterization, Clin. Res. 19, 545

(1971). —

[67] Report of ad hoc committee on acid-base terminology, in Current Concepts of Acid-Base

Measurement, C. G. Nahas , ed., Ann. N.Y. Acad. Sci. 133_, 251 (1966).

[68] Schmidt, J. E., Medical lexicographer; normal pH, Mod. Med. p 119 (June 29, 1970).

[69] Lamont, A., "Hypercapnia" versus "Hypercarbia" , Anesthesiology, 22, 324 (1961).

[70] McCarthy, E., quoted in Hechinger, G., The insidious pollution of language. The Wall

street Journal (October 27, 1971).

[71] Auden, W. H., quoted in Time (November 1, 1971).

[72] Astrup, P., Ultra-micro-method for determining pH, PCO2 and standard bicarbonate in

capillary blood, in A Symposium on pH and Blood Gas Measurement, R. F. Woolmer, ed.

,

pp. 81-93 (Little, Brown & Co., Boston, MA, 1959).

89

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

TRI-SLIDETM CALCULATOR FOR HENDERSON-HASSELBALCH EQUATION

AND C02RREC°t-02-SLIDETM FOR TEMPERATURE CORRECTIONS

OF pH, Pc02, AND P02

Harry F. Weisberg

Division of Biochemistry

Mount Sinai Medical Center

Milwaukee, Wisconsin 53201, USA

The general form of the Henderson-Hassel balch equation

may be rewritten for the bicarbonate:carbonic acid system with the usual apparent pK' of

6.10 (in whole blood or plasma at 37 °C body temperature) as eq (2).

[HCOi]

pH = 6.1 + log

[H2CO3]

The total carbon dioxide content (CO2 CT) is essentially the sum of the bicarbonate {oomhined

CO2) and carbonic acid {free CO2); the latter is proportional to the partial pressure of

carbon dioxide (PCO2 symbol as used by respiratory physiologists) and includes the dissolved

(free) carbon dioxide.

[CO2 CT] = [HCO3] + [H2CO3] (3)

[HCOi] = [CO2 CT] - [H2CO3] (4)

[H2CO3] =aPc02 = 0.03 PCO2 (5)

Since [HCU3] and [H2CO3] are not directly determined, eq (4) and (5) are substituted

into eq (2) giving rise to eq (6).

[CO2 CT] - 0.03 PCO2

pH = 6.1 + log (6)

0.03 PCO2

1. TRI-SLIDE

If any two of the three unknown quantities in eq (6) are known, the third can be

calculated [1-3]^. The TRI-SLIDE Calculator (fig. 1) (first designed in 1963 with revisions

in 1965, 1969, and 1971) has been redesigned; it reduces all calculations to one setting

of the special slide rule [4-6].

The calculator. has a table summarizing the three possible settings. If the pH and PCO2

and/or carbonic acid values are known, setting the PCO2 value on the left hand movable scale

(or the [H2CO3] on right hand movable scale) opposite the pH value on center section allows

one to read, opposite the red arrow (at 7.61 on pH scale), the [CO2 CT] on the left hand

Figures in brackets indicate the literature references at the end of this paper.

91
scale. If the pH and [CO2 CT] values are known, the PCO2 and/or [H2CO3] values are obtained

by setting the [CO2 CT] on left scale opposite red arrow and reading the PCO2 and/or [H2CO3]

opposite the pH. Finally, the pH can be determined if the [CO2 CT] and PCO2 and/or [H2CO3]

values are known; the [CO2 CT] is set opposite the red arrow and the pH read opposite the

PCO2 and/or [H2CO3].

Normal average values (and ranges) for adult males and females for arterial and venous blood are given in a table on the calculator (fig. 1); in addition, "normal" ranges are bracketed in red and blue (for arterial and venous blood, respectively) on the various slide

Figure 1. TRI-SLIDE calculator.

7.34-

7.44)

7.35-

7.45)

7.31

741)

7.32 7

42)

X

S OT y 0

1-5

32-42) 35-45) 42

55]

39

52; 3 2 n z0

A:

1.20(1.0!

1.11

(0.96

V:

138(1

2(

1.29

(i.i; |]z >

<> i-1.35)

1-1.26) )

1.65J

M.561 C

m

v>

<: a* >0

•"E

1p TO>

Z

<: » m

^E -HE

3.? CO-d

0

-2TO

+2) 1?

[H2C03]

SX mo gi If*

sill*

§? fa 3 -ml

A a % ^

q - a> 'J S aw r

> Sd3 g. 1 f S ?.

It

CD

CD

C/3

§5

.0eu

KNOW

pH and Pco? or[H2C03] pH and [COj CT]

[CO2 CT] and PCO2 [CO, CD

SET

RCO2 or [H2CO3]

[CO2 CT]

OPPOSITE OPPOSITE

pH

pH

PCO2 or [H2CO3]

[CO2 cri

PCO2 or [H2CO3]

pH

[H2CO3] = 0.03 X PCO2 [HCOsT = [CO2 Cri - [H2CO3]

wvimti fflVlVD Iheo.COjCI ItiM.COjCT

(fflniol/l) [« + ] [Hcr] (mmol/D

(m\/\) pH [H2CO3] 3€i 37.4 100 7.00 8.0/1 29.8 31.4

1, Theor. COj CT = 25 + 28 (7.40-pH)

2 TlieorC02CT„„ = 27 + 28 (7 37-pH)

3. Theor.ECFMzCTirt = 25 + :2 (7 40-pH)

4. Theor. ECFc02(rr,„ = 27 + 12 (7-37-pH) 6UID£ TO THERAPY

The following formula (based on extracellular fluid volume of 25% of body weight) is used

as a "guMe" to the parenteral therapy of the Bicarbonate Deficit or Bicarbonate Excess compofKnt. It is best to administer hall of the calculated volume and reevaluate the pttitirt before therapy is continued.Negative answer denotes: ml 1/6 molar Sodium lactate (or bicarbonate,

acetate, citrate, gluconate, etc.) for Bicarbonate Deficit ("metabolic acidosis" or "Base Deficit"). 1.5 X kg X JCT = OR

Positive answer denotes; ml 1/6 molar Ammonium chloride for Bicarbonate Excess ("metabolic alkalosis" or "Base Excess"). iCT("BE/D") = Actual CO2 CT - Theoretical CO2CT SOLUTION mmol/l FACTOR SOLUTION mmol/l FACTOR

"1/6 molar" 166 1.5 4.2% NaHCOs 500 0.5 5.0% NaHCOj 595 0.4 7.5% NaHCOa

(50 ml "ampoul") 892 0.28 8.4% NaHCOs

(40 ml "ampoul") 1000 0.25

Gastric or «3 70 3.5

(for NH^n

2.14% NH,CI 400 0.6

C1I974 AMERICAN Slid*-Charl Cor WhMlon, III. «0I87. frinl*d In U.S.A.

92
rule scales. The carbonic acid, bicarbonate, CO2 content, and ACT (defined below) are expressed as millimoles per liter (mmol/1). The values for PCO2 are given in terms of

pressure (P) as millimeters of mercury (mmHg) and have been expressed as "torr" units; they

also can be expressed (in keeping with the recommendation of the lUPAC and IFCC) in terms of

pressure (p) as kilopascals (kPa).

One pascal (Pa) equals 1 newton (N) per square meter (N/m^) ; 1 bar equals 100 000 N/m^

or 100 kPa; 1 mbar equals 0.1 kPa or 0.75 mmHg whereas 1 mmHg (or 1 "torr") equals 1.333 22

mbar or 0.133 322 kPa; and 1 kPa equals 7.5 mmHg or 10 mbar. Equation (5) gave the concentration of carbonic acid utilizing the factor of 0.03 (at 37 °C) and the partial pressure of

carbon dioxide (PCO2) expressed in mmHg; for the "proposed" SI notation, eq (5) can be

rewritten as

[H2CO3] = 0.23 PCO2 (7)

in which the partial pressure of carbon dioxide (symbols pC02 or PCO2) is expressed in kPa. Table 1 compares the gas values in terms of mmHg and kPa.

The delta content (ACT) has the same connotation as Base Excess/Deficit.

ACT (or "BE/D") = Actual [CO2 CT] - Theoretical [CO2 CT] (8)

Table 1. Comparison of "gas" values.

PRESENT (mmHg)

(mmHg = 7.5 x kPa)

"PROPOSED" (kPa)

(kPa = 0.133 X mmHg)

PC02

A: 40 (35-45)M

37 (32-42)F

V: 46 (42-55)M

43 (39-52)F

A: 5.3 (4.7-6.0)M

4.9 (4.3-5.6)F

V: 6.1 (5.6-7.3)M

5.7 (5.2-6.9)F

P02 A: 95 (75-100)

V: 40 (30-50)

A: 12.7 (10-13.3)

V: 5.3 (4.0-6.7)

P50 (T50) "26.6" "3.5"

A positive delta content (ACT) signifies Bicarbonate or Combined CO2 Excess (metabolic

alkalosis) and is equivalent to Base Excess (BE) or "Base Excess Positive." A negative

delta content (ACT) signifies Bicarbonate or Combined CO2 Deficit (metabolic acidosis) and

is equivalent to Base Deficit (BD) or "Base Excess Negative."

The "theoretical" [CO2 CT] for arterial or venous blood [in vitro and in vivo or "extracellular") can be obtained from (a) the window on the calculator (which also gives the

equivalent hydrogen ion concentration as nanomoles per liter (nmol/1) and the ratio of

bicarbonate.-carbonic acid for various pH values); (b) e.g..

Theoretical CO2 CT

Theoretical CO 2 CT

Theoretical ECF

Theoretical ECF

art

ven

CO 2 CT

CO 2 CT

art

25 + 28 (7.40

27 + 28 (7.37

25 + 12 (7.40

27 + 12 (7.37

pH)

pH)

pH)

pH);

(9)

(10)

(11)

(12) ven

93
or (c) the Weisberg Acid-Base Balance Evaluation Diagram (fig. 2) from the intersection of

the actual pH of the specimen and the in vitro or in vivo respiratory buffer lines.

6.8 6.9 7.0 7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8

0

'

I M t I

I I I I I I

I I I 1^1 I I I

I I I I I ;

I I 1 I I I 1 I I II 1 iTl I I I I

I I I I I 1 pH 6.a *.9 7.0 7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8 [H ^ ] nmol/l 159 126 100 80 63 50 40 32 25 20 16

IMCOj- l S *^ JL i2 'iJ i5 20 2S 32 40 SO (H,CO,l T I 1 1 1 1 I I I 1 1

ACID-BASE BALANCE EVALUATION DIAGRAM (1974 REVISIONS

Figure 2. Acid-base balance evaluation diagram.

| The buffering capacity--the number of moles of acid needed to decrease pH by one unit | in one liter of a molar solution—was termed the molar buffer values by Van Slyke. Woodj bury [7] introduced the term Van Slyke or slyke (abbreviated si) for the molar buffer , value, defined as the ratio of the change in base [7] or bicarbonate [8] to the change in pH i as the measure of "buffer capacity" of blood due to the presence of hemoglobin and plasma I proteins. Hemoglobin has a buffer value of 3 sl/mmol (or 1.86 si per g/dl) and plasma

protein has a value of 0.1 sl/g. The entire relationship can be described in eq (13), |

!

sl/1 = (1.86 X Hb g/dl) + [TP g/dl x (1 - Hct)] (13) j;

in which the Hb is hemoglobin, TP total protein, and Hct hematocrit. | Substitution of "average" values of 14 g/dl and 6 g/dl for hemoglobin and total protein,

respectively, with a hematocrit of 0,45 gives 29.3 sl/1 of blood. Filley [8] gives 29 sl/1 i

I 94 :

I
as a normal value for blood of pH 7.3 to 7.5. For "in vivo" conditions as exemplified in a 70 kg individual with a blood volume of 7.1 percent of body weight ("average" of 7.6 percent

for average male and 6.6 percent for average female), there are 146.5 si in the 5 1 of

blood. The "interstitial fluid" volume in both sexes is about 12 percent of body weight and

contains an average of 0.66 g/dl protein (with a hematocrit of "zero"); therefore 0.66 g/dl

X 10 X 0.1 X 8.4 1 yields 5.5 si. The grand total of 152 si divided by the 13.4 1 of extracellular fluid yields a concentration of about 11.3 sl/1. Factors of 28 si and 12 si are utilized in eqs (9) and (10), and (11) and (12), respectively, and for the theoretical CO2

content values on the TRI-SLIDE table (fig. 1) and in the in vitro and in vivo buffer lines

on the Weisberg Acid-Base Balance Evaluation Diagram (fig. 2).

The in vivo buffer curve for carbon dioxide (expressed as [HCO3]) is curvilinear rather

than rectilinear [9-11] and expressed as eq (14) by Davenport [9] to give the value for

bicarbonate at specific values of PCO2 (which fall along the buffer curve).

[HCO3] = 31.39 X PCO;

PCO2 + 12.95

(14)

The same curvilinear relationship holds for the carbon dioxide titration curve when expressed

as [CO2 CT]. It is best to estimate the ACT (or BE/D) from the in vivo buffer data when the

pH is below "7 AO" ("acidosis") and from the in vitro buffer data when the pH is above

"7.40" ("alkalosis").

2. Diagnosis

Seven out of 13 possible conditions of acid-base imbalance can be diagnosed (fig. 3)

from the pH and PCO2 and/or [H2CO3]. The CO2 content or the actual bicarbonate value is

used to distinguish a "mixed" imbalance from respiratory conditions; and the ACT or base

excess/deficit (BE/D) is needed to differentiate between acute and chronic (partial compensated) respiratory conditions. The ACT or BE/D is also used in calculating the amount and

type of fluid therapy required.

PCO2 and/or [H2CO3]

[CO2 CT] or [HCOi]

mixed vs respiratory

ACT or BE/D

acute vs chronic respiratory

pH

Ni

9

t10

8 (9 or 10) 11 (12 or 13)

N12 13

Figure 3. Flow chart for differential diagnosis of acid-base imbalance (see fig. 4)

Key to numbers in chart:

1234

56

78

910

1112

13

normal acid-base

resp acid, comp (met alk, comp)

resp alk, comp (met acid, comp)

met acid, acute

met acid, partial comp

met alk, acute

met alk, partial comp

mixed resp acid and met acid

resp acid, acute

resp acid, chronic or partial comp

mixed resp alk and met alk

resp alk, acute

resp alk, chronic or partial comp

95
The same diagnostic area numbers in figure 3 are used for diagnosis with the evaluation

diagram as exemplified in figure 4. The "average line" of compensation is shown rather than

overlapping "areas."

6 8 6 9 7 0 7 1 7 2 7 3 7 4 7.5 7 6 7 7 7 8

' I I I I I I I I I I

I I I I I I I I J I I I I I I L I I I I I I I I I 1 I I I I III I I I I

PH

[H ) nmol/l 159

IHCO,

I M I I

I M F II

I I ! 1^1 I I I

I M M I M I II I I M I

I I IT| I I I I | I I I 6.8 6.9 7.0 7 1 7.2 7.3 7 4 7.5 7 6 7 7 7.8

Ralio

63

12 i IM,CO,l 1

ACID-BASE BALANCE EVALUATION DIAGRAM (1974 REVISIONl

Figure 4. Acid-base diagram with lines for diagnosis and "confidence limits" (see fig. 3)

3. Therapy

The physician should utilize the laboratory data in the management of a patient with

acid-base imbalance. Equation (15), based on an extracellular fluid volume of 25 percent of

body weight, is used as a "guide" to the parenteral therapy of the Bicarbonate Deficit or Bicarbonate Excess component.

Negative answer denotes: ml 1/6

molar sodium lactate (or bicarbonate, acetate, citrate, gluconate,

etc) for Bicarbonate Deficit

(metabolic acidosis or "Base

Deficit")

1.5 X kg X ACT (or "BE/D") = <! OR (15)

Fositive answer denotes: ml 1/6

molar ammonium chloride for Bicarbonate Excess (metabolic alkalosis

or "Base Excess")

The 1.5 "therapy factor" applies to 1/6 molar solutions— about 166 mmol of cations and of

anions (see table 2 for factors for other solutions)! It is best for the physician to administer half of the calculated volume and reevaluate the patient before therapy is continued [1 ,12].

96
Table 2. Factors for parenteral fluid therapy.

9(11 OULUIITTflN 1 1 Uli lltlllu 1 / 1 mmn IMMKJ 1/1 1 / 1

"1/6 molar" 166 1.5 4 2% NaHCOs 500 0.5

5 0% NaHCOg 595 0.4

Gastric or #3

(for ml)

70 3.5 7. 5% NaHCOa

(50 ml "ampoul")

892 0.28.

2.14% NH4CI 400 0.6 8. 4% NaHCOa 1000 0.25

(40 ml "ampoul")

4. C02rrec°t-02-Slide

The pH, Pco2, and P02 values are usually determined at 37 °C with the instruments

available in the clinical laboratory. This slide rule (fig. 5) is used to determine the

respective values at 37 °C if they were determined at other temperatures, e.g., water bath

inaccurate or under special conditions. It, is also used to determine the values "corrected"

to the temperature of the patient.

C02RREC*1-02-SLIDr

Designed 1974 By Harry F. Weisbcrg, M D.

TEMPERATURE °C

37»C

40

lihlililililijH|li|ili|ilili|ilililili{ihlilili|iliTililil)lilililil ' I I I I

I

I lll|lll[|llll||lll|llf -

I I I I [I I I l|l

6 18 20 25 30 35 40 45 50 60 70 80 90 100

TEMPERATURE CONVERSION

°F 50 55 60 65 70 75 80 85

i lilili|i

| i|ililil| lilililiKlilil|i|i'lilili| ih,

"C 10 15 20

"F 85 90 95 100 105 110 115 °F

TEMPERATURE "C

37°C

45 40-W-3S 30 25 20 15 10 ^

ii|iiii|iiii]iln|iili|lili|liil|ihl|i(ii|iili(li(i|iiilp^ (WHOLE BLOOD) 7 40 7.50 7.60 7.70 7 80 7 90 8.00

On respective scale, set deter- mined value opposite temperalure of test (usually 37 C), opposite temperature ol patient, read "corrected" value

for PcO; or pH,

ACT or BE/D ®

P02 ("FINAL"

-25 0 +25

kii|i||if^ nil liii|ini|iii{|iii CORRECTION) ;o 80 90 100

*J^Set PO; from -c^ opposite ACT or BE/D; opposite :ero, read

"final" corrected Poj.

Delermine % 0; saturation on iT; from "corrected" Po,

value Xl- ~

C 1

SET _ I CONDITIONS I "7 ( FOR Poi J DHERMINATION

Poi CORRECTED

Poj DETERMINED

P02»T37°,7.«I,OBE/D % 0; SATURATION

I I I I I I I I

[

f ll I

I M ll|llll|l l lf|f 10 15 20 25 30

37°C

10 1 5 20 25 30 3^''*^ *° *^

lilllllilllllilllilililil Jil|{ili{llllili|lii|ili|lll|ll||ll|lli|lil|^

|iiii|iiil|iiil|nil|iiil|iiii|iiii|iiil|iiilp^

6.6 6.7 6.8 6.9 7.0 7.1 7.2 7.3 7.4 7.5 7.6 7.7 7.8 7.9 8.0 25 30 35 40 45 50 60 70 80 90 100 150 200 250 300 350 400 500 600

I I

I I I

I

I

I I

I I

I IIm!||I|I|II|! I I 1 I I I II ll Illllllllllltl|llll1lllllnillllllllllll I I I I llllll I I I I I I II I I II I I IliiiliiiilnilliiiJiiN

I

I

I lll|llli|ini|llii|iH>| , I li|>lll|llli|lnl|liil|i<ii|Mil|.lil|Mii|ini|i|,,i|i||| I , , I

I

I lll|llll|illi|liii|.ili|iltl|iiM| 25 30 35 40 45 50 60 70 80 90 100 150 200 250 300 350 400 500 600 mmHg

I tjjn ll j|ij|llilli|i|^l

I

I j

[ M|ilii'iilii|iljiiiLijij|imjiiiiliii|l J

j

l|li|il^ I I I I I I I I 1 1 1 1 1 1 liiiiliiiiliiijliiiiliiiij

I

1

1 1 1

1

40 50 60 95 96 97 98 98.5 99.5 99.6 99.7 99.8 99.9 %

Figure 5. C02rrec°t-02-Slide.

It is imperative to correct pH, PCO2, and P02 values to the patient's temperature

rather than assuming 37 °C as for the thermostated water bath. Random patients were tested

with an IVAC Model 811 electronic thermometer (°F) with the oral probe placed into the

sublingual pocket for about 15 seconds. The accuracy of these thermometers is 0.15 °F in

contrast to 0.5 °F for glass (mercury) thermometers. Figure 6 shows the range of temperatures

found in 383 patients--94. 5 to 101.7 °F (or 34.8 - 38.5 °C); the average temperature was

97.6 °F or 36.5 °C. The "discrepancy" will be greater with patients subjected to hypothermia

or those with high fevers.

Equation (16) is utilized on the C02rrec°t-02-sl ide to calculate the pH of whole blood at the patient's temperature (pHo^.) from the pH value as it was determined, usually at 37 °C

(pH o) [13].

^ 37

WB pHo^ = pH o + 0.0147 (37 - °t) . (16)

97
Since plasma or serum [14] and cerebrospinal fluid contain less protein, the factors are different and the respective formulas are given in eqs (17) and (18).

P/S pHo^ = pH o + 0.012 (37 - °t)

^ 37

CSF pHo^. = pH c + 0.003 (37 - °t)

t 37

(17)

(18)

90"

80"

70 —

60-

50-

40-

30^

20 —

10 Figure 6. Range of temperatures

found in 383 patients.

94.1.'94.6- 95. 1^95.6-' 96.1 -'96. 94.5 95.0 95.5 96.0 96.5 97.0 97.5 98.0 98.5 99.0 99.5 100.0 100.5 101.0 101.5 102.

0

•F

In addition, the actual P02 of the patient can be determined if the percentage oxygen

saturation ("corrected") and the pH at the actual temperature of the patient are known.

Finally, the C02rrec°t-02-sl ide allows one to correct the P02 to the temperature of the

patient, and to correct P02 values to 37 °C, 7.40, and zero ACT (BE/D) for entrance into

the oxyhemoglobin dissociation curve. The Po2/% O2 saturation scale (F) on the slide rule

applies to "adult" hemoglobin A; the oxyhemoglobin dissociation curve of the newborn is different [15]! Other factors (2,3-DPG, type of hemoglobin, etc.) and various disease

states can cause a shift of the position of the oxyhemoglobin dissociation curve.

The A (temperature) and B (pH) scales are used to set the conditions at which the test was done or to conditions desired for calculated results. The C scale shows P02 correated

to 37 °C and 7.40, whereas D scale shows determined P02; the E scale is used to correct P02

for existing changes in ACT or BE/D. The F scale is the "normalized" (37 °C, 7.40, and zero ACT or BE/D) oxyhemoglobin dissociation curve (Po2/% O2 saturation) up to 600 mmHg and

99.9 percent oxygen saturation.

Table 3 compares temperature corrections for various hypothetical conditions as calculated

by different nomograms or slide rules.

.6-*97.1-'97.6-'98.1-'98.6-' 99. 1-' 997^^*100. l-'lOO.eilOI. 1^101 .1'

98
Table 3. Comparison of corrections of pH, PCO2,

and P02 to temperature of patient.

I, If patient's temperature is

A. pH 7.40^

C02rrec°t-02-slide

Severinghaus Calculator [16] , Siggaard-Andersen Nomogram [17]

Kelman & Nunn Nomogram [18]

B. Pco? 40 mmHg^

C02rrec°t-02-slide

Severinghaus Calculator [16]

Siggaard-Andersen Nomogram [17]

Severinghaus Factors [19]

Kelman & Nunn Nomogram [18]

Greenburg & Moulder Nomogram [20]

Rattenborg [25]

C. P0 7 90 mmHg^

C02rrec°t-02-slide

Severinghaus Calculator [16]

Severinghaus-Astrup Nomogram [19,21,22]'

Severinghaus Factors [19]

Severinghaus Factors [23] , Kelman & Nunn Nomogram [18]

Kelman & Nunn Factors [18]

Rattenborg [25]

Instrumentation Laboratory Nomogram [26]

II.

If Patient at 34°C

C02rrec°t-02-slide

Severinghaus Calculator [16]

Siggaard-Andersen Nomogram [17]^

Severinghaus Factors [19]

Kelman & Nunn Nomogram [18]

Greenburg & Moulder Nomogram [20]

Rattenborg [25]

III.

A. If Patient at 34°C

C02rrec°t-02-slide

Severinghaus Calculator [16] , Siggaard-Andersen Nomogram [17]

Kelman & Nunn Nomogram [18]

B. If Patient at 34°C (P02 = 90 mmHg)

C02rrec°t-02-slide 113

Severinghaus Calculator [16] 113

Severinghaus-Astrup Nomogram [19,21,22]^ 120

Kelman & Nunn Factors [18] 115

Instrumentat ion Laboratory Nomogram [26] 112

^ As determined on laboratory instrumentation

' at 37°C. ' ^ Modified from 38° to 37°C by substracting ' 0.02 pH.

20°C 30°C 35 C 40°C 45°C

7.65 7.50 7.43 7.36 7.28

7.67 7.51 7.43 7.36 7.29

7.66 7.50 7.43 7.35 - 7.65 7.50 7.43 7.36 -

17.5 28.5 36.5 46.5 59

17.2 28.8 36.7 45.5 53.3

18 29 36.5 45

29.7 36.7 45.5 56.3

19.2 29.6 36.8 45.6 - 18 29 36.5 45 56

19.5 30 37 46 57

00 D 1 ou 1 U/

LI . C / 0 111 ^ 34 60 80 105 140

27.5 54.9 78.3 111.8 158

1 53.2 79.5 105.8 132

36 61.2 81 103.5

60.8 80.7 105.9

27 55 79 - - 59 79 105 140

For PCO2 val ues^

30 40 50 60 70

26 34.6 43.3 52 60.5

26.2 35 43.6 52.5 61 27 35 43 51 59

26.4 35.2 44 52.8 61.6

26.

1

34.8 43.5 52.2 60.9

26 34 44 52 61 26.5 35.5 44 53 61

For pH values^

00 7. 10 7.20 7. 30 7.50 7.60

045 7. 145 7.245 7. 345 7.545 7.645

040 7, 142 7.243 7. 348 7.550 7.650

045 7. 140 7.245 7. 340 7.540 7.640

045 7. 145 7.245 7. 345 7.545 7.645

101 91 82 65 58

102 91 81.5 65.5 58.3

106 95 86 68 60

102,5 92 82.5 66 59

100 90 80 65 58

Modified from 38° to 37°C by adding 2 mmHg.

Using 96% saturation line.

Radiometer Chart No. 984-204.

99
By definition, the P^o (or T50) value is the partial pressure of oxygen (corrected to

37 °C, pH 7.40, and BE/D (or ACT) of zero) corresponding to 50 percent oxygen saturation;

normal value is about 26.6 mmHg or 3.5 kPa. The methods to obtain the P50 value are timeconsuming. Canizaro, et al. [24] have developed a nomogram which allows estimation of the

P50 value from a single venous blood sample on which the P02 [oorreated to patient's body

temperatuve andpH) and the percent oxygen saturation are determined. The estimated P50

values compare favorably (r = 0.92) with the actual P50 values if the oxygen saturation

(venous blood) is between 25 and 75 percent.

References

[1] Weisberq, H. F., A better understanding of anion-cation ("acid-base") balance,

Surg. Clinics N. Amer. 39, 93 0959).

[2] Weisberg, H. F., The Henderson-Hasselbalch equation is simple. Scientific Exhibit,

ASCP and CAP Annual Meeting, Seattle, Washington (1961).

[3] Weisberg, H. F., Water, Eteatrolyte, and Acid-Base Balance; Eormal and Pathologic

Physiology as a Basis for Therapy, 2nd ed. , 533 pp. (Williams & Wilkins, Baltimore,

MD 1962).

[4] Weisberg, H. F., pH and CO2 content; electrolyte and acid-base balance. Commission

on Continuing Education, Check Sample CC-36, (Dec, Amer. Soc. Clin. Path., 1-33 (1965).

[5] Weisberg, H. F., Calculator for the Henderson-Hasselbalch equation, in Manual for

Procedures for the Applied Seminar on Clinical Pathology of Respiratory Diseases

,

F. W. Sunderman, ed. , pp. 239-245 (Institute for Clinical Science, Inc., Philadelphia,

PA 1972).

[6] Weisberg, H. F., Water, electrolyte, acid-base, and oxygen, in Todd-Sanford's Clinical

Diagnosis by Laboratory Methods, I. Davidsohn, and J. B. Henry, eds., 15th ed., pp. 772-

803 (W. B. Saunders Co., Philadelphia, PA 1965).

[7] Woodbury, J. W. , Regulation of pH, in Physiology and Biophysics, T. C. Ruch, and

H. D. Patton, eds., 19th ed., pp. 899-934 (W. B. Saunders, Philadelphia, PA 1965).

[8] Filley, G. D., Acid-Base and Blood Gas Regulation, (Lea and Febiger, Philadelphia,

PA 1971).

[9] Davenport, H. W. , The ABC of Acid-Base Chemistry, 6th ed., Rev. (U. Chicago Press,

Chicago, IL 1974).

[10] Rispens, P., Significance of Plasma Bicarbonate for the Evaluation of H^ Homeostasis,

Thesis, U. of Groningen, Netherlands (1970).

[11] Rispens, P., Zijlstra, W. G., and Van Kampen, E. J., Significance of bicarbonate for the

evaluation of non-respiratory disturbances of acid-base balance, Clin. Chim. Acta, 54,

335 (1974).

[12] Weisberg, H. F., Parenteral fluid therapy in adults, in Current Therapy 1975,

H. F. Conn, ed. , pp. 421-433 CW. B, Saunders Co., Philadelphia, PA 1975).

[13] Rosenthal, T. B., The effect of temperature on the pH of blood and plasma in vitro,

J. Biol. Chem. 173 , 25 (1948).

[14] Siggaard-Andersen, 0., The Acid-Base Status of the Blood, 2nd ed., (Williams and Wilkins,

Baltimore, MD 1964).

[15] Nelson, N. M., et al., A further extension of the in vivo oxygen-dissociation curve for

the blood of the newborn infant, J. Clin. Invest. 43, 606 (1964).

100
[16] Severinghaus, J. W., Blood gas calculator, J. AppZ. Physiol. Z\_, 1108 (1966).

[17] Siggaard-Andersen, 0., Titratable acid or base of body fluids, Ann. N.Y. Aaad. Sai. 133 , 41 (1966).

[18] Kelman, G. R. and Nunn, J. F., Nomograms for correction of blood P02, PCO2, pH and base excess for time and temperature, J. Appl. Physiol. 21_, 1484 (1966).

[19] Severinghaus, J. W., Blood gas concentrations, in Handbook of Physiology, Section 3^, Respiration, 2, 1475 (1965).

[20] Greenburg, A. G. and Moulder, P. V., Temperature coefficients for PCO2 and pH in whole

blood. Arch: Surg. 91, 867 (1965).

[21] Astrup, P., Engel , K., Severinghaus, J. W., and Munson, E., The influence of temperature

and pH on the dissociation curve of oxyhemoglobin of human blood, Sound. J. Clin. Lab.

Invest. ]]_, 515 (1965).

[22] Severinghaus, J. W., Oxyhemoglobin dissociation curve correction for temperature and pH

variation in human blood, J. Appl. Physiol. J_2, 485 (1958).

[23] Severinghaus, J. W. , Use of single-line nomograms, in Biology Data Book, 2nd ed., 3^, 1875 (1974).

[24] Canizaro, P. C. et al. , A technique for estimating the position of the oxygen-hemoglobin

dissociation curve, Ann. Surg. 180 , 364 (1974).

[25] Rattenburg, C. C, Blood Gas Evaluator (Graphic Calculator Co., Barrington, IL, 1970).

[26] Instrumentation Laboratory, Inc., Instruction Manual for Model 182 CO-Oximeter, revised

edition (1972).

101

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

AIDS FOR EVALUATION OF ACID-BASE IMBALANCE-DIAGRAMS,

NOMOGRAMS, AND SLIDE-RULES

Harry F. Weisberg

Division of Biochemistry

Mount Sinai Medical Center

Milwaukee, WI 53201, USA

In 1962 I had accumulated 12 acid-base diagrams spanning the period, 1921-1959 (table

1). Since then I have been able to add 7 time charts (1917-1974), 17 nomograms (1923-1974)

[45,46,162,166-180]^ , 8 mathematical factors (1957-1967), and 12 slide rules (1958-1974)

[181-192]. The original 12 diagrams have expanded to 39 for the period 1921-1959.

Table 1. Evolution of acid base diagrams. Reproduced with permission

of the publisher, Williams and Wilkins [1].

Vertical Horizontal

coordinate coordinate Other

(ordinate) (abstcissa) coordinates

Author year Value Log

scale?

Value Log

scale?

Value Log

scale?

Van Slyke 1921 CT^ PC02 PH^

Peters 1923 CT Yes PC02 Yes pH

Peters & Van Slyke 1931 HCOs"'^ or CT

. PH

Shock & Hastings 1934 HCO3" Yes pH PC02^ Yes Shock & Hastings 1935

Davenport 1947 HCO3" pH PCO2 Yes Weisberg 1950

Westerfeld

Weisberg

1953^

1956' CT pH PCO2 Yes

Winters et at. 1958 CT Yes pH PCO2 Yes Weisberg 1961 CT [H^]^ PCO2 Yes Singer 1951 BB^ PCO2 pH and

CT Yes

Peirce et al. 1959 PCO2 pH BB Yes Brewin et al.

Astrup

1955^

1956/ PCO2 Yes pH

Astrup et al. 1959 PCO2 Yes pH BB, BE,?

and SB

CT--carbon dioxide content; pH-- reciprocal of hydrogen ion concentration;

HC03--bi carbonate; dpco2--partial pressure of carbon dioxide; ^[H+]--hydrogen

ion concentration; fBB--buffer base; %E—base excess; and 'iSB--standard bicarbonate.

Figures in brackets indicate literature references at the end of this paper.

103
Figure 1 illustrates the fecundity of the "sixties" for the production of acid-base

aids. The earliest diagram I have been able to unearth was published in 1914 by Christiansen,

Douglas and Haldane [3]. Table 2 is the latest accumulation of 105 diagrams (1914-1973) used

in the evaluation of acid-base imbalance. The references listed in the 1° column are the

originals I have been able to find, whereas those references in the 2° column are corrections,

modifications, and variations based on the originals.

50-1

40-

30-

20-

Diagrams D 105

Nomograms 17

Slide Rules ^ 1

2

im.

1911-20 1921-30 1931-40 1941-50 1951-60 1961-70 1971-

Ftgure 1. Histogram of acid-base aids. Modified from Weisberg [2].

Key to Symbols :

Table 2. Comparison of 105 acid-base diagrams. Modified

from Weisberg [1 ,2]

.

Triaxial coordinates

aH^ Hydrogen ion activity

A Acid, expressed as ± Pco2 (Kintner)

Alk. Res. Alkaline reserve (McClendon)

B Base, expressed as ± base (=G=BE) (Kintner)

BB Buffer base

BE Base excess

BEj^g Base excess, whole body (Russell st at.)

BE/D^^P Base excess/deficit of extracellular fluid (S.iggaard-Andersen)

BE/Dy. c Base excess/deficit for extracellular fluid with hemoglobin at

5 g/dl CRooth)

C^7 Hydrogen ion concentration (OpH)

CSF Cerebrospinal fluid (McClendon)

CT Total carbon dioxide (as mmol/1 or vol %, etc.)

CY CO2 capacity (Goldberger)

aA-B a Acid-base (Astrup)

104
Table 2. Comparison of 105 acid-base diagrams (continued).

ECF Extracellular fluid

[H+] Hydrogen ion concentration

h\o Hydrogen at 40 mmHg (Frumin)

[HC03~] Bicarbonate ion concentration

[HCOs'l^o Bicarbonate at 40 mmHg (Siggaard-Andersen) ; as SB [H2CO3] Carbonic acid (and dissolved CO2) concentration

LA Lactic acid

N Normality of h"*" (McClendon)

"N" Normality of titration of alkaline reserve (McClendon)

EO2] Concentration of (dissolved) O2

O2 CT Total oxygen (vol %)

PCO2 Partial pressure of carbon dioxide

"PCO2" Assumed from % CO2 in a alveolar air (McClendon)

pHi^o pH at 40 mmHg (Kappagoda et at.)

PR Serum protein (Stadie et at.)

R Ventilation ratio (Kim et al.)

S Solubility coefficient for CO2 in plasma at 38 °C

SB Standard bicarbonate at PCO2 40 mmHg, 38 °C, Hb oxygenated

SBy^ito Standard bicarbonate at pH 7.40 (Van Slyke)

t" Temperature (Stadie et al.) (Lenfant)

Ti+o Bicarbonate at 40 mmHg in vivo (Collier et al.)

TB Total buffer (Coats)

VR Ventilation ratio

WB Whole body

Reference Values Plotted On Year Author(s) 1° 2° Ordinate Log? Abscissa Log? Others Log?

1914 Christiansen et al. 3 4,5 CT PC02

1916 Lewis et al. 6 7 pH CT

1916 McClendon 8 PCO2 pH

CT

1917 McClendon et al. 9 PCO2 Yes Eh+] Yes

1917 McClendon 10 pH

rH+] Yes

PC02 Yes [HCOi] Yes

1917 McClendon 10 CT Alk. Res. Yes

1917 McClendon et al. 11 PCO2 Yes

pH Yes Alk. Res. Yes

1917 Parsons 12 13 pH PC02

105
Table 2. Comparison of 105 acid-base diagrams (continued).

Values Plotted On Year Author(s)

Reference

1° 2° Ordinate Log? Abscissa Log? Others Log?

1917 Parsons 12 13-15 PC02

1918 McClendon 16 "PC02" Yes [H^]

pH

Yes

1918 Straub & Meier 17 CT PCO2

1919 Haggard &

Y. Henderson

18 CT PC02

1920 Henderson 19 EHCO3] [H^]

1921 Henderson 20 [HCO3] pH

1921 Van Slyke 21 22-24 CT PC02

1921 Van Slyke 21 [HCO3] pH

1922 Van Slyke et al. 25 CTLHCO3J pH

1922 Doisy et al. 26 pH [HCOi]

1923 Peters et al. 27 28 [HCOi] Yes [H2CO3] Yes

1923 Peters 29 CT Yes PC02 Yes

1923 Peters 29 24 CT Yes PC02

[H2CO3]

Yes

Yes

1923

1924-

Cullen & Jonas

^ 25 Hastings et al. j

30

31 32,33,193 CT pH

1924 Bock et al. 28 23 [HC03"J pH

1924 Bock et al. 28 CT pH

1925 Murray & Hastings 34 [HCO3] Yes pH

1925 Austin & Cullen 35 [HCOi] PC02

1925 Austin & Cullen 35 36 CT pH

CSF "N" Yes

pH Yes [H2CO3]

[H2CO3]

pH

pH

PC02

Yes

[H2CO3]

[C1-]

P02

Hb02

pH Yes [H2CO3]

Yes

Yes

[H2P0^] Yes [HP04^"] Yes [PO^^"] Yes [Ca2+] Yes [COs^"] Yes

[H2CO3]

Pco2 Yes

106
Table 2. Comparison of 105 acid-base diagrams (continued).

Reference Values Plotted On

Year Author ( s

)

1° 2° Ordi na fp 1Luyna?:

1LOgnn?: u uners Log f

1925 Stadie et at. 37 [HC03]

r K pH

OI 4- 0 t pH [HCOi]

[H2CO3]

PC02

Yes

Yes Yes 1927 Eisenman 38 CT Yes PCO2 Yes

1927 Dill et al. 39 CT Yes PCO2 Yes

1928 Henderson 36 CT

[H2CO3J

LO2J

PCO2 -

1928 Henderson 36 PCO2 pH CT Yes

1931 Peters & Van Slyke 23 40 CT Yes PCO2 Yes pH

1 ))o 1 nastings & Steinhaus°

/n 4^-46 PCO2 Yes pH [HCOi] Yes

1932 Douglas & Havard 47 pH

CTL 1 PCO2 -

McCl endon° 48 pH rC02 Yes [HCOi] Yes

1947 Davenport 49 45,50-54 [HCOi] pH PC0 2 Yes 1948 Clark 55 56-61 ,165 [HCO3] PCO2 pH Yes

1 yb 1 Singer BB PCO2 CT

pH Yes

Yes 1954 Astrup 1 63,69,77,

164

PCO2 Yes pH -

1955 Cranston et al. 64 [H2CO3] -

1956 Roberts et at. 65 PCO2

pH

CT -

1 yoD roppei et au. 00 C7 CO 0/ ,00 rC02 L

\ yoD MS irup 03 rC02 Yes pH SB Yes

1958 Winters et al. 71 72 CT Yes pH PCO2 Yes

1958 J0rgensen 73 SB Yes PCO2 Yes pH

1959 Go! dberger 74 CT

1 CY

pH

1959 Peirce 75 PCO2 pH BB Yes

107
Table 2. Comparison of 105 acid-base diagrams (continued).

Reference Values Plotted On

Year Author(s) 1° 2° Ordinate Log? Abscissa Log? Others Log?

1959 Astrup 76 PC02 Yes pH BB

SB

A A D AM-b

Yes Yes

I es

1960

1 you

Astrup et at. ] Si ggaard-Andersen ]77

7Q/o

46,52,

7Q QQ /ts-oy

PC02 Yes pH BB

CDBE

Yes

Vqc 1 es Yes

1 ybu de la Huerga onyu,

91

L. 1 pH Vac rC02

1961 Wei s berg 92 CT rH''"T pH Voc I cb

PCO2 Yes

1961 Lenfant 93 [H2CO3] Yes pH CT Yes

1 951 Lenfant 9o ru rr\ 1 LH2CU3J Vat- Yes PC02 t

1961 We is berg 92 1 CT pH

[HCO3]

LH2CO3J

Yes

PCO2 Yes

1962 Campbel 1 94 95,96 PCO2

Ln2'-U3J

[HCO3]

pn Yes

1 Co

1 ytD£ Campbel

yf Q7y / r CU2

[H2CO3]

1 Co Yes EHCOg] 1 CO un[H+]

pH

1 Co Yes

1962 Nunn 98 [HCO'a] PC02 pH Yes

1963 Si ggaard-Andersen 81 99 PC02 Yes pH [HCOi] Yes

1963 Weisberg (revised) 100 101-103 CT dH[H+]

[HCO3]

Yes

Yes

PC02 Yes

1963 Robin 104 PH

LH2CO3J

PC02 [HCOi] Yes

i ybo rink a iNanas t Ub

pH Yes

PC02 J nous J VR

Vqc ICS

1 964 Wni tehead 1 06,

107

rC02 Yes Vqc I es

1964 Coats 108 [HCOi] [H2CO3] pHTR Yes

1964 Da r row 109 [HCO^] PC02

[H2CO3J

pH

[HCO3]

Yes

Yes [H2CO3]

1965 Woodbury 110 15,52,

111-114

[HCOi] PH^

[H+J Yes PC02 Yes

108
Table 2. Comparison of 105 acid-base diagrams (continued).

Reference Values Plotted On

Year Author(s) 1° 2° Ordinate Log? Abscissa Log? Others Log?

1965 Schwartz {et al.

)

115,

116

13,117,

118

[HCOi] PC02

1965 Schwartz (et al.

)

115,

116

117-119

pH Yes

PC02

1965 Mithoefer et al. 120 121 [HCOi] Yes PC02 Yes [H^]

pH Yes

1965 Owen et al. 122 PC02 [H^]

pH Yes EHCO3]

1965 El kinton 123 124 [HCO'g] Yes PCO2 Yes pH

1965 Siesjo & Ponten 125 PC02 Yes [HCOs]

S-PC02

SB Yes

1966 Instrumentation

Lab'

126 PC02 Yes pH CT Yes

1966 Lennon & Lemann 14 [HCO3]

1966 Lennon & Lemann Id 1 c. / , 1 CO , 140

PC02 iHCOi] -

1966 Kim et al. 129 CT PC02 pH

RO2 CT

Yes

1966 Stinebaugh & Austin130 131,132 PC02 CT pH Yes

1966 Young 133 [HCO3] Yes pH

[H+] Yes PC02 Yes

1967 Cohen 134 135-138 [H+]

pH Yes

PCO2 [HCOi] Yes

1967 Albert et al. 127 PCO2 BE -

1967 Albert et al. 127 139 PC02 -

1 / n 1 Uci L c: u CLL' . 127 PC02 nHpn

1967 Siggaard-Andersen 85 89,141 PC02 Yes pH Yes ^^(ECF)

[HCOi] 40

Yes

Yes

1967 Kintner 142 CT pH

[H+] Yes

PC02 A

B

Yes Yes

Yes

1967 Winters 57 59 PC02 [HCOi] pH Yes

1967 Winters 57 59,143 PC02 BE pH Yes

1967 Winters 57 59 BE PC02 pH Yes

109
Table 2. Comparison of 105 acid-base diagrams (continued).

Reference Values Plotted On Year Author(s) 1° 2° Ordinate Log? Abscissa Log? Others Log?

1967 Srouji 144 BE PC02 _

1968 Engel et at. 13 BE PC02

1968 Goldring al. 14(1 1 "tO pH EHCO3J PCO2 Yes

1969 Li & Holder 147 pH SB PC02 Yes

1969 Nunn 112 PC02 Yes pH

[H+] Yes

1969 Weisberg 148 149-153

' CT

[HCO3]

Yes

Yes

PCO2

[H2CO3]

Yes Yes

[H2CO3J

1970 Gilbert & Auchincloss

PC02 -

1970 Heisler & Schorer 155 89 [HCOi] pH PCO2

BB

BE

Yes

Yes Yes

1970 Kappagoda et at. 156 PC02 Yes pH pHitO

1970 Rooth 157 PC02 ^^Hbs

^^Hbs

pH Yes

1970 Rooth 157 PC0 2 pH Yes

8°Hb5

1970 Blair 158 PC02 BE

LA

pH -

1971 Stephens 159 CT pH Yes PCO2 Yes

1971 Slonim & Hamilton 160 [HCO^J PCO2

[H2CO3]

VR Yes

pH

1972 Kintner 161 [HCOi] pH PCO2

BB Yes

1972 Russel et al. 162 PC02 Yes pH ^^WB Yes

1973 Visser 163 PC02 Yes pH Yes

110
References

[1] Weisberg, f/ater^ Electrolyte and Acid-Base Balance, Znd edition, 533 pp. (Williams and

Wilkins Co., Baltimore, 1962).

[2] Weisberg, InteTyretation of Electrolyte, Acid-Base, and Oxygen Imbalance, 4th edition,

3rd revision (Private mimeographed, 1975).

[3] Christiansen, Douglas and Haldane, J. Physiol. 44, 244 (1914).

[4] Chambers and Kleinschmidt, J. Biol. Chem. 55, 257 (1923).

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[6] Lewis et al. , Brit. Med. J. 2, 517 (1916).

[7] Astrup, Ugeskr. f. Laeg. 116 , 758 (1954).

[8] McClendon, J. Biol. Chem. 28, 135 (1916-17).

[9] McClendon et al. , Carnegie Inst. Washington Vl_, 23 (1917).

DO

[11

[12

[13

[14

[15

[16

[17

[18

[19

[20

[21

[22

[23

[24

[25

[26

[27

[28

[29

McClendon et al. , J. Biol. Chem. 30 265 (1917).

McClendon et al. , J. Biol. Chem. 31 519 (1917).

Parsons, J. Physiol. 51 440 (1917).

Engel et al., J. Appl. Physiol. 24-, 288 (1968).

Lennon and Lemann, Ann. Int. Med. 65, 265 (1966).

Ichiyanagi et al., Resp. Physiol. 7, 310 (1969).

McClendon, J. Amer. Med. Assoc. 70, 977 (1918).

Straub and Meier, Bioahem. Z. 89, 156 (1918).

Haggard and Henderson, Y., j. Biol. chem. 39, 163 (1919).

Henderson, J. Biol. Chem. 41, 401 (1920).

Henderson, J. Biol. Chem. 46, 411 (1921).

Van Slyke, J. Biol. Chem. 48, 153 (1921).

Austin et al. , J. Biol. Chem. 54, 121 (1922).

Peters and Van Slyke, Quantitative Clinical Chemistry, Vol. I, Interpretations

(Williams and Wilkins Co., Baltimore, 1931).

Van Slyke, Bull. N.Y. Acad. Soi. ]0_, 103 (1934).

'

Van Slyke et al. , J. Biol Chem. 53, 277 (1922).

Doisy et al., J. Biol. Chem. 54 305 (1922).

Peters et al. , J. Biol. Chem. 55, 709 (1923).

Bock et al., J. Biol. Chem. 59, 353 (1924).

Peters, J. Biol. Chem. 56, 745 (1923).

Ill
[30] Cullen and Jonas, J. Biol. Chem. 57, 941 (1923).

[31] Hastings et al. , J. Clin. Invest. 1, 25 (1924-25).

[32] Westerfeld, Fundamentals of Aoid-Base and Electrolyte Bioohemistry (Syracuse, N mimeographed edition, 1953).

[33] Weisberg (1956).

[34] Murray and Hastings, J. Biol. Chem. 65, 265 (1925).

[35] Austin and Cullen, Medicine 4, 275 (1925).

[36] Henderson, Blood (Yale Univ. Press, New Haven, 1928).

[37] Stadie et al., J. Biol. Chem. 66, 901 (1925).

[38] Eisenman, J. Biol. Chem. T\_, 587 (1927).

[39] Dill et al., J. Biol. Chem. 73, 251 (1927).

[40] Miller, Acta Med. Scand. 165, Suppl . 348 (1959).

[41] Hastings and Steinhaus, Amer. J. Physiol. 96, 538 (1931).

[42] Shock and Hastings, J. Biol. Chem. 104, 585 (1934).

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[45] Rispens, Thesis, U. of Groningen (1970).

[46] Mohler (1971), personal communication.

[47] Douglas and Havard, J. Physiol. 74, 471 (1932).

[48] McClendon, Science 81, 569 (1935).

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Chicago, 1958).

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Chicago, 1969).

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[56] Filley, John Hopkins Med. J. 120 , 355 (1967).

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112
Arbus, Canadian Med. Asen. J. 109, 291 (1973).

Ayres, i?esp. Care 19, 280 (1974).

Singer, Mer. J. Med. Soi. 221 , 199 (1951).

Brewin et at., Guy's Hoep. Reports 104, 177 (1955).

Cranston et at., J. Physiol. 127 . 380 (1955).

Roberts et al., J. Clin. Invest. 35, 261 (1956).

Poppell et al. , J. Lab. Clin. Med. 47, 885 (1956).

Bane et al. , J. Appl. Physiol. 10^, 151 (1957).

Austin, Amer. Heart J. 69, 691 (1965).

Astrup, Scand. J. Clin. Lab. Invest. 8, 33 (1956).

Visser and Maas, Clin. Chim. Acta 5_, 850 (19G'.).

Winters et al. , J. Clin. Invest. 37_, 640 (1958).

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J0rgensen, Ugeskr. f. Laeg. ]2^, 993 (1958).

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Philadelphia, 1959).

Peirce, Trans. Amer. Soo. Artif. Int. Orgs. 6^, 240 (1960).

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Boston, 1959).

Astrup et al.. Lancet 1035 (1960).

Siggaard-Andersen, Scand. J. Clin. Lab. Invest. 1^, 311 (1960).

Rosevear (1960), personal communication.

Siggaard-Andersen, Scand. J. Clin. Lab. Invest. U, 598 (1962).

Siggaard-Andersen, Scand. J. Lab. Clin. Med. |5^, Suppl . 70 (1963).

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Goldberger, A Primer of Water, Electrolyte and Acid-Base Syndromes, 3rd edition (Lea

and Febiger, Philadelphia, 1965).

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Thompson and Mullen, J. Surg. Res. 1_0, 117 (1970).

Mason, Fluid Electrolyte and Nutrient Therapy in Surgery (Lea and Febiger, Philadelphia,

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Siggaard-Andersen, The Acid-Base Status of the Blood, 4th edition (William and Wilkins

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113
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118] Levesque (1972), personal communication.

119] Van Ypersele de Strihou and Frans, Presse Med. 75, 1797 (1967).

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122] Owen et al.. Lancet Z, 660 (1965).

123] Elkinton, Ann. Int. Med. §2, 893 (1965).

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126] 1 ns LrumciHa l i uri i_a uu r a lu ry, iiic. ^laooy.

127] Ml Deri ev au. } Ann. mv. Mea. dd, \\yo/i.

128] riciLc aiy. ^ /inn. mv. lyi&a. uoo \ i y / u j

.

129] i\i[n ev au. , v . Hpp u . rnyst-o u. ^i, looo ^i^dd^.

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145] Goldring et al., J. Clin. Invest. 47, 188 (1968).

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148] Weisberg (1969).

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150] Weisberg (1972).

151] Weisberg (1974).

115
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116
[183] Weisberg (1963).

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117
I

t

j

i
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977

THE OVERALL FIRST IONIZATION EQUATION OF CARBONIC ACID AS

RELATED TO CO2 IN GAS PHASE: A NEW pK

A. H. J: Maas

Department of Cardiology and Thoracic Surgery

University Hospital

Utrecht, The Netherlands

B. F. Visser

Department of Anesthesiology

University Hospital

Nijmegen, The Netherlands

For the evaluation of acid-base disturbances in biological fluids, the carbonic

acid/bicarbonate buffer system is in common use. In all studies this system is described by

the 1st ionization equation of carbonic acid as related to the physically dissolved carbon

dioxide in the solution: the well-known Henderson-Hasselbalch equation. As a result of

various applications, originally for calculation of the pH of plasma, later on the pCOa and

nowadays the actual bicarbonate concentration, standard bicarbonate concentration or totalCO2 concentration, this equation is written in different forms. However, little emphasis

has been placed on an exact formulation which accounts for the non-ideality of the system.

By virtue of the thermodynamic theory of equilibrium, we have shown that the 1st ionization

equilibrium of carbonic acid, which relates the bicarbonate concentration and the partial

pressure of CO2 to the pH of the solution, is preferable to the conventional HendersonHasselbalch equation [1]^.

1 . Theory

If there is equilibrium between CO2 in the gas phase and that in the liquid phase, we

can consider the overall equilibrium:

CO2 (gas) + H2O JHCO3 + H"^ (1)

of which the equilibrium constant K^^ is defined as:

aH^ • aHCOa i/HCOg aH"^ • cHCOg

^ig = = X . (2)

aH20 • fCOz aHzO • gCOz pCOz

Suffix g in Kin refers to CO2 in gas, a, f, y, g, a and p symbolize respectively, activity,

fugacity, activity coefficient, fugacity coefficient, concentration, and partial pressure.

For practical purpose we define a practical coefficient z{q:

aH20 • ^C02 aW^ • cHCOi

K{n = X Kiq = (3)

J/HCOg PCO2

^Figures in brackets indicate the literature references at the end of this paper.

119
Usually one considers the 1st ionization equilibrium as related to the sum of freely dissolved

CO2 and the hydrated CO2. The starting-point in that case is the overall 1st ionization

equilibrium, related to the dissolved carbon dioxide:

COaCliquid) + H2O j H"^ + HCO3 (4)

of which, the equilibrium constant is defined as follows:

aH"*" • aHCOg yHCOi qH"^ • CHCO3

X = = X . (5)

aHaO • aC02 aH20 • 1/CO2 ^^^^z

The suffix 1 in Ki-, refers to CO2 in the liquid phase. Now the CO2 concentration is equal

to the solubility s times pC02 minus the concentration of H2CO3:

0CO2 = S ' pCOz - CH2CO3 (6)

which equation by means of the hydration constant K^^ can be written as:

aH20 • 1/C02 • K. 0CO2 = S • pC02/(l + -) (7)

2yH2C03

or in a simplified form:

<3C02 = S . pC02/(l + a) . (8)

If we replace eC02 by S • pC02/(l + a) in equation (5) then Ki-, is changed into the conventional form:

ytiCOl • (1 + a) aH"^ • CHCO3

jC = X . (9)

aHzO • 1/C02 S • PCO2

For practice we again define a coefficient easy in operation:

aH20 • Z/C02 ' CHCO3

K{. = X = . (10)

' yHCOi • (1 + A) ' S • PCO2

To be able to choose between the two practical equations we first write them in the usual

form. We take the negative logarithm of these equations and apply the notation -log = p. This gives us for equation (10):

+ aH20 • yC02 oHCOl CHCO3

paH = pKi. - log + log =

p^i-, + log (11

)

z/HCOs • (1 + a) S ' PCO2 S • PCO2

120
which equation represents the exactly defined expression of what is now known as the HendersonHasselbalch equation and for equation (3):

^ aHjO • gC02 CHCO3 CHCO3

paH = p^ig - log + log = pi^ig + log (12)

yHCOs pCOa pCOa

which equation we propose to call the modified Henderson-Hasselbalch equation. Since this

equation contains one variable less, namely the solubility coefficient 5, this modification

of the Henderson-Hasselbalch equation is the most suitable equation for description of the

carbonic acid/bicarbonate buffer system. Assuming g-COa equals 1 (see ref [1]) and a equals

zero, it will be clear that the variability of p^ig is determined by the ratio aH20/y\]C0l

and the variability of pKi-. overmore by yC02, which coefficient strongly depends on lipid

content of the solution.

2. Experimental Section

In order to be able to apply the 1st ionization equilibrium of carbonic acid as related

to CO2 in gas,j«/e established the values of the ionization constant (pi^ig) and practical

coefficient (p?ig) over a large range of application [2].

A. pZig

The true ionization constant pKig has been determined in water using (NaHCOs + NaCl)

solutions of ionic strength 0.01 < I < 0.500 in equilibrium with (CO2 + H2) gas mixtures

applying electromotive force measurements with a hydrogen electrode cell without a salt

bridge from which the exact pKig value results. We found:

pZig (25°C) = 7.720 pZig (38°C) = 7.810 .

pZig, a thermodynamic constant, is independent of solution composition and varies with tem- perature only. For the temperature range 25-40 °C, a useful equation could be derived subtracting the equation

pZi^ = 3374J62 ^ Q_Q3272y . ^^jqj (13)

from Shedlovsky and Mclnnes [3] and

log 5q = ^-^^y^^ + 0.0148483T - 13.670 (14)

from data of Bartels [4] and Austin [5] for pure water resulting in the relationship:

pZig = pZi^ - log = 0.017877 - 1.037 (15)

where T is the absolute temperature. Our measured data agree within 0.005.

121
B. pXjg

The practical coefficient

p^ig = pXiq - log = pH - log CHCO3 + log pCOa (16)

was determined in (NaHCOs + NaCl) solutions, cerebrospinal fluid (CSF), plasma and serum, measuring pH, c?HC03 and pC02. We use the symbol ~ to indicate that pH in consequence ofjthe

residual liquid junction potential in the measurement is not exactly identical to paH'^.pK-^n

and pK{g are related to p^ig and pK{g according to the following equation

+ pXlg - pKig = pK[g - \)K{g = pH - paH . (17)

3. Bicarbonate Solutions

Using a hydrogen electrode cell including a salt bridge, we measured the practical

coefficient for various ionic strengths and found by extrapolating to infinite dilution an approximate value of the ionization constant p^ig, which differed airaa n.n05 from pKig

indicating that the residual liquid junction is very small. From these measurements, we

also derived an empirical equation for yHCOs, which in the form as recalculated by SiggaardAndersen [6] reads:

log yHCOg = -(0.50 + O.OOlAr) /f + 0.131 (18)

where at = r - 288.15 K.

As a first approximation, log aH20 decreases linearly with I according to the

equation [2]:

log aH20 = -0.0151 (19)

Combining equations (15), (18), (19) and adding 0.005 for the residual liquid junction

potential give a valid approximation of pZ|q as a function of ionic strength at temperatures

15-40 °C.

^

pK{g = 102|-404 ^ 0.017877 - 1.042 - (0.50 + O.OOIat) /i + 0.1457 . (20)

Our measured results deviate less than 0.005. For the activity coefficient yH^ , we recalculated our data and found the relationship:

log yti^ = -(0.50 + 0.001A7) /f + 0.81 . (21)

Apart from the problem of the residual liquid junction potential, the usual glass

electrode deviates from the fundamental hydrogen electrode in the pH determination [7].

122
Further, the bicarbonate determinations in body fluids are subject to error of analysis,

because carbonate and carbamate are enclosed in the present titrimetric and gasometric

analysis [8]. To be able to understand quantitatively the influence of the disturbing

"bicarbonate" fract^ions and the error of the glass electrode, we have determined the practical coefficients pz{g of a bicarbonate solution isotonic with CSF and plasma, composition

0.025 mol/1 NaHCOs + 0.135 mol/1 NaCl over the range 6-8 pH and performed simultaneously the

required pH measurements with the hydrogen electrode cell as well as with the glass electrode

cell at 25 and 38 °C (fig. 1).

7.55

7.54-

7.53

7.52-

7.51

7.50-

25° —* ^ » -^A.

\

Pi<^g(cb)— O

P^ig —

1 -pH

7.55-

7.54-

7.53-

7.52-

7.51-

7.50

25^

m &-.-fi..p.

ffl A' A A

A

pl<'g(cb,gl)---

pK^g(gi) •••••

pKjg —A

pH

6.0 6.5 7.0 7.5 8.0

—I 1 1 1—

6.0 6.5 7.0 7.5

\

\

\

Vb

8.0

7.64

7.63

7.62'

7.61

7.60-

7.59-

7.64

7.63

7.62

7.61

7.60-

7.59'

38°

S.^j»

t. < A ^ 0\p-..

pKjg(cb,gl)-'

pKjgCgl) ••

pK'

pH

6.0 6.5 7.0 7.5 8.0 eTo eTi tTo tT? 8.0

Figure 1. The uncorrected and corrected practical coefficient pZjg of (NaHCOs + NaCl) solution (j = 0.16), determined with the hydrogen electrode cell and

with the glass electrode cell as a function of pH (see text).

123
In the two graphs on the left-hand side, the overall 1st ionization coefficient, determined with the hydrogen electrode cell.^is plotted as a function of the pH. The broken line

indicates the uncorrected coefficient p^ig Ccb) , and the drawn line the coefficient pK'ig

corrected for the carbonate error. In the graphs on the right-hand side, the coefficient,

determined with a cell with the glass electrode js plotted as a function of the pH. The

broken line gives the uncorrected coefficient pK'iQ (cb, gl), the dotted line the coefficient

pK'ig (gl) corrected for the carbonate error and tne straight line the coefficient pK[q

corrected for the carbonate error and the glass electrode. These experiments affirm that

the pK'ig does not vary with the pH, as is established^by definition. On the other hand.

Si ggaard-Andersen [9J explains the variation of the pK'ig with the pH he found, by the

presence of non-dissociated NaCOs besides COl~ ions. His experiments have been carried out

with a glass electrode of the type used by us. By applying the right correction, however,

we have come to the conclusion that the variation of pZig (gl) with the pH should be

explained from the property of the glass electrode itself.

4. CSF, Plasma and Serum

Altogether, of 15 samples of CSF and 12 samples of plasma and sera the coefficient

pFig (cb, gl) was determined at 25 °C &nd 38 °C.

7.55-

7.54

7.53

7.52

7.51

7.50

CPS 25 °C

•

•

• •V•V •

* v pKig(cb,gl)

•

1 pH

6.2 6.6 7.0 7.4

Figure 2. The practical coefficient pK[g (cb, gl ) of CSF and plasma (serum) as a function of pH (see text).

124
group of p^ig (cb, gl) we have taken the mean value, indicated by a. Through it, we have

drawn a line at sight. At pH smaller than 7, p^ig (cb, gl) is constant, and at pH higher

than 7, pK{q (cb, gl) decreases. This decrease is the same as that of the discussed bicarbonate solution. From this, we may conclude that the dependence of this practical coefficient on the pH in the physiologic pH range must be due to the influence of the 2nd dissociation of carbonic acid on the bicarbonate determination and the non-ideal cell with the glass

electrode for the pH determination.

The p^ig (cb, gl) values of CSF (7.527 at 25 °C and 7.617 at 38 °C) and plasma (7.529

at 25 °C and 7.624 at 38 °C) do not appear to differ significantly, indicating that the

ionic strength of both fluids is equal. So it is also misleading to use the old values

-pK[-\ (cb, gl) = 6.13 for CSF and pK{-\ (cb, gl) = 6.10 for plasma (at 38 °C) , suggesting that

the ionic strength of plasma is higher than that of spinal fluid. For simplicity we propose

to take a mean quantitity for CSF and plasma, to be designated by the asterisk, pK*g.

In the graphs pK*^ is drawn as a function of the pH (dotted line). Finally, we deduced

from eq. (15) and the formula of mentioned dotted line [2] an approximation for pxtn valid

valid in CSF and plasma over the pH range of 6.0-7.8 pH at temperature 15-40 °C.

.I02|^, 0.017877

- 1.21813 - 0.0012(pH-7.0) - 0.0406(pH-7.0)2

The precision of the pK*g value obtained for CSF or plasma of normal subjects, is better

than 0.01 p^ig. Extreme pathological variation in the ionic strength of the plasma, for

example J = 0.160 ± 0.030, causes a variation in pKi^ of about ± 0.015 as verified with

eq. (20).

5. Conclusion

On theoretical and practical grounds, we recommend the use of the practical equation of

the overall 1st ionization equilibrium of carbonic acid as related to the gas phase for

calculations of the bicarbonate concentration (or total CO2 content) of CSF or plasma. At

pH = 7.40 and t = 37 °C the pK*n value is equal to 7.615 for CSF, plasma and serum.

References

[1] Maas, A. H. J., Visser, B. F. , van Leeuwen, A. M. and Overbeek, J. T. G. , An overall

equation for the first ionization equilibrium of carbonic acid, Pflugers Arch. 304,

20 (1968).

[2] Maas, A. H. J.^van Heyst, A. N. P. and Visser, B. F., The practical equilibrium

coefficient (p2'ig) for the first ionization of carbonic acid in solutions of sodium

bicarbonate, cerebrospinal fluid, plasma and serum at 25 and 38 °C, Clin. chim. Acta

33, 325 (1971).

[3] Shedlovsky, Th. and Mclnnes, D. A., The first ionization constant of carbonic acid,

0-38 °C from conductance measurements, J. Amer. Chem. Soo. 57, 1705 (1935).

[4] Bartels, H., and Wrbitzky, R. , Bestimmung des C02-Absorptions-koeffizienten zwischen

15 and 38 °C in Wasser and Plasma, Pflugers Arah. 27]_, 162 (1960).

[5] Austin, W. H., Lacombe, E. , Rand, P. W. and Chatterjee, M. , Solubility of carbon

dioxide in serum from 15 to 38 °C, J. Appl. Physiol. 18, 301 (1963).

[6] Siggaard-Andersen, 0., ed. , The Acid-Base Status of the Blood, 4th revised edition,

p. 34 (Munksgaard, Copenhagen, 1974).

[7] Maas, A. H. J., pH determination of body fluids with a micro glass electrode and a saturated KCl bridge in the cell, Clin. Chim. Acta, 28, 373 (1970a).

125
Maas, A. H. J., A Titrimetric method for the determination of actual bicarbonate in

cerebrospinal fluid and plasma or serum, Clin. Chim. Acta, 29, 567 il970b).

Siggaard-Andersen, 0., The first dissociation exponent of carbonic acid as a function

of pH, Scand. J. Clin. Lab. Invest. U, 587 (1962b).

126
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

TOWARDS A PHYSIOLOGIC NOMENCLATURE FOR IE VIVO DISTURBANCES

OF ACID-BASE BALANCE

Jordan J. Cohen

New England Medical Center Hospital

Boston, Massachusetts 02111, USA

In considering disturbances of acid-base equilibrium from a clinical vantage point, it

would seem axiomatic that a fundamentally physiological frame of reference must dominate

any conceptual scheme offered as an aid to diagnosis and therapy. It is abundantly clear

that both normal physiologic regulation as well as patho-physiologic disturbances of acidbase equilibrium in the living organism involve a large array of complexly interacting,

dynamic processes. Although the simplest of these processes, that of buffering, can for some purposes be conceived of in strictly chemical terms, a thorough analysis of even this

process must incorporate many features which are strictly within the domain of physiology

and which are not adequately encompassed by purely chemical notions. I refer here to such

features as compartmental ization , recruitment, endocrine influences, "slow equilibration,"

membrane transport, and the like. Add to this the physiologic "regulatory" functions of

the respiratory and renal systems, each of which is impinged upon by a veritable welter of

influences, and one can scarcely deny that true comprehension of this system poses an enormous challenge.

As has been amply discussed, this challenge cannot be circumvented by systems of acidbase analysis based exclusively on the in vitro behavior of blood [1]^. There would appear,

in fact, to be no alternative to using the empiric behavior of the intact living organism

itself, as the foundation for any systematic and satisfactory analysis of acid-base metabolism.

Beginning with this assumption, with which I would hope we could all agree, our task

would appear to be to develop a nomenclature commensurate with the physiologic complexities

that are known to exist. As a point of departure for developing such a nomenclature, there

are compelling reasons for utilizing the time-honored, carbonic acid/bicarbonate buffer

pair. As we all know, this buffer pair occupies a unique position in the living organism

since it provides the biologic linchpin connecting the chemical process of buffering, on the one hand, to the physiologic process of acidity regulation, on the other. This unique

property stems, of course, from the fact that carbonic acid and bicarbonate ion concentrations are regulated by semi -independent physiologic control systems. Moreover, extrapolating

directly from the isohydric principle, one can state that alterations in body fluid acidity

can only occur as a consequence of alterations either in carbonic acid or in bicarbonate

concentrations (or both). Disturbances of acidity can be classified with little ambiguity,

therefore, into "respiratory" and "metabolic"^ subgroups.

Considerably more ambiguity is encountered, however, in attempting to develop terms

which connote the directional changes in acidity that are produced by these two classes of

disturbances. In the first place, there is the difficulty encountered with respect to the

secondary changes in bicarbonate and carbonic acid concentration which are evoked by

primary disturbances in the countervening variable. For example, should the secondary

hyperventil ization induced by an initial reduction in bicarbonate concentration be given a separate term such as "respiratory alkalosis" when it is part and parcel of the organisms

1 Figures in brackets indicate the literature references at the end of this paper.

^Although the term "metabolic" is somewhat narrow in scope in that it fails to communicate

adequately the frequently overriding. influence of renal mechanisms on bicarbonate con- centration, it is far too well engrained to be easily supplanted.

127
inherent response to this challenge and when, in fact, the patient remains acidemic. Even more troublesome is the difficulty encountered in certain mixed acid-base disturbances in which the interaction of more than one primary disturbance, and their respective physiologic

responses, produces truly paradoxical deviations in acidity. For example, we have recently

described an experimental situation in which chronic hyperventilation causes a significant

elevation in hydrogen ion concentration [2].

Based on the considerations that I have all too briefly summarized above, I would

offer the following set of recommendations for a physiologic nomenclature adequate to deal

with in vivo disturbances of acid-base balance.

A. The acid-base status of an individual at a given moment in time should be characterized only by:

1) a direct measure of the "respiratory" component, i.e., carbonic acid con- centration (H2CO3) or carbon dioxide tension (PCO2) and

2) a direct measurement of the "metabolic" component, i.e., bicarbonate concentration

(HCO3).

B. The level of blood acidity, which is determined by the levels of the co-existing

respiratory and metabolic components, should be quantitated in terms of the actual hydrogen

ion activity, symbolized by pH or [H"*"] or a|^+.

C. Deviations from normal acidity, irrespective of cause, should be termed "acidemia"

and "alkalemia," respectively,

D. Deviations from the normal level of the respiratory component, irrespective of cause, should be termed "hypocapnia" and "hypercapnia ," respectively.

E. Deviations from the normal level of the metabolic components should be termed

"reduced bicarbonate concentration" (? hypobicarbia) and "elevated bicarbonate concentration"

(? hyperbicarbia) , respectively.

F. Disturbance of acidity initiated by changes in the respiratory component should

be termed "respiratory alkalosis" and "respiratory acidosis," respectively. Changes in the

respiratory component occurring in response to primary changes in the metabolic component

should be termed "secondary hyperventilation" and "secondary hypoventilation," respectively,

and should not be termed respiratory alkalosis and respiratory acidosis.

G. Disturbances of acidity initiated by changes in the metabolic component should be

termed "metabolic acidosis" and "metabolic alkalosis," respectively. Changes in the metabolic component occurring in response to primary changes in the respiratory component should

be termed "secondary elevations in bicarbonate concentration" and "secondary reductions in bicarbonate concentration," respectively, and should not be termed metabolic alkalosis and

metabolic acidosis.

H. For the respiratory disturbances, the adjectives "hyper-acute," "acute,"

"transient," and "chronic" should be used to designate the duration of the acid-base

disturbance with respect to the secondary physiologic adjustments in bicarbonate concentration.

1) "Hyper-acute"--the virtually instantaneous, and largely theoretical, interval

prior to the titration of tissue buffers.

2) "Acute"--the interval following the immediate titration of tissue buffers and

prior to a measurable renal contribution.

3) "Transient"--the period during which the renal contribution is accumulating.

4) "Chronic"--the indefinite period following completion of the renal contribution

to the new- steady state.

I. The adjective "uncomplicated" should be used to designate an abnormal state of

acid-base equilibrium produced by a single primary disturbance, including the influence of

128
the secondary physiologic response appropriate to the specific degree and duration of the

primary change.

J. The adjective "mixed" should be used to designate a state of acid-base equilibrium

produced by the simultaneous presence of more than one, independent disturbance.

K. The terms "correction" and "repair" should be used to designate the restoration

to normal levels of the component responsible for initiating an acid-base disturbance.

It should be noted that the terms "compensated" and "uncompensated" have been avoided

in this scheme since these terms are commonly misconstrued to mean success and failure,

respectively, in returning the level of acidity to normal; it is now abundantly clear that

unimpeded physiologic responses evoked by initiating disturbances of acid-base equilibrium

should not be expected to restore hydrogen ion concentration to control levels.

REFERENCES

[1] Schwartz, W. B. and Relman, A. S., A critique of the parameters used in the evaluation

of acid-base disorders. New Eng. J. Med. 268 , 1382 (1963).

[2] Cohen, J. J., Madias, N. E., Wolf, C. J., and Schwartz, W. B., Evidence that the renal

response to hypocapnia is not homeostatic in origin. Clin. Res. 23^, 430A (1975).

129
I

I

I

I

I
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

MINIMAL ACCEPTANCE CRITERIA FOR ACID-BASE NOMOGRAMS

Jordan J. Cohen

New England Medical Center Hospital

Boston, Massachusetts 02111, USA

Acid-base nomograms can be classified into two major subgroups: (a) those providing a graphic representation of the chemical equilibrium of the carbonic acid-bicarbonate and/or

other buffer systems, and (b) those concerned with delimiting the range of physiologic

responses to be anticipated, on empiric grounds, in response to primary disturbances of

acid-base equilibrium of a specified magnitude.

The first class of nomogram has proven useful as a convenient means for specifying the

value of one (unmeasured) variable when other components of the equilibrium reaction have

been independently analyzed. These nomograms are also helpful in the teaching of acid-base

physiology since they can often be used as a backdrop for the visual representation of the

various pathways of disturbances. Such nomograms have also, on occasion, been modified and

extended to permit evaluation of certain "derived" parameters which have been offered as aids to the analysis of clinical disturbances of acid-base equilibrium. To the extent that

such nomograms reflect accurately the actual chemical equilibrium upon which they are based,

one can certainly have no quarrel with their use as handy, calculational aids. It is only

when such nomograms make available certain derivative acid-base parameters that are subsequently misused or misinterpreted by the unsophisticated that problems arise. Clearly,

this is not an indictment of such nomograms themselves, but does raise certain risk-benefit

considerations that might provide a useful focus for our discussion. To put my own bias

squarely on the table, let me say that I can find no compelling reason for developing

nomograms or algorithms that permit the calculation of so-called derived parameters since an adequate and straightforward analysis of acid-base disturbances can be based directly on the

actual concentrations of bicarbonate and carbonic acid. To state the issue in a slightly

different way, the information content required to analyze a clinical acid-base disturbance

is no less when using derived parameters than when using the actual values for bicarbonate

and carbonic acid (PCO2) themselves.

The second class of nomograms, that is concerned with delimiting the empiric physiologic

ranges for the anticipated responses to a given acid-base disturbance, has gained wide

popularity in recent years. This type of nomogram has, of course, been made possible owing

to the availability of experimental data which serve to characterize the response of otherwise normal living organisms to graded degrees of both respiratory and metabolic acid-base

disturbances. A typical nomogram of this type was published recently by Arbus [1]^. In order for such nomograms to be used appropriately as a clinical aid, the delimited zones must be based upon steady-state observations in individuals known to have uncomplicated

acid-base disturbances. Moreover, to the extent that animal data are utilized for such

purposes, it must be clearly understood that extrapolation to human beings is tentative and

speculative. Although the applicability of this type of nomogram to the clinical evaluation

of acid-base disturbances has been widely acknowledged, care must be taken to avoid over interpretation and over simplification. For example, a common misconception is that values

that fall within the zone ("confidence band") of a given uncomplicated acid-base disturbance

necessarily indicate the presence of an uncomplicated disturbance. In fact, it can be shown

that the co-existence of dual or triple ("mixed") disturbances may cause acid-base parameters

to fall well within a given confidence band [2].

Figures in brackets indicate literature references at the end of this paper.

131
For this and other reasons, it must be emphasized that such graphic aids cannot replace

the intelligent synthesis of all relevant clinical and laboratory data in achieving a rational diagnosis and therapeutic plan.

REFERENCES

[1] Arbus, 6. S., An in vivo acid-base nomogram for clinical use. Can. Med. Assoc. J. 109 , 291 (1973).

[2] Arbus, G, S., Hebert, M. D., Levesque, P. S., Etsten, B. E., and Schwartz, W. B,,

Characterization and clinical applications of the "significance band" for acute

respiratory alkalosis. New Eng. J. Med. 280 , 117 (1969).

132
National Bureau of Standards Special Publication 450. Prcc£tci1ngs of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

A PHYSIOLOGICAL APPROACH TO ACID-BASE DIAGNOSTICS

Poul Kildeberg, and Knud Engel

Departments of Clinical Chemistry and Pediatrics

Odense University

DK 500 Odense, Denmark

The organism may be conceived as a pluricompartmental buffer solution through which

titratable acid and base, because of a relatively high buffer value (b) and the presence of

specific control mechanisms, are transported at varying rates at a nearly constant extracellular pH. Correspondingly, the clinical acid-base physiologist is concerned with acid

and base rather than with acids and &asee--being more often than not in no position to imply

specific molecular moieties: his principal task is to evaluate extent and control of

hydrogen ion donation in the body, i.e.^ to assess sources of gain and loss, fluxes,

distribution, concentrations, turnover, and exchange of titratable acid or base. The acidbase physiologist, then, is primarily concerned with the three key variables of buffer

chemistry: pH, B, and titratable acid (TA). In terms of Van Slyke's definition [1],^

pH = - /e'^dTA + pHq (1)

where pHg is the titration end-point.

Amounts of hydrogen ion transfer taking place under circumstances

defined by end-point specification are stated in terms of the acid-base

chemical quantity, TA. Magnitudes of that quantity are expressed as the product of a dimensionless number (numeric value) and a unit. The

basic unit concerned is (milli) mole of titratable hydrogen ion, conven- tionally named (milli) equivalent. The use of this unit makes it possible

to describe the total titratable acidity of any known solution as the

algebraic sum of fractional titratable values contributed by individual

specific Br^nsted acids (and bases) making up the system. Although the

milli-equivalent (meq) is being abandoned in current clinical chemical

usage [2], it remains particularly convenient for the buffer chemist.

The meq is derived from the mmol by qualification with respect to component (H"*") and expressed by multiplying the basic unit by a dimensionless quantity, N, the value of which depends on the acid properties

(strength) of the molecular moiety concerned. For a given nonionic

acid, meq = mmol titratable H"*" = N-mmol of acid. Accordingly, TA (meq)

= C'N'mmol of acid where C is the numerical factor of the millimolar

concentration. N is the average number of negative charges conferred

per molecule of the nonionic conjugate acid at the neutral pH (pHo). In general , N is equal to

pH(.-pK, 2pHn-pK,-pK2 i-pH^-pKn ... -pK10 ' + 2-10 ^ ' ^ ... + i-lO " ' ^

pHf.-pK, 2pHn-pK,-pK„ i-pHn-pKn ... -pK1 + 10 ^ ' + 10 ' ... + 10 " ' ^

^ Figures in brackets indicate literature references at the end of this paper.

133
where pK-j is the pK^ value corresponding to the i'th dissociation step of

a pluribasic acid. For an arbitrary solution of an acid and its salt(s), we get

TA(meq) = N-C^^^^^ - [b""^] - 2[b"-2].... i[B"-'"]

where C^otal ''s the millimolar buffer total concentration and b"~^ are conjugate bases of the acid HB". At very low or very high levels of pH,

the equation must be extended by the terms +[H''"] and -[OH"].

To the extent that different categories of Br0nsted acids and bases occurring in the

body exibit physiological specificity in the sense that they conform with characteristic

patterns of origin, transport, and elimination, the acid-base physiologist must evaluate

titratable contributions by such categories separately. Therefore, a very fundamental

question is this: Do we know categories of acids and bases the unspecific net titratable

contributions by which are subject to organ-specific control and elimination? Viewing the

organism as a "pH-stat" which within limits of the capacities of the organ systems involved

ensures constancy of the extracellular pH, we may rephrase the question: Do we know

categories of titratable acid and base which are subject to independent control in relation

to a common extracellular pH?

Whereas the total TA of the body fluids (as determined by closed system titration) is of limited physiological interest, the unique role of cavhonic acid in acid-base metabolism

was early recognized; and the development by Van Slyke [3,4] of methods suitable for CO2

measurement in biological fluids heralded fifty years of acid-base physiological research.

However, as an immediate consequence of the introduction of quantitative CO2 measurements

in biological chemistry, the operational concept of non-oarbonic (non-volatile) acid arose which has been a constant source of controversy. By means of various modifications of the

CO2 technique, clinical physiologists and clinical chemists have endeavored to express— in

semiquantitative machine language: "CO2 capacity" [5], "CO2 combining power" [6], "alkaline

reserve capacity" [7], "standard bicarbonate" [8], etc. --concentrations of non-carbonic

acid without ever questioning the relevance of this concept. It is characteristic of the

situation that when Astrup, et al. [9,10] finally succeeded in measuring (by CO2 titration)

exact concentrations of titratable non-carbonic acid in blood, an extensive debate evolved

which did not establish any such relevance [11-18]. The difficulties were compounded by a persistent demand by clinicians for diagnoetioally (physiologically) specific acid-base

measurements.

Measured concentrations of non-carbonic acid (base) in biological fluids include: (1)

titratable equivalents originating in reversible processes of intermediary metabolism or represented by exogenous (dietary) metabolizable "organic" acid, and (2) titratable equivalents originating in irreversible reactions or gastro-intestinal absorption of non-metabolizable acid and base. The kidney contributes to acid-base homeostasis by exclusively

influencing the extracellular concentration of acid (base) belonging to the second category

which in turn cannot be changed by extrarenal mechanisms unless the capacity of the kidney

in this respect is overcharged. Correspondingly, changes in the level of renal function

cannot per se bring about variations in extracellular concentrations of carbon dioxide or metabolical ly reversible (eliminable) equivalents which depend on pulmonary function and

metabolic processes, respectively. Physiological coupling of variations in extracellular

concentrations of titratable carbonic acid (CA = 0.9522'[C02]tQtal ) metabolizable and

non-metabolizable, non-carbonic acid takes place by the operation of renal, metabolic, and

respiratory control mechanisms in relation to a common "set-point" pH. For example, a primary rise in the extracellular concentration of non-metabolizable base will elicit a corrective response by the kidney and compensatory responses by the respiratory centers

[18a] and soft tissues [19-22].

1. Definitions

The chemical concept of non-carbonic acid is devoid of physiological specificity:

implicit in current knowledge exist three rather than two relevant categories of (titratable)

acid and base [22],

134
Titratable acid (TA) (chemical language)

I

'

1

I

1 MA NA CA (physiological language)

(soft (kidney) (lung)

tissue)

where MA, NA {"net aoid") , and CA represent titratable values of metabol izable acid, non- metabolizable, non-carbonic acid, and carbonic acid, respectively. Each of these three

categories of titratable acid may be given an explicit operational definition:

Net acid (NA) s titratable acid on titration to an end-point at pH 7.40, temperature,

37 °C, ionic strength of normal blood plasma, and CA = MA = zero meq/^.

Metabol izable acid (MA) e titratable acid on titration to an end-point at pH 7.40, tem- perature 37 °C, ionic strength of normal plasma, and NA = CA = zero meg/£.

Carbonic acid (CA) = titratable acid on titration to an end-point at pH 7.40 temperature

37 °C, ionic strength of normal plasma, and NA = MA = zero meq/£.

Obviously primary definitions are required for two of these quantities. Carbonic acid is quantitatively defined as the titratable value of the stoichiometrical CO2 content/concentration of a given system. Disregarding carbamination and "CO2 fixation" we have (pK^ = 6.11).

1.0502-CA = (C02)tQtal " + H2CO3 + HCO^ + CO3

where CA is equal to the millimolar content of bicarbonate plus twice the millimolar content

of carbonate at pH 7,40. Also, for each system or organism concerned, a decision must be

made with respect to the distinction between rsversihle and non-reversible equivalents.

2. Methodology

Concentration of CA are measured a.m. Van Slyke or by the equilibration method [10].

Concentrations of NA may be determined by direct titration following in vitro ashing [23,24].

This method, however, assumes an identity between (the extent of) in vitro and in vivo

processes of oxidation which does not always obtain, particularly not in the growing organism

where neutral sulfur is stored in new protein. A more convenient approach, which has the

further advantages of allowing separation of "addition disturbances" from "subtraction

disturbances" and of statements of titratable values at specified levels of ionic strength,

is based on the principle of electroneutrality. Thus, for any sample the concentration of NA is given by

- ^^^An-]^^ - EKjCat^]^^ (2)

where [] are stoichiometrical millimolar concentrations. Kg and Kc denote the average

number of negative and positive charges, respectively, per molecule at pH 7.40; and the ^ subscript "nm" signifies non-metabolizable ionic species (in biological fluids largely Na , K+, Ca++, Mg"^"^, CI", SOi;, and phosphate). Accordingly, the balance of net acid [22,25-31]

is given by

NAB = 1.8.Bp + B^^. - B^j^, - B^, - 2 • B^^,, - 2 • B^^,,

(3)

135
where "B" are millimolar balances representing differences between oral and parenteral

intake and renal and gastrointestinal losses. It will be noted that eq. (3) assumes identity between endogenous production (gain) and renal excretion (loss) of sulfate ion. It is apparent that the concept of NAB is deeply embodied in the older medical acid-base literature [30] and that no genuine discrepancy exists between the older "anion-cation terminology" and the acid-base chemical concepts of "acid" and "TA".

Finally, MA is obtained as the difference between NCA, determined by titration at a

P-ri of zero mm Hg, and NA.

3. Normal Acid-Base Metabolism

As a result of endogenous metabolic activity, very large quantities of CA and MA are subject to exchange and cyclic turnover, respectively, at an extracellular pool size (CA

+ MA) of about 525 meq in the normal adult. Because the pH of extracellular water is approximately 7.42, an extracellular pool of net base of the same magnitude must be

maintained. The kidney accomplishes this by continuously releasing hydroxyl ions (accompanied by non-metabolizable cation, i.e., net base) to the peritubular venous return at

a controlled rate, in the adult about 6300 meq/24 hours. Because in the physiological

range of pH free hydroxyl ions exist only in negligible concentrations, OH" ions generated

by the renal tubules are instantaneously and veversibly neutralized by endogenous CA and

MA. Extracellular net base, therefore, is represented by bicarbonate and "metabolizable

organic anion" (MOA"). Of the net base concentration of normal plasma (41 meq/£), ap- proximately 17 meq/Ji or 41 percent are represented by MOA" (partly protein). Net base

absorbed from the gastrointestinal tract and net acid (sulfuric acid, uric acid, bilirubin,

etc.) and net base (urea, creatinine, etc.) originating in irrevevsihle metabolic reactions

as well as net base communicated by bone contribute to the extracellular pool. The net

effect of such processes is normally given by the rate of renal acid excretion, equal to the difference between the rates of tubular net base generation and glomerular net base

loss [22].

Urinary net acid is partly represented by the metabolizable cation, NH^.

By generating ammonia, the kidney does net ultimately remove hydrogen

ions from the body fluids. It does, however, replace extracellular NA

by (somewhat larger quantities of) MA which can be eliminated by extrarenal

mechanisms.

12365

A H2CO3

-10V

OH"

6310

7365

MOA

5000

OH +H2CO3 + MOA

w

[h 003] + [moA"]

(24) , (17)

1300 [hco3]+[moa]

75

(27) (8)

-75 NH.

-10

Figure 1. A quantitative outline of acid-base metabolism. P = plasma

compartment, g - glomerular membrane, F = glomerular ul traf i 1 trate.

136
The acid-base homeostatic "capacity" of the kidney is reflected by the fact that

renal net base exchange (-20 - 0 meq/24 hours) amounts to only a fraction of a percent of

the rate of renal net base turnover. An outline of the gross quantitative aspects of

acid-base metabolism is shown in figure 1.

4. Acid-Base Status: Clinical Evaluation

The term "acid-base status" implies a set of relevant variables which adequately

characterize the acid-base composition of a given sample medium. The acid-base status of

biological fluids (blood, plasma, urine, cerebrospinal fluid, milk, gastric juice, etc.)

lould always be expressed in the same chemically unambiguous terms.

CA

NA

MA TA pH = pKg + log

[B]

[A]

(4)

where the right-hand equation refers to selected relevant acid-base pairs, usually the

HCO3/H2CO3 system. Representative values for some biological fluids are listed in table L

The relationship between the quantities recommended here and some commonly used plasma

acid-base variables is given by

TA - CA = NCA = NA + MA = NBB - BB - K = -BE - K (5)

where NCA is titratable non-carbonic acid, NBB is "normal buffer base" [10], BB is "buffer

base" [32], BE is "base excess" [10], and K is a constant, approximately 23.8 meq/£.

Note: If 25 mmol of CO2 (pKi 6.11), 17.2 mmol of lactic acid (or 103 g of albumin, TAgib " 0.167 meq/g) and 41 mmol of sodium hydroxide are added to one liter of water in a closed system, the resulting solution

will have a pH of 7.40, a P-^ of 40 mm Hg, and a BE of zero meq/£.

LU2

Table 1. Representative values for acid-base status

of some biological fluids.

CA MA NA (meq/L)

TA pH

Blood plasma^ 24.0 17.0 -41 .0 0 7 4

Whole blood^ 21.7 1.0 -17.4 5. 3 7 2^

Normal acid urine 1.2 0^ 5 6. 2 5 8

Human milk 4.9 25.0 -24.0 5 9 6 6

Cow's milk 4.1 66.0 -60.0 10. 1 6 5

arterial ized whole blood, hemoglobin concentration 15 g/100 ml filtered metabol izable acid minus secreted metabol izable base (NH3);

50 - 50 = 0 meq/l

hemolyzed blood

137
The pH-stat concept implies three mutually non-exclusive causes of deviations in

extracellular pH, applicable to each regulatory organ system involved: set-point deviation, functional errors, and overcharging by excessive loads. Because of the com- plexity of the biologic pH-stat careful operational definitions of these are mandatory.

-eve-vekarg-inct- Vie- may undeics-tand "loading with acid.,.or-Jias€-r €xogenous -wiit. respect to the

regulatory organ considered and sufficient to produce a stable change in extraclTllilar pH.

Malfunction may be taken to imply intrinsic disease of the organ concerned; and--by

exc^us^ on- -deviations in set-point may be indicated by functional changes in the absence

of (local) disease and acid-base loading. In such terms, diabetic keto-acidosis represents

malfunction of soft tissues and overcharging of kidney and lung; obstructive or restrictive

lung disease with chronic hypercapnia and a normal extracellular pH represents malfunction

of the brain-lung system; and steriod alkalosis and acute salicylism involve set-point

deviations for kidney and lung, respectively.

Any such combination of causes may lead to a primary change in the TA of extracellular water, the associated change in pH setting off secondary processes of correction

and compensation. The resulting ATA is the algebraic sum of concurrent changes in the

concentrations of CA, MA, and NA. Each such change is brought about by a combination of

processes of retention and redistribution and changes in the volume of solvent, due in

turn to osmotic redistribution of water across the cell membrane or to changes in water

balance. Balance changes may be accounted for by identifying several sources of input and

output. This approach to systematic evaluation of clinical acid-base disturbances is

shown schematically in figure 2. As an example, figure 3 shows the suggested steps in the

input

out

put bal

set pointmolfunct load

1

sec corr prim sec A\A comp

1 1

1 pH

result ATA

CAUSES

(etiology)

OBSERV.

(din. chem.

distr

MECHANISM

(physiology)

distr bal

input urine

r T

bal

acid

gain

base

loss RTA

rate

Mm'

'grad

lim'

prim

Z\NA

sec. changes result

AHA

solv

vol

distr

ure- mica

NH3

def?

Figure 2. Interpretative steps in the

evaluation of the acid-base status. Figure 3. Evaluation of a primaryrise in the concentration of

extracellular net acid.

diagnostic interpretation of an observed rise in extracellular NA, which on the basis of

the pattern of the acid-base status of blood, the total clinical situation of the patient,

and existing physiological knowledge is considered to reflect primary accumulation of NA

in the extracellular compartment. It may be pointed out that a normal acid-base status of

138
blood does not preclude a significant disturbance of acid-base metabolism. For example,

in hyperparathyroidism primary redistribution of body net base (release of skeletal base)

may lead to negative net base balance in the absence of significant change in the acidbase status. Because a primary rise in extracellular NA rarely if ever is brought about

by processes" of" redistribution, such changes may be evaluated in terms_of the -balance of

net acid and changes in the volume of solvent ("contraction alkalosis" and "dilution

acidosis", respectively) [33,34]. Contraction alkalosis due to negative water balance and

a secondary readjustment of the renal tubular set-point contributes significantly to the

acid-base status in "gastric alkalosis" [25]; and extracellular dilution acidosis due to osmotic redistribution of water across the cell membrane can be produced experimentally by

infusion of hypertonic solutions of aprotes such as sodium chloride and glucose [35]. A positive balance of net acid is a frequent cause of primary increases in extracellular NA. Defining the net acid output as urinary net acid [22,31], and output disturbance leading

to NA accumulation is by definition a renal acidosis of which two types are generally

recognized: the acidosis of uremia and renal tubular acidosis (RTA). Renal tubular

acidosis in turn may be due to failure to establish an adequate trans-tubular pH gradient

("distal" or "gradient limited" RTA [36]), to a relatively low rate of (proximal) tubular

net base generation ("proximal" or "rate limited" RTA [37]), or possibly to impaired

ammonia excretion [38]. Finally, input disturbances involve extrarenal sources of net

acid, in particular oral intake of NA ("addition acidosis") and gastro-intestinal losses

of net base ("subtraction acidosis").

References

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value to the dissociation constant of the buffer and the concentration and reaction of

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[2] Dybkaer, R. and J0rgensen, K. , Quantities and Units in Clinioal Chemistry^ including

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139
[15] Schwartz, W. B. and Relman, A. S., A critique of the parameters used in the

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Berry, M. N. and Scheuer, J., Splanchnic lactic acid metabolism in hyperventilation,

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Eldridge, F. and Salzer, J., Effect of respiratory alkalosis on blood lactate and

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Haldi, J., Lactic acid in blood and tissues following intravenous injection of sodium

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Camien, M. N; and Gonick, H. C, Relationship of renal "net acid" excretion to titratable

ash-acidity (ashrTA) in diet and feces, Proc. Soo. Exper. Biol. Med. (N.Y.), 126 , 45 (1967).

Gonick, H. S., Goldberg, G., and Mulcare, D., Reexamination of the acid-ash content of

several diets, Amer. J. Clin. Nutr. 2]_, 898 (1968).

Degn, J. K. , Wamberg, S., Engel K., and Kildeberg, P., Metabolic alkalosis in obstructive

vomiting. Volume depletion and balance of net acid. Acta Paediat. Soand. 63^, 537 (1974).

Goodman, A. D., Lemann, J., Jr., Lennon, E. J., and Relman, A. S., Production, excretion

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Lemann, J., Jr., Lennon, E. J., Goodman, A. D., Litzow, J. R., and Relman, A. S., The

net balance of acid in subjects given large loads of acid or alkali, J. Clin. Invest.

44, 507 (1965).

Lennon, E. J., Lemann, J., Jr., and Litzow, J, R., The effects of diet and stool

composition on the net external acid balance of normal subjects, J. Clin. Invest. 45,

1601 (1966).

Shohl , A. T. and Sato,. A., Acid-base metabolism. I. Determination of base balance,

J. Biol. Chem. 58, 235 (1924).

Wamberg, S. T., Engel, K., and Kildeberg, P., Balance of net base in the rat. Reference values in relation to growth rate, Biol. Neonat. 28_, 171 [1976).

Singer, R. B. and Hastings, A. B., An improved clinical method for the estimation of

disturbances of the acid-base balance of human blood. Medicine, 27, 223 (1948).

Cannon, P. J., Heinemann, H. 0., Albert, M. S., Laragh, J. H., and Winters, R. W.,

"Contraction" alkalosis after diuresis of edematous patients with ethacrynic acid,

Ann. Int. Med. 62, 979 (1965).

140
< [34] Winters, R. W. , Scaglione, P. R., Nahas, G. G., and Verosky, M., The mechanism of

acidosis produced by hyperosmotic infusions, J. Clin. Invest. 43, 647 (1964).

[35] Sotos, J. F., Dodge, P. R., and Talbot, N. B., Studies in experimental hypertonicity.

II. Hypertonicity of body fluids as a cause of acidosis, Pediatvios, _30, 180 (1962).

[36] Morris, R. C, Jr., Renal tubular acidosis. Mechanisms, classification and implications, mew Eng. J. Med. 281, ^405 (1969).

[37] Rodriguez-Soriano, J. and Edelmann, C. M., Jr., Renal tubular acidosis, Ann. Rev. Med.

29, 363 (1969).

[38] McCance, R. A., Matheson, W. J., Gresham, G. A., and Elkinton, J. R., The cerebroocular-renal dystrophies: a new variant, Arch. Dis. Childh. 2S_, 240 (1960).

141

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

CO2 SOLUBILITY, pK' AND RELATED FACTORS IN ACID-BASE BALANCE

William H, Austin

Maine Medical Center

Portland, Maine, U.S.A.

1. The Role of pK', CO2 Solubility, and Temperature

in Acid-Base Measurement

An appreciation of the problems of pK' and CO2 solubility necessitates an understanding

of some of the features of the Henderson-Hasselbalch equation that are not always emphasized.

With this in mind, a brief restatement of this equilibrium reaction, which provides a system for the examination of acid-base balance, is made here. The thermodynamic equation

may be derived from the equilibrium reaction, starting at the partial pressure of CO2 to the dissociated hydrogen ion activity.

PCO2

^

dissolved CO2

^

H2CO3 H"^ + HCO3 (1)

The first consideration, usually not made a part of the familiar constant, is the solubility

of CO2 itself.

The first step involves the uptake of CO2 in the plasma. The constant which is necessary for the understanding of this phenomenon is the solubility coefficient a., (Bunsen's

coefficient) [1]^, or "S", the solubility factor [2]. Neither appears in the collected

constant (pK'), though it is mathematically possible [3,4]. Some consideration has been

given to this, but classically it is excluded from the "overall" constant. The latter,

often called the first dissociation constant, or apparent pK' , is in reality a collection of

the ionization and hydration constants and will be treated later. CO2 solubility may be

given as Bunsen's solubility coefficient a, (ml C02/ml plasma), often anglicized as "a", and

frequently mistaken for the CO2 factor "S", which is millimoles of COa/mm HgPC02.

Some use a, "a", and "S" interchangeably, as noted recently [5]. This point should

be clarified. The original solubility of 0.541 ml/ml [6] was determined by Bohr in 1905,

and has an interesting history. Van Slyke [7], in 1928, restudied the solubility, using

apparently four samples: two of ox serum, one of normal human serum, and one of normal

human oxylated plasma. He obtained values of 0.509, 0.511, 0.506, and 0.511 ml C02/cm3

for these respective samples. The more familiar figure of 0.510, or "S" = 0.0301, is therefore in reality an average of solubilities on ox and human plasma and serum with a water

content of 0.925 g/cm^ at 38 °C. Eisenman, et at. [8], Danowski and Gilmore [9], and we

[10] found different water contents, and used 0.935 in our solubility studies. The recalculation of Van Slyke's data, using a higher water content, interestingly enough yields values

almost identical to ours, as referred to in the "Handbook of Physiology" [11], and the Ad

Hog Committee on Methodology of the New York Academy of Science in 1966 [12]. This value is "S" = 0.0306.^

^Figures in brackets indicate the literature references at the end of this paper.

2 Van Slyke astutely recognized [7] that the actual value of a or S is irrelevant if it is used with the appropriate pK' . Note that "ml" or "cc" is used exactly as cited in the original papers.

143
For years, the variation of CO2 solubility in plasma was assumed to CO2 parallel that

of water. This was found not to be the case (fig. 1) [10,13], and Severinghaus and others

made appropriate corrections which were reflected in pK's, especially at the lower temperatures [10,13]. The variations at 37 °C are ^mall by clinical standards, and should not pose- any problems-.

Water--Bartels et ai. "' Extrapolated serum--Dill & Forbes

Extrapolated plasma--Severinghaus

et al . Human serum--author

Ox plasma--Bartels et ai.

15

"20" 25 30

temperature °C

35 40

1.00

. .95

.90

.85 c0

. .80 +-> 3

ml

sol

.70 s.<u0.

.65

CN 00

• .60

.55

. .50

Figure 1. Solubility of carbon dioxide in water, serum and

plasma from 15 °C to 38 °C. Solid line indicates CO2 solubility in plasma does not bear a constant relationship to water (top line), hence to extrapolated values for plasma

CO2 (Severinghaus). (Reproduced, with permission, from the

J. Appl. Physiol. 19, 893-96, 1964).

Attendant alterations in disease states, blood constituents, viscosity, and ionic

strength could conceivably alter CO2 solubility. These have been considered [14-19], and

if the solubility factor is assumed, for the sake of discussion, to be fixed, then these

changes, if any, should be reflected in the pK' . This is the constant for the balance of

the equilibrium reaction:

Dissolved CO2 HpCOc Hco; (2)

To interrelate the solubility and these reflected variations in pK' , the origin of this

constant must be examined. From this "overall" reaction, a pK' may be derived.

The "overall" equilibrium constant falls into two sets of reactions:

1. The Hydration reaction:

dissolved CO2 =^ H2CO3 which may be quantitatively expressed as (C02)K. = (H2CO3) with as the hydration constant.

2. The Ionization reaction:

H2CO3

^ H + HCO3 - quantitatively expressed as (H2C03)K. = (H )(HC03)

with K. as the ionization constant.

144
These terms may be expressed as concentration or activity. In actual practice, _H (or pH)

is an activity, since this is what the electrode measures, and CO2 content (HCO3) is concen- tration, usually from gasometric measurements. Titration values for HCO3 are less commonly

used. For this presentation, the practically measured pH (or H+) will be an activity, and

bicarbonate-related entities will be concentrations. Since both dissolved CO2 and (H+) + (HCO3) are in equilibrium with (H2CO3), the hydration and ionization reactions may be equated,

and constants collected.

(K^ X K. = K'), resulting in (H"^) = K' x (C02)/(HC03) (3)

Although this derived equation is theoretically correct, it does not correspond to the

experimental operations involved in the determination of the "overall" equilibrium constant

(K' or pK'). One experimentally fixes the CO2 concentration by equilibrating blood, plasma

or serum with a given partial pressure of carbon dioxide gas. This operation fixes the

concentration of dissolved carbon dioxide plus the true carbonic acid in proportions determined by the hydration coefficient K^. This being the case, the numerator "CO2" of the

equation must be CO2 (dissolved) plus H2CO3 (actual), as this is what the analytic determination produces, rather than what the thermodynamic determination alone gives. There are

a few bicarbonate ions^ which arise, but they arc insignificant. The "overall" reaction

must be rewritten as:

(H^) = K' . (CO2) + (H2CO3) /(HCO3) (4)

Since it has been recognized that this relationship is more properly expressed logarithmically, one may rewrite the equation as:

(HCO^)

pH = pK' + log (5)

(CO2) + (H2CO3)

The term (CO2) + (H2CO3) is often referred to as H2CO3, creating the false impression that

all of the denominator of the Henderson-Hasselbalch equation is carbonic acid as such. The

dissolved CO2 may be potential H2CO3, but not actual carbonic acid. This has been a point

of confusion to many students of carbonic acid chemistry.

The use of pK' , rather than K' , has been subjected to the same criticism as the use of

pH and H"^ concentration; however, the logic is the same. The most cogent, and sometimes

less appreciated, argument has been cited by Davis [20]. Essentially, he notes that a more

meaningful description of the equilibria constants is given in their logarithmic form. This

is not an accident, nor is it pure convenience nor perpetuation of a misconception. The

thermodynamic behavior of a substance is directly related to a potential, the logarithmic

function of activity--in this case the log (negative) of the constant K', or pK'

Davis has expressed it well, "...the most meaningful property describing dissociation

equilibria of acids, whether in chemical or physiological systems, is not the dissociation

constant itself, but the logarithm of this dissociation constant, pK' , which is directly

proportional to free energy of dissociation and clearly a measure of the strength of an acid in aqueous solution." Buffering capacity, employing the constant being discussed, is

related to the free energy of dissociation, and is not coincidental to the units used [20,21].

^Virtually all of the bicarbonate which appears in actual practice arises from the salt of

the weak acid, in this case H2CO3. The origin of the bicarbonate does not affect the logic

of the derivation; however, the large quantity of these ions from sodium bicarbonate is very

important to the buffering capability of the bicarbonate buffer system.

145
Reasoning in terms of the reactants may be a more appropriate approach. Chemical and

physiological activities are more precisely related to the potential, which in turn is a direct function of pH. Therefore, participation of hydrogen ions in physical and physiological reactions is determined by ionic potentials, which are logarithmic expressions of ionic

concentrations. (This may be extended to include the other reactants. The Nerst equation,

which deals with potentials, applies for similar entities, chemically determined, which

react physiologically to the logarithm of molarity.)

Digressing again to the example of pH, which may be more easily visualized, but applicable to pK, one finds hydrogen ion concentration can never really be derived accurately

from activity. In theory, there may be pH changes with no H""" change, if there is an independent change in activity coefficients [20]. pH, hence H+, is an apparent entity. pH results

from emf with an electrode calibrated with buffers of "assigned" values. Though assigned,

these values are probably very accurate; however, the pH one gets when blood is measured is relative. The blood or plasma is a biological fluid, differing greatly from the buffer.

The pH "number" is obviously only relative. There must be distortion 'also of the "real" pH

through inherent errors in the measurement itself, apart from the buffer. Studies of liquidjunction potentials show this [22], so how, especially in view of the other points, can one arrive at a meaningful H"*" concentration?

The argument of putting acid-base measurement in familiar terms is a pitfall. One

takes this entity, calls it a concentration or equivalent, and immediately assumes that the

physiologist or clinician can now understand it and can relate this value to other parameters

The end result is like putting a planet and a grain of sand side by side and saying, "Now we

are dealing with mass, and all is clear." The student, through working with concentration,

has to be constantly aware of the diversity of the magnitude of concentration and, not as has been the case, think of H"*" in the same scope as Na+ or HCO3 concentrations. It is easy

to see how the dimensions of H+ can be thrown out of proportion after thinking of Na"*", HCO3, CI", K+, etc. It is too easy to lose track of the diminutive nature of hydrogen ion

concentration, especially since it is the intensity of acidity that we are seeking. A change in using the inexact concentration concept can only foster this, whereas pH has the

dimensionless nature which would diminish this tendency without the need for clarifying that

H''' is not a true concentration, and is calculated or arrived at with the many reservations

mentioned above, keeping in mind the magnitude of the quantity.

Also, conversion to concentration does not have the advantage of keeping all of the

parameters in the same dimension. In addition to the gigantic difference in magnitude, we

are now accustomed to using PCO2, a far more easily acceptable entity than H2CO3, which is

not without its conceptual problems, (i.e.^ H2CO3 is really dissolved CO2 and true HjCOg—

a point of much misunderstanding). It makes no sense to say that H"^ concentration is similar

to bicarbonate concentration (HCO3), and then relate these two to a pressure (PCO2), especially if this is done in the name of clarity. Filley [23] has stated, with references, an excellent analysis of this controversy.

Since it is more useful to use the partial pressure of carbon dioxide (PCO2) in the

equation, the solubility factor must be introduced to make (CO2) + (H2CO3) = PCO2. This

gives to the familiar:

(HCO^)

pH = pK' + log (6)

PC02- S

or, for pK' determinations made in the studies discussed:

CO2 content

pK' = pH - log ( 1) (7)

S • PCO2

'CO2 content determined gasometrically in most studies cited above.

146
The original pk"s determined by Severinghaus in 1956 [16] have been widely used. A value

for pK' has been determined in many studies, and has been said to be a function of pH [16,24]

and temperature [10], and are affected by disease states [17] and biochemical abnormalities

[151. Re-examination of pK' strongly indicates to us that neither pathologic conditions

[26] nor abnormal blood constituents [15], exclusive of temperature, alter the pK' in a practical way. In fact, we could not find significant variations with pH within the clinical range [27]. This is actually similar to the classical studies of Severinghaus [16], and

is confirmed by Nunn [28]. By significant, we would mean less than 0.01 between 7.10 and

7.70 Cfig. 2). This is not to say that electrodes and other factors described by Maas [4],

pK' = 6.132 -0.0042 pH± O.OII

•

• •

- • •

•

• • • • rr^.—•

pK' • • pH

1 1 1

1 1 1 7.10 7.20 7.30 7.40 7.50 7.60 7.70

Figure 2. pK variations with pH in human blood. (Reproduced, with permission, from the Archives of Internal Medicine , 126 , 699 1970).

Severinghaus [16], and Siggaard-Andersen [24] do not give rise to pH/pK variations, but

that for all practical purposes, a pK' of 6.10 on whole blood pH (not corrected for "suspension", liquid-junction effect) plasma or serum CO2 content, gasometrical ly determined, and

fixed, tonometered CO2 pressures, is suitable for clinical, and even research, purposes at

37 °C [27]. It is of interest that pK' determinations made using the CO2 electrode revealed

the same pK' , but with a much wider standard deviation. The pK's are termed "operational",

since no attempt is made to estimate or calculate pHs of plasma because only precise values

for whole blood pH were known, and "correction" could not be satisfactorily made despite

estimated differences of 0.01 by some [2]. CO2 content may be determined accurately by the

Van Slyke method, though all of the CO2 is not from HCO3 or dissolved CO2. Carbonate and

carbamate certainly figure into the picture, but are of questionable significance in the

overall picture.

The PCO2 electrode, although not used as the primary method of classical assessment of

partial pressure of CO2 for pK' work, has shown limitations which, though small, might be

considered by some to be important.

The magnitude of the variation in PCO2 measured by electrode may be seen in an indirect

way by comparing pK and S.D., using actual PCO2 measurements and electrode-measured PCO2.

This reflection in the S.D. of the pK is from ±0.013 to ±0.032 [26]. In terms of actual

partial pressure, the standard estimate of the error for actual and electrode PCO2, measured

from approximately 20 mmHg to 50 mmHg, is ±3.5 mmHg. Other studies over a larger range,

with 142 samples, give slightly smaller but significant differences of ±2.11 mmHg [29].

147
Figure 3 breaks down values for pK' as determined in patients and in normal samples by

tonometering high and low CO2 concentrations. The significance is in the similarity. Rigid

techniques were observed, as in previous studies [26], and significant differences were not

present. This is very consistent with another fifty determinations, using randomly selected

patient blood (fig. 4). Whole blood pH and electrode PCO2 measurements were used with Van

5% CO2

NORM, and ABNORM

6.098± .010

PK' TOTAL (56)

6.101 ± Oil /A

10% CO2

NORM.and ABNORM

6.104+ .011

ABNORMAL (28)

6.102 ± .013 NORMAL (28) 6.101 ± .009

(14) I0%C02

6.104+ .013 / \

(14) 5% CO2

6.097 ± .009

(14) 5% CO(14) I0%C02

6.101 ± .013 6.104 ± .008

Figure 3. pK's derived from blood of normal and ill patients by tonometry

with gas mixtures of approximately 5 percent and 10 percent. No apparent

differences noted under various circumstances.

6.03 6.04 6.05 6.06 6.07 6.08 -6.09 6.10 6.11 6.12 6.13 6.14 6.15 PK'

Figure 4. pK' determinations derived from tonometered whole blood with

fixed partial pressures of CO2 and with PCO2 determined by PCO2 determined by PCO2 electrode. Hatched boxes represent two samples, one for

each category. Note the narrow S.D. with, the tonometered method.

148
Slyke CO2 contents as above. The mean was 6.099. Again, the only difference is in the

S.D, of ±0.030, presumably due to CO2 electrode variations. When twenty-eight of these

samples were examined with the tonometry method, the mean was 6.012 and the S.D. was ±0.013.

Perhaps the most cogent point arising from the pK' studies relates to the calculated

CO2 contents' relationship to actual, directly determined Van Slyke C02's. This variable (or

HCO3) is most frequently used at present, and it is significant in that 54 of 59 samples

(most referred to in studies presented here) are within 3 millimoles/liter of the directly

determined figure, with the maximum spread of 5 millimoles/liter occurring only once. This

is seen with the simple uncorrected pK' of 6.100. Using pK's corrected to pH [10,11] changes

the calculated values by less than 1.0 millimole in all instances.

A relatively unexplored, but not totally unrealistic, concept would involve the inclusion of the solubility factor into the pK' itself. Maas [4] has explored this, and Van Slyke [7] himself alluded to this interrelationship many years ago. Simple, logarithmic

conversion of "S" (0.0306 to 1.50) added to 6.10 would yield 7.60 as the pK' , the only

constant in the entire equation. In effect, this is what we actually do {i.e., use an appropriate pK' for a given "S"); however, conceptual and logistical factors make this claim

of simplification debatable.

Measurement, by agreement, may be made at 37 °C. Buffers and calibrating gases will

not be treated here, but the current practice of using pH/C factors of -0.0148 (0.0150)

[30-32] and temperature/PC02 factors [33,34] can persist and yield universally acceptable

values. Since Nunn [28] and Severinghaus [33] have good evidence of the stability of CO2

content at different temperatures, this allows determinations of pH and PCO2 at 37 °C and

calculations of CO2 content (or bicarbonate) which are still valid where pH and PCO2 are corrected. The other route of correcting pH and PCO2 to body temperature then, with appropriate solubility factors and pK's calculated to the third variable [11,32] (HCO3 or CO2

content), is also valid but not necessary, and it is cumbersome. Tables [36], slide rules

[35], and diagrams [37,38] will simplify the calculations, and whichever is preferred may be

acceptable, as long as it yields the same answers for the calculated third variable. This

matter, again, is under scrutiny; here standardization should be achievable.

2. Summary

We must use the most refined techniques: stable, sensitive electrodes and electrometers,

accurate buffers, good CO2 electrodes and calibration, and careful sampling and storing--all

of which are being discussed, along with the best possible pK's and CO2 solubilities; however,

practicality must prevail. It is our contention that a pK' of 6.100 and an "S" of 0.0306

at 37 °C would not be unreasonable [10-12], and that whole blood pHs be used. Slight variations

of pK' with pH would not be unacceptable, but are not essential. The changes in pH with

temperatures of -0.015/°C (0.0148) [30,32], as well as PCO2 (a log PCO2 = F-T at 37 °C; F = 0.019) have been well studied and seem to be reasonable standards [28,34].

The misconception which may arise from pK' studies is that the value determined is the

appropriate one for the given sample. Because it is the result of two or three determinations, this makes it subject to the accumulation of all the errors inherent in each measurement. This is apparent from electrode and tonometer groups. For example, on one occasion

an electrode PCO2 sample gave a value of 6.030, but 6.102 when determined by tonometry.

Another sample gave 6.147 and 6.097, respectively. This suggests to me that there is no "normal" for a given laboratory. pK' may be affected by various factors as noted, but the mean arrived at from the many studies is the operational one for practical use. Considering

all the variables, true deviation of the calculated "third factor" is very hard to systematically quantitate, but must be extremely small even using these standard factors outlined

here.

References

[1] Bohr, C. and Bock, J., Bertimmung der absorption Einiger Gase un wasser Beiden temperature zwischen 0 und 100°, Weid. Ann. Physik. U. Chem. 44, 318 (1891).

149
[2] Severinghaus , J. W. , Stupfel , M, , and Bradley, A. F., Accuracy of blood pH and PCO2

determinations, J. Appl. Physiol. 9_, 189 (1956).

[3] Austin, W. H. and Stinebaugh, B. J., The "not so apparent" pK' , J. Me. Med. Assoc. 60,

10 (1969). ~"

[4] Maas, A. H. J., Eet Koolzvurevenwioht in Natriumbiaarbonaatoplossingen, Liquor Cerebrospinaliis en Menselijk Bloedplasma (Bronder-Offset, Rotterdam, 1967).

[5] Austin, W. H., Notations in acid-base balance. Am. Heart Jour. 70, 574 (1965).

[6] Bohr, C, Absorptionscoefficienten des blutes und des blut plasmas for gase, Skand.

Arch. Physiol. ]!_, 104 (1905).

[7] Van Slyke, D., Sendroy, J., Hastings, A., and Neill, J., Studies of gas and electrolyte

equilibria in blood, J. Biol. Chem. 78, 765 (1928).

[8] Eisenman, A. J., MacKenzie, L. B., and Peters, J. B., Protein and water of serum and

cells of human blood, with a note on the measurement of red blood cell volume, J. Biol. Chem. 116^, 33 (1936).

[9] Danowski, T. S. and Gilmore, G. H. , Relationships between serum water or weight and

protein concentration, j. Lab. Clin. Med. 35, 67 (1950).

[10] Austin, W. H. , Lacombe, E., Rand, P. W. , and Chatterjee, M., Solubility of CO2 in serum

from 15-38 °C., J. Appl. Physiol. 18, 301 (1963).

[11] Fenn, W. and Rahn, H., eds.. The Handbook of Physiology: Respiration, 699 pp. (Williams

and Wilkins Co., Baltimore, Md., 1965).

[12] Nahas, G. G. , Ed., Report of the Ad Hoc Committee on Methodology, Current Concepts of

Acid-Base Measurment, Annals of the New York Academy of Science, New York City, 274 pp. (Edward Bros., Inc., Ann Arbor, MI., 1966).

[13] Bartels, H. and Wrbitzky, R. , Determination of carbon dioxide absorption coefficients

between 15 and 18 degrees C in water and plasma, Pflueger Aech Ges Physiol. 2T\_, 162

(1960).

[14] Ludbrook, J., Estimation of PCO2 by means of the Henderson-Hasselbalch equation, in pH

and Blood Gas Measurements, R. F. Woolmer, ed. , p 34 (Little, Brown and Co., Boston,

MA, 1959).

[15] Austin, W. H., Ferrante, V., and Ritchie, R. , Effect of abnormal plasma constituents on the pK' of whole blood. Am. Jour. Clin. Path. 51_, 799 (1969).

[16] Severinghaus, J. W., Stupfel , M. , and Bradley, A. F., Variations of serum carbonic acid

pK' with pH and temperature, J. Appl. Physiol. 9^, 197 (1956).

[17] Trenchard, D. , Noble, M. , and Guz, A., Serum carbonic acid pK' abnormalities in patients

with acid-base disturbances, J. Clin. Sci. 32, 189 (1967).

[18] Sinclair, M. J., Hart, R. A., Pope, H. M. , and Campbell, E. J. M., The use of the Hen- derson-Hasselbalch equation in routine medical practice, Clin. Chim. Acta. 1_9, 63 (1968).

[19] Hastings, A. B. and Sendroy, J., Jr., The effect of variations in ionic strength on the

apparent first and second dissociation constants of carbonic acid, J. Biol. Chem. 65,

445 (1925).

[20] Davis, R. P., Logland: A Gibbsian view of acid-base balance. Am. J. Med. 42, 159

(1967).

[21] Huckabee, W. E., Henderson vs. Hasselbalch, Clin. Res. 9, 116 (1957).

150
[22] Austin, W. H. and Littlefield, S., The difference in apparent pH of blood and buffer

caused by raising the liquid junction from room temperature to 37.5 °C, J. Lab. and

Clin. Med. 67, 516 (1966).

[23] Filley, G. F., Acid-Base and Blood Gas Regulation, 213 pp. (Lea and Febiger, Philadelphia, PA, 1971).

[24] Siggaard-Anderson, 0., The first dissociation exponent of carbonic acid as a function

of pH, Soand. J. Clin. & Lab. Invest. U, 487 (1962).

[25] Thornton, J. A. and Nunn, J. F. , Accuracy of determination of PCO2 by the indirect

method, Guy's Hospital Report, 109, 203 (1960).

[26] Austin, W. H. , Ferrante, V., and Anderson, C, Evaluation of whole blood pK' in the

acutely ill patient, J. Lab. Clin. Med. 12^, 129 (1968).

[27] Austin, W. H. and Ferrante, V., The operational value of whole blood pK' to pH, Avoh.

Int. Med. ]2^, 699 (1970).

[28] Nunn, J. F., Bergman, N. A., Bunatyan, A., and Coleman, A. J., Temperature coefficients

for PCO2 and PO2 of blood in vitro, J. Appl. Physiol. 20, 23 (1965).

[29] De Raedt, M. , Vandenbergh, E., and van de Woestijne, K. P., Direct and indirect determination of partial pressure of CO2 in the arterial blood of patients with respiratory

insufficiency, Clin. Soi. 35, 347 (1968).

[30] Rosenthal, T. B., The effect of temperature on the pH of blood and plasma in vitro, J. Biol. Chem. 173, 25 (1948).

[31] Greenburg, A. 6. and Moulder, P. V., Temperature coefficients for PCO2 and pH in whole

blood, Aroh. Surg. 91, 867 (1965).

[32] Austin, W. H., Lacombe, E. H., and Rand, P. W., pH-temperature conversion factors and

PCO2 factors for hypothermia, J. Appl. Physiol. 19, 893 (1964).

[33] Kelman, G. R. and Nunn, 0. F., Nomograms for correction of blood PO2, PCO2 and base

excess for time and temperature, J. Appl. Physiol. 2]_, 1484 (1966).

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PO2 of blood in vitro, J. Appl. Physiol. 9, 201 (1956).

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Maldistribution, 170 pp. (Lea and Febiger, Philadelphia, PA, 1973).

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151

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DEFINITION OF OXYGEN SATURATION AND

CHARACTERIZATION OF OXYGEN-HEMOGLOBIN AFFINITY

Arthur L. Malenfant

Instrumentation Laboratory Inc. 113 Hartwell Avenue

Lexington, Massachusetts 02173, USA

1. Introduction

The function of hemoglobin is to transport oxygen from the lungs to the tissues. In order to carry out this function optimally, the hemoglobin must take up a maximal amount of

oxygen during circulation through the lungs, and then yield this oxygen to the tissues perfused

during circulation through the body. The ability of the blood to carry out both these

functions is determined by the affinity of the hemoglobin for oxygen. Specific definition

of oxygen saturation is needed.

The characterization of the relationship between the P^, and percent oxyhemoglobin

U2

saturation--the oxyhemoglobin dissociation curve--has been found to be increasingly complex.

It is now known that there are a number of factors which affect this relationship, resulting

in shifts in the oxyhemoglobin dissociation curve. The significance of these shifts centers

around the fact that the amount of oxygen available to tissues at a given Pq^ is changed.

Factors known to affect the oxyhemoglobin dissociation curve are pH, temperature, P^q

and 2,3-DPG (fig. 1). Increase in pH or decrease in the value of any of the other factors

shifts the dissociation curve to the left [1,2]^. When all these have a "normal value" the

oxyhemoglobin dissociation curve has a P50 (Pq for 50 percent saturation of the hemoglobin

with oxygen ) of 26.7 mm. In addition to the normally observed factors which influence the

oxygen-hemoglobin affinity, carbon monoxide, present because of environmental pollution and

in cigarette smoke, also effectively shifts the dissociation curve to the left [1,7].

Both acute shifts in P50 [3] and what can be considered to be adaptive shifts in P50

[4] have been observed.

2. Measurement Techniques

The techniques employed in characterizing oxygen saturation--and more importantly, the

hemoglobin-oxygen affinity--of blood must necessarily include measurement of a number of

parameters. Direct photometric measurement of oxygen saturation of blood is straightforward

Figures in brackets indicate the literature references at the end of this paper.

153
Figure 1. The normal dissociation curve for whole blood

at pH 7.4 —, dissociation curves for hemoglobin

solution at pH 7.22 under varying conditions of P^q

and 2,3-DPG/Hb ratios.

as long as factors such as carboxyhemoglobin and methemoglobin are taken into consideration

[5,6]. The typical concentrations of methemoglobin, however, are low enough to not significantly affect accuracy. Failing to characterize the presence of these species can lead to erroneously high values for saturation, however.

Other methods include direct manometric measurement of oxygen content of a particular

sample compared with the oxygen content of a sample saturated with oxygen by appropriate

means [8]. Oxygen content can also be determined using a fuel cell to consume the oxygen

liberated from blood samples, otherwise treated in essentially the same manner as the

manometric technique above. Comparing oxygen content with oxygen capacity yields oxygen

saturation defined as a percentage of "available hemoglobin"--hemoglobin available for

combination with oxygen [9]. Comparing oxygen content with that oxygen capacity theoretically available through measurement of total hemoglobin concentration yields an oxygen

saturation defined as a percentage .of total theoretical oxygen capacity [6]. The molecular

154
weight of hemoglobin has been accepted as 64,458 [10]. Each molecule of hemoglobin combines

with four molecules of oxygen. Therefore, in traditional terminology, 1.39 ml of oxygen at STP will combine with 1 gram of hemoglobin. Hemoglobin concentration expressed in grams

(g/100 ml) multiplied by 1.39 ml of oxygen yields the theoretical oxygen capacity of the

blood sample. The relationship in more acceptable terms is 4 moles of oxygen per mole of

hemoglobin. Hemoglobin concentration of 2.327 millimoles (15g%) theoretically binds 9.308

millimoles of oxygen (20.85 ml) when complete saturation is achieved.

3. P50 Nomogram

Use of measurements typically employed in blood gas analysis, i.e., pH, Pq^ and P^q^

combined with the measurement of 2,3-DPG will also yield measurement of oxygen saturation

which is accurate [11].

For normal hemoglobins the influence of pH on the oxyhemoglobin dissociation curve has

been well characterized and has been reconfirmed; the effect of P^q at fixed pH has been

established. The characterization of oxygen hemoglobin affinity when changing the ratio

between 2,3-DPG and hemoglobin concentrations has been completed as well (fig. 2). Availability of all of these relationships has enabled construction of an empirical nomogram

which predicts oxygen saturation which results at 37 °C. This nomogram relates the logarithm

of the Pp. required to give 50 percent saturation to changes in pH, Pp^, , and 2,3-DPG/Hb

U2 LU2

ratio (fig. 3). The nomogram can be used to calculate any of the parameters when the other

three are known, and could be used as a means of indirectly measuring 2,3-DPG if that were

of interest.

Figure 2. Values of log P50 of human blood vevsus 2,3 DPG/Hb ratio at zero

Ppp, and Ppp, = 41 mmHg.

155
I I J \ I I I I U7.0 7.1 7.2 7.3 7.4 7.5 7.6 7.7 pH

Figure 3. Values of log P50 as a function of pH and P^^ : a) normal

human blood at 37 °C; b) human blood with no 2,3 -DPG;^c) human

blood with 2,3-DPG/Hb ratio = 0.75; d) human blood at P^q = 0 and

2,3-DPG/Hb ratio = 2.3 and Pp^ = 41 mmHg and 2,3-DPG/Hb ratio = 2.5.

The work which led to the data which supports this nomogram was carried out by a group

comprised of researchers from Instrumentation Laboratory, Inc., the University of Milan and

University of California at Santa Barbara. The aim of the work was to extend studies which

have been performed on the effects of these hemoglobin effecters from diluted and purified

hemoglobin solutions to whole human blood under near physiological conditions of temperature,

Ppp, , pH, etc.

The oxygen dissociation curve (ODC) for a concentrated (ea. 25 percent) sample of

human hemoglobin in 0.1 M KCl serves as a basis for determining the effect of the individual

species. These ODC's were obtained using the "open" tonometric technique at a constant pH

of 7.22. This pH was chosen since it has been demonstrated [15] that an extracellular pH of

7.4 closely corresponds to an intracellular pH of 7.22. Thus, at the same intracellular

pH only two cof actors, CO2 and 2,3-DPG are required to make the ODC of a hemoglobin solution

correspond to the ODC of whole blood.

156
Table 1. Values of log P50 as a function of pH and

Pp-, interpolated through the data.

2,3-DPG

Hb

P^«

CO2

pH: 7.0 7.2 7.4 7.6

0 1 .260 1 .148 1 .038 0,.926

0 20.874 1,.285 1.820 1,.154 1,.088

41.535 1,.328 1.274 1 .220 1,.166

57.794 1 .342 1.305 1 .272 1,.232

0.75 0 1,.528 1.432 1 .336 1,.240

41.535 1,,.545 1.468 1 .396 1,.316

0 1,.575 1 .484 1 .394 1,.306

1 20.874 1,.584 1.496 1 .412 1,.330

41.535 1 .592 1.512 1 .436 1,.356

57.794 1,.604 1 .530 1 .456 1,.380

2.3 0 1 .716 1 .628 1 .540 1,.452

2.5 41.535 1 .756 1.672 1 .592 1 .512

The position of the ODC along the P^, axis will be described in terms of its P50 (Pn U2 U2 at which the hemoglobin is 50 percent saturated with oxygen). Plotting log P50 versus

pH yields straight-line relationships which are a function of P^g and 2,3-DPG. Table 1 indicates the values for log P50 interpolated from the data generated at varying levels of

P^Q^ and 2,3-DPG.

These data can be either represented in a nomogram or mathematically derived from an empirical formula. Both the nomogram and the empirical function allow calculation of pH,

Pqq » log P50 or 2,3-DPG/Hb ratio when the other three are known. The final monogram

(fig. 4) has been completed by adding a scale for the P^, and a corresponding scale for S^, , U2 U2

assuming that the relation between Pq and Sq conforms to the classical Hill's equation [12].

100 - = log K + n log Pp

U2

At 50 percent saturation log K = -n log P50. The value of n was assumed to be 2.7. Clearly

this equation can only be assumed to give a rough estimate of the relation between S^. and P^

U2 U2

Other workers have also found that n varies for normal human blood when CO2 and 2,3-DPG are varied [13]. With these limitations in mind, it is possible to obtain the function Sq^ vs. Pq^

derived from Hill's equation at a definite value of log K = (-n log P50).

157
46-

42 -

40

38 - 36-

3 4-

3 2- -1.5 30

2 8--

2 6-

.-1.4

2 4-

2 2-'

20--1.3

18-.

1 6-_

1 2 -

1 0 - P50 (mmHg)

-1.6

-1.2

-10

log Pso

® REF Pc02

1 0 -

SO2 (•/.)

-- 1 0

3 0

4 0

5 0

6 0

7 0 --80

9 0

9 5

REF 2.3 PPG ^

Hb

20-

--7. 8

3 0-

40 -

50-

6 0' 70-

80-

9 0

• 100-

12 0-

1 40- 16 0- 180-

20 0-

P02

( mmHg)

-7.

7

-7.6

7.5

-7.4

-7.3

-7.2

-7.

1

-7.0 PH

Figure 4. The oxygen dissociation curve and P50 nomogram.

Use of the P50 nomogram is relatively straightforward (fig. 5). If it is required to find log P50 for values of pH = 7.4, P^q = 40 mm, G = 1.0 the steps are as follow: (1)- Line 1, connecting P^q with Ref. P^q is traced; 2) G = 1.0 and Ref. 2,3-DPG/Hb are connected

by line 2; (3) A straight line connecting pH = 7.4 with the log P50 scale is traced through

point of intersection of lines 1 and 2. The value of P50 = 26.7 mm is read on the appropriate scale. As shown in figure 6, the value of G may be established by knowing P50, P^q , and Pn . U2

Figure 7 indicates the means by which an GDC may be- constructed by knowing P50, P^q , and Pf. . U2

By inserting values of G, pH and P^q into the mathematical expression which is

expressed below, values of log P50 can be directly calculated.

158
1 2

1 0 - Pso (mmHg)

4 6-

1,1.- 42 -

40

3 8-

3 6-

3 4

3 2-

30 -

2 8--

2 6-

2 4-

2 2-

20

1 8 -

1 6--1.2

-16

-1.5

-1.4

•1.3

-1.0 log Pso

Hb

2 0-

30-

40 -

5 0-

60 -v^O-l \

8 0 90-

100-

12 0-

1 40 160-

180-

200

P02

( mmHg)

7. 8

-7.7

7.6

•7.5

•7.4.

7.3

7.2

-7.

-7.0 PH

Figure 5. Use of nomogram to determine P50 from measurement

of 2,3-DPG/Hb ratio, P^q and pH (see text).

log P50 = 1.9729 x 10-3G2pH P^q - 0.6002 x 10-3G pH P^^q^ + 3.9274 x 10-3

pH P^Q^ + 206.5 X 10-3G2pH - 247.79 x 10-3G pH - 410.31 x 10-3

pH + 15.85 X 10-3G2 P + 1.157 x 10-3G P„ - 25.948 x 10-3

LU2 1,1)2

P(>Q^ - 1.6337 G2 -h 2.2427 G + 4.125.

Evaluation of errors which will occur in values for P50 at pH = 7.0, 7.2, 7.4 and

7.6; Pj,Q = 0, 20, 41 and 57; and G = 0, 0.75, 1, 2.3 and 2.5 has been carried out by use

159
Figure 6. Use of nomogram to determine 2,3-DPG/Hb ratio from

measurement of P50, Pr-n. and pH (see text).

of the mathematical expression. Table 2 gives the results obtained for log P50 using the

mathematical expression and can be compared with the results interpolated from the empirical

data. In the range of values which encompasses most physiological and pathological conditions, the maximum discrepancy of log P50 is 0.037 at a value of log P50 = 1.632 which is

about 2 percent. In the range of G from 0.1 to 1.0 the maximum deviation is 0.011 for log

P50 = 1.328. Overall, it is likely that the error in the determination of log P50 or G in

the range of the variables indicated in the nomogram does not exceed 3 percent of the

desired quantity.

160
1.6

-1.5

4 6-

4* -

42 -

40

3 8-

36-

3 4-

3 2-

30 -

2 8-

26 -

2 4-

2 2-

20--1.3

1 8 -

' «--1.2

1 2 -

1 0 - Pso (mmHg)

1 0 -

SO2

-20

-- 3 0 2 0--

4 0

5 0 60

7 0 ® REF PCO2 30-

+ 80

9 0

9 5

40 •

5 0-

70-

80-

90-

10 0-

12 0-

1 40-

1 60- 180- 200-

-10

log Pso

( mmHg)

-7.0 PH

Figure 7. Use of nomogram to construct oxygen dissociation curve

from measurement of P50 (see text).

This work is by no means final, and use of the nomograms can safely be applied only to blood with normal hemoglobin concentration in the absence of pathological hemoglobins. We

have no data available at this time, for instance, on blood from individuals who are anemic.

In these individuals G is altered (mainly by decrease in hemoglobin concentration) from its

normal value [14].

The purpose here is to demonstrate the requirements for an accurate assessment of

oxygen saturation of blood samples. The several species which will affect the relative

hemoglobin-oxygen affinity must be accounted for in order to accurately make an assessment.

Direct photometric measurement of oxygen saturation, combined with measurement of

blood gas parameters pH, P^q , and Pq yields the required information for correct assess- ment of hemoglobin-oxygen affinity.

161
Table 2. Values of the function log P50 = f(pH, Pp^, , G)

derived from equation. 2

L ) 0— Ur

b

Hb

p^C02nH 7 n 7 ? 7 &. 7

0 1 1 1

1 .1/1 ^ 71 1 . ucjy 1 . UU/

0 20.87 1,.285 1 .219 1 .154 1 .088

41.53 1,.317 1 .267 1 .218 1.169

57.79 1,.;342 1 .305 1.269 1.232

0.75 0 1,.528 1 .432 1 .336 1 .240

41.535 1,,545 1 .469 1.392 1 .316

0 ]

,

573 1 1 . o^C

1 20.87 1

.

,584 1 .499 1 .415 1 .330

41 .53 1,,595 1 .516 1 .437 1.358

57 7Q 1

1 1 1 1 /ICC

2.3 0 1.,426 1 .448 1 .471 1.493

2.5 41.53 1.,625 1 .595 1.565 1 .535

References

[I] Shappel, S. D., Lenfant, C. J. M., Anesthesiology, 37, 127 (1972).

[2] Fallon, K. D. et at.. Analyzer, 4, 11 (1974).

[3] Valeri, C. R., Collins, F. B., X){l\lth Scientific Meeting of the Blood Research Institute,

p. 397.

[4] Metcalfe, J. et al. , Circulation Research, 25, 47 (1969).

[5] Siggaard-Andersen, 0., et al. y Clin. Chim. Acta, A2, 85 (1972).

[6] Malenfant, A. L., et al., 20th National Meeting, AACC (August 19, 1968)

[7] Astrup, P., Forsvarsmedicin, S_, 199 (1969).

[8] Van Slyke, D. D., Neil, J. M., J. Biol. Chem. 61, 523 (1924).

[9] Clauvpl, M., Schwartz, K., Clin. Chem. ]±, 3, 253 (1968).

[10] Eilers, R. M., Amer. J. Clin. Path. 47, 212 (1967).

[II] Muletti, A. et al.^ Physiological Basis of Anaesthesiology , M. M. Mushin, et al., ed.

p. 95 (Piccin Medical Books, Padua, 1975).

[12] Hill, A. v., Proc. Physiol. Soc. J3., 4 (1910).

[13] Garby, L., et al. , Acta Physiol. Scand. 84, 482 (1972).

[14] Astrup, P., Rorth, M., J. Clin. Lab. Invest. 31, 311 (1973).

[15] Agostoni, A. et al. , Science 182 . 300 (1973).

162
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

PROBLEMS ASSOCIATED WITH THE DEFINITION OF MEASURED AND CALCULATED

QUANTITIES IN BLOOD pH AND GAS ANALYSIS

Robert W. Burnett

Clinical Chemistry Laboratory

Hartford Hospital

Hartford, Connecticut 06115, USA

For most students, the subject of acid-base balance has the reputation of being a singularly complex and confusing one. This seems to me to be due in large part to the

multitude of methods which have been used and quantities which have been defined to express

the acid-base status of an individual. Much confusion has been caused by the existence of

several methods, that might all be in common use, and that measure similar, but not exactly

the same, quantities. Another obstacle to understanding is the existence of several terms

to denote a single quantity, or group of quantities that differ from each other only in a trivial way. Moreover, many modern and widely used textbooks contribute to the problem by

devoting a disproportionate amount of space to a discussion of methods and quantities that

are generally agreed to be outmoded, and giving inadequate space to a discussion of legitimate differences of opinion or differences in usage which exist in this field today.

My objective in this presentation is to outline some of the problems that exist in the area of definition of quantities used in blood pH and gas analysis. I think it will be

clear that many non-trivial questions exist that will not, of course, be answered at this

meeting, but that do deserve our careful consideration.

The quantities that perhaps cause the fewest problems, in terms of their meaning,

happen to be the three most commonly measured, namely pH, P^q and Pq . This is not to imply that no questions exist in the definition of these. The pH of an unknown solution is

defined and determined in an operational way by comparison with the pH of an accepted

standard solution. Systematic biases in blood pH measurements between laboratories often occur however, even when both are using the same pH standard. These can be the result from

differences in the type of liquid junction, differences in bridge solution composition,

different temperatures of the salt bridge, or junction potential differences due to the

so-called "cell effect" seen with whole blood. While it would be worthwhile to work

toward eliminating such inter-laboratory variations, it must also be remembered that ac- curacy as a concept in pH measurements with whole blood is much more difficult to deal with

than with most other measurements. Quoting from Bates' monograph [1]^,

"...the experimental pH can never be an exact measure of either

the concentration or the activity of hydrogen ion It is

safe to say that no quantitative interpretation of measured

pH values should be attempted unless the medium can be classified as a dilute aqueous solution of simple solutes."

The need, therefore, is to establish a convention for the measurement of pH in blood, so that consistent and reproducible numbers can be obtained in different laboratories and with

different instruments.

With regard to the gas tensions, P^q and Pq , definition is not a problem, but again.

^Figures in brackets indicate the literature references at the end of this paper.

163
consistent measurement of these quantities is often a problem. Differences in electrode

response time at different gas tensions and differences in measured gas tensions between

gas and liquid standards are two common reasons for inconsistent and inaccurate measurements.

These problems are actually ones of standardization, and perhaps of instrument specifications,

and should be discussed later in this symposium.

Another measurement which is routine in most clinical chemistry laboratories is CO2

content, but this term has been used in connection with many different methods, not all of

which give equivalent results. The determination has been made by measuring the volume of

CO2 released from acidified serum or plasma or measuring the color change in a weakly

buffered indicator solution on reaction with an aliquot of this CO2, for which different

indicators and buffers have been used. The measurement can be made with or without the

addition of base as a preservative, with or without equilibration of the sample to a known

CO2 tension, and after centrifugation of the blood at either body temperature or room temperature. Clearly, there is a need to adopt a single definition for this quantity, and

then to determine the extent of systematic bias that is present in the several methods that

are available.

Some other quantities that are often measured directly, relate to the oxygen transport

status of an individual. These too can have more than one meaning, depending on the particular method of measurement. Hemoglobin may be measured as oxyhemoglobin alone, or in

combination with reduced hemoglobin, carboxyhemoglobin or methemoglobin. Likewise, oxygen

saturation, which is actually a derived quantity, can be calculated with or without terms to include the effect of carboxyhemoglobin or methemoglobin. Spectrophotometric methods with

measurements at one, two, three or four wavelengths have all been used, and will give

different results with the same blood sample if carboxyhemoglobin or methemoglobin is present. To add to the confusion, there is also the very common practice of calculating

oxygen saturation from a measured pH and P« . The equations or nomograms used for this

U2

purpose are constructed from a standardized oxyhemoglobin dissociation curve--not necessarily

the same one--which is only valid under a single set of conditions, including CO2 tension,

2,3-diphosphoglycerate concentration and concentrations of carboxyhemoglobin and methemoglobin.

A systematic bias is always present when all these conditions are not met by the specimen of

blood being analyzed. Moreover, it is quite common to find that two of these nomograms give

significantly different answers, particularly at high oxygen saturations.

Increased interest in the actual oxyhemoglobin dissociation curve which is obtained

in a given sample of blood has led to the more common calculation of P50, the oxygen tension

which corresponds to 50 percent saturation. Here too, careful definition and consistent

conventions are needed in order to avoid inter-laboratory differences. First, the meaning

of oxygen saturation must be established, and second, agreement must be reached on the

values for constants to be used in the Hill equation, Adair equation, or other relationship

used to calculate P50.

A final quantity of interest related to oxygen transport is oxygen content. This may

be measured directly by either the manometric Van Slyke method or by a technique in which

all the oxygen in the sample is reduced electrochemically . Alternatively, oxygen content

can be calculated as (K x total hemoglobin x oxygen saturation). This calculation will be

meaningful only if two conditions are met. First, total hemoglobin and oxygen saturation

must be expressed in units that are consistent with each other, and second, a value for K must be agreed upon. The theoretical value for K can be calculated from the molecular

weight of hemoglobin A and the molar volume of oxygen at standard temperature and pressure,

and is equal to 1.39 ml O2 per gram of hemoglobin. This value is not always considered the

best to use however, due to reports of an unexplained systematic bias between calculated and

measured oxygen contents. Again, it is clear that a convention needs to be adopted in order

to make it possible to obtain values that are meaningful and consistent among laboratories.

Returning to the definition of acid-base status, I have described some problems having

to do with measured quantities and I would like to discuss now some factors related to derived quantities. These are sources of confusion and inconsistency, again because a given

quantity is often calculated using different definitions or different values for constants.

First, consider the quantities that may be calculated using the Henderson-Hassel balch

equation. If we assume that accurate values of pH and P^q are available, the HendersonHasselbalch equation allows calculation of plasma bicarbonate and CO2 content. The condition

164
that must be met in order to insure consistent results among laboratories is, of course,

that a single value for the dissociation constant of carbonic acid be adopted and a single

number for the solubility coefficient of CO2 in plasma be used. Careful experimental work

has resulted in apparently accurate values for these constants, so this should not present a serious problem.

A more difficult quantity to reach a consensus on in the past has been base excess. This reflects the fact that base excess can be defined in relation to various fluid com- partments in the body, including plasma, whole blood, interstitial fluid, or some combination

of these. The parameter that is necessary to define base excess mathematically is the

buffer capacity of the fluid. Although an early definition of base excess was based on the

buffer capacity,of whole blood, the most widely accepted definition at present is based on the buffer capacity of extracellular fluid (i.e., whole blood together with interstitial

fluid) which has been determined in vivo by several groups. Unfortunately, although most

workers who are active in the field of blood pH and gas analysis seem to agree on the

advantages of the so-called extracellular fluid base excess, several widely used textbooks,

published within the last five years, do not mention this quantity at all, but persist in

using the older definition. Clearly, when there is indeed consensus among experts, then the

conventions that have been agreed upon must be effectively publicized if they are to influence

current teaching.

Finally, reference should be made to the many other terms that have been used in the

past and in some cases continue to be used to describe various aspects of the acid-base

status of an individual. This list includes alkali reserve, bicarbonate reserve, buffer

base, buffer anions, CO2 combining power, acid-base ratio, standard bicarbonate and CO2

capacity. Some of these terms are redundant; some are simply no longer useful. All of

these terms may be found in recently published textbooks. The need for derived quantities

can perhaps be placed in perspective by keeping in mind that acid-base status is completely

determined if only two quantities are known, for example, pH and P^q . The justification for

continuing to use other quantities must rest on a real practical advantage in the understanding

or treatment of acid-base abnormalities.

In summary, present usage of both measured and derived quantities for blood pH and gas

analysis is often inconsistent and confusing. This area of clinical chemistry would benefit

greatly from international agreement on the mathematical definition of quantities and on values for appropriate constants. It would then also be necessary for the areas of agreement

to be widely publicized. One change which I feel would be a significant improvement would

be for scientific journals to adopt uniform criteria for nomenclature and definition of

terms in this area. This has been done successfully by other journals in the areas of

spectrophotometry and statistics.

The most important reasons for precise definition of terms and agreement on conventions

are that they will make possible measurements whose results are reproducible among laboratories, it will greatly facilitate the interpretation and comparison of data generated in

different laboratories, and it will make the job of those who apply the results of these

measurements to patient care less complex. The desirability of achieving these goals is

something that can certainly be immediately agreed upon.

Reference

[1] Bates, R. G. , Determination of pH, 2nd edition, p. 103 (John Wiley & Sons, New York,

1973).

165

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

ARTERIAL AND VENOUS BLOOD SAMPLES IN ACID-BASE BALANCE

William H. Austin

Department of Research

Maine Medical Center

Portland, Maine 04122 USA

The field of pH and blood gas measurement has stood as the stepchild of anesthesiology,

pulmonary medicine, nephrology, endocrinology, and research itself. It has never gained the

status of a full-fledged discipline. Because of these multi-faceted relationships, different

concepts concerning sampling, measuring, calibrating, reporting, and interpreting have

arisen.

A simple example of this is seen in the use of PCO2 by the respiratory physiologists

and PCO2 by others. In this simple example, the symbol "P" is used to mean pressure, whereas

it is also used in the designation of a constant, such as pK' . The respiratory physiologists

have achieved some measure of uniformity by agreeing upon symbols and nomenclature for

international use. However, "a", the solubility factor (ml C02/ml plasma) and "S", the

coefficient (mm PC02/ml plasma) are still used interchangeably.

This could be logically extended to the entire field of acid-base balance, as it is used in the various disciplines mentioned above. The logical way to achieve consistency and

uniformity is to make pH and acid-base balance a specialty in itself. This would be done

with the understanding that this area would not create isolation.

In doing this, there must be general agreement in the many areas of this subject. This

workshop, I believe, has addressed itself primarily to this point. Many of these areas are being discussed here, and no single paper can encompass the entire field. One area in which

there has been debate is that of the site of the sample taken to be measured [1-4]^. It is obvious that arterial blood is essential when dealing in matters of pulmonary physiology;

however, there has been no universal agreement regarding the role of venous blood in the

other physiologic areas. Gambino [5-7] has been an outspoken advocate of venous blood, and

has used venous blood where pH and PCO2 have been of primary concern. For others, the

addition of PO2 is essential, and dictates the origin of the sample. I am not aware of any

clear definition of the role of venous blood measurements, despite consideration of mixed

venous (brachial and hand) blood. Much has been written about "arterial izing" venous blood,

and using various techniques to ensure the uniformity of venous samples. There is evidence

of the similarity of venous and arterial blood when drawn from the brachial vessels under

controlled conditions; however, we have encountered some dissimilarities despite rigid

sampling and measuring techniques. At times, there may be very little difference, but it

has been difficult to determine the conditions which produce A-V dissimilarities.

Under controlled conditions, the PCO2 of arterial blood will change dramatically, relative to venous PCO2, with alteration in ventilation. Figure 1 indicates this point. In the studies illustrated here, blood samples were taken from the inferior vena cava and distal

aorta in each of eight dogs, before and following a twofold increase in minute ventilation

[8], The dogs were intubated, and had the appropriate vessels cannulated, with a one-hour

steady state during which the PCO2 measured by electrode was kept within ±3-4 mm. The

figures show changes which occurred at various intervals in both the arterial and venous

systems. The rapidity of the fall of PCO2 in arterial blood, and the lag in decrement for venous blood illustrates just one facet of this problem. The lowering of the arterial

partial pressure of carbon dioxide is consistent with the change in ventilation, and the venous pressures reflect changing carbon dioxide levels at the tissue level.

1 Figures in brackets indicate literature references at the end of this paper.

167
Figure 1. The average and mean changes

in venous {•) and arterial (l) PCO2

resulting from a two-fold increase in

minute ventilation. (Reproduced with

permission of The Journal of the Maine

Medical Association.

)

0 2 4 6 8 10 12 MINUTES

Each value is appropriate for the situation in the unsteady state, and there is no "right" or "wrong" value. There is, then, the logical extension of this simple maneuver to the patient. There is no axiom which states that venous and arterial blood should have a fixed difference in either sample, or will represent one physiologic set of circumstances.

The misconception occurs when one tries to correct arterial for venous, or venous for arterial, blood sample values. The major issue is, "which sample of blood reflects the clinical

state, and does it provide the information necessary for the appropriate treatment of the

patient?"

Table 1 illustrates the A-V differences which were measured in fifty consecutive pa- tients with various disease states. Numbers 32,35,36,37,39,42, and 46 suggest abnormalities

if the venous blood is observed. This is not the case. No statistical analysis can be made

here, but neither the nature of the illness nor its severity appears to correlate with the

differences observed. That there are differences is no surprise to some, but may be challenged by others. Yet the facts must speak for themselves. Although there is no systematic,

objective proof of the matter, the arterial values appear most consistent with the patient's

clinical state. In those instances where one would not anticipate abnormalities, venous blood measurements suggest that there are. All this is done using the criteria set forth

for drawing and measuring pH and blood gases [6]. The reasons for these differences are not

apparent, and no obvious technical errors were made in the determinations. The conditions

for sampling and measuring have been made elsewhere, and strict adherence to these has been

used. Reliability of calculated variables has been stated [1-5]. Transient changes in venous blood flow may account for some of these differences; however, it was impossible to predict when the wide A-V differences would occur. It could be argued that the venous measurements in reality reflect the tissue state, which is perhaps desirable, but this may

only be representative of the area from which the sample is taken. Here a mixed venous blood sample would be the logical site, but not practical to obtain.

I have been an advocate of the use of venous blood, but over the years I must concede

that I find myself relying more and more on arterial values, even in disease states which

are not primarily of respiratory origin. Now technicians can easily obtain samples, which

obviates a previous objection, and frequent punctures have not been a problem. The circumstances are such that the calculated CO2 content, or bicarbonate, made when the pH and PCO2 are measured, is generally quite similar, or differs by a consistently small amount. This means to me that the Henderson-Hasselbalch equation is giving correct CO2 contents, or bicarbonates, but that the pH and PCO2 values may be misleading. On other occasions, abnormal

pH's and PC02's were found on venous blood, when no disorder existed after subsequent arterial

samples were taken. I have become more and more convinced of the reliability of arterial pH

and PCO2 as representative of the severity of the disease state, even though venous CO2 con- tents are usually consistent. This appears to apply more to PCO2 measurements which we have

found to be extremely variable in venous blood.

168
Table 1. A-V blood gas differences in fifty patients.

pH PCO2 TCO2

A V A A V A A V A

1

1 • 7 46 7 43 03 41 .

46.6 c0 . 1 29. 6 310 1 .4

9c • 1/ • 41*T 1 7 37 n4 43.5 49.0 r0.

rb 27. 7 28. 5 0.8

0 • 7 4fi 7 43 • uo 40.4 44.

3. 9 28. 8 30. 0 1 .2 4. 7.41 7.38 .03 42.0 45!o 3. 0 26! 7 26! 8 o!i

5. 7.44 7.41 .03 36.

41 .5 4. 9 25. 0 26. 4 1 .4

6. 7.46 7.43 .03 37.7 41.5 3. 8 26. 9 27. 6 0.7

7. l.Zl 7.35 .02 53.4 58.0 4. 6 31. 1 32. 1 1.0

8. 7.46 7.43 .03 29.8 31.3 2. 5 21

3 20. 8 0.5

9. 7.52 7.46 .06 38.5 42.5 4. 0 31. 6 30. 2 1.4

10. 7.43 7.39 .06 42.5 48.5 6. 0 28. 2 29. 5 1.3

1 1 • 7 dl 7 35 06 36 7 41 n 5. 3 23 3 23. 2 0.1

1\ 9c . 7 /in 7 3Q m *+VJ . u 47 n 6 4 28 5 3.3

1 0 • 7 4^ 7 3Q 04 41 4 4ft n 6. 6 77 5 29. 2 1 .7 14. 7.41 7.37 .04 49.2 56.0 6. 8 31. 3 32! 6 r.3

15. 7.44 7.42 .02 42.0 46. 5 4. 5 28. 5 30. 3 1 .8 16. 7.42 7.39 .03 48.0 51.0 3. 0 31. 2 31. 0 0.2

17. 7.49 7.44 .05 46.0 52.6 6. 6 35. 2 35. 4 0.2

18. 7.40 7.40 .00 49.0 50.0 1. 0 31. 0 31

1 0.1

19. 7.38 7.35 .03 74.0 79.0 5. 0 44. 2 44. 2 0.0

20. 7.41 7.36 .05 40.5 52.5 12. 0 25, 8 30. 0 4.2

21 7 4"? 7 38 05 47. 5 57.

10.0 29 2 34 3 5.1 22. 7 ^R 7 34 04 41 .

50.0 8. 5 24 7 27 1 2.4

23 7 43 7 39 04 46. 5 56.0 9. 5 28 6 34 1 5.5

24. 7.41 7.35 .06 37^4 46!4 9. 0 23 8 25 7 1.9

25. 7.49 7.45 .04 30.5 39.0 8. 5 23 4 27 1 3.7

26. 7.42 7.37 .05 47.0 57.0 10.0 30 6 33 1 2.5

27. 7.40 7.37 .03 59.0 70.0 11. 0 36 8 40 8 4.0

28. 7.48 7.36 .12 35.5 48.6 13. 1 26 2 27 8 1.6

29. 7.37 7.32 .05 60.0 74.0 14. 0 35 1 38 3 3.2

30. 7.42 7.36 .06 36.5 54.0 17. 5 23 8 31 9 8.1

31

7. 53 7.40 . 1

36.0 52.0 16. 0 30.2 32. 5 2.3

32. 7.40 7.33 07 43.

62.19 0 27. 4 32. 8 5.4

33 7.45 7. 38 • 07 33.5 47.0 13 5 23. 4 28. 1 4.7

34. 7.46 7^42 !o4 36! 3 48!4 12 1 25. 9 31! 5 5.6

35. 7.43 7.29 .14 36.8 60.0 23 2 34. 2 29 3 4.9

36. 7.53 7.45 .08 34.0 51.0 17 0 28 6 35 6 7.0

37. 7.39 7.34 .05 43.7 61.0 17 3 26 6 34 5 7.9

38. 7.17 7.14 .03 77.8 102.0 24 2 30 0 36 0 2.7

39. 7.37 7.33 .05 42.0 57.0 15 0 27 8 32 0 4.2

40. 7.25 7.17 .08 51.5 74.0 22 5 23 5 28.0 4.5

41

7. 51 7 45 06 48 8 60 0 11 2 39 8 44 0 4.2

42. 7.40 7 34 • 06\J\J 39 4 57 0 17 .6 25 4 32 0 7.5

43 7 43 7 37 06 36 6 59 0 22 4 25 4 35 0 9 6 44. 7.49 7.40 .09 36.4 58.0 21 .6 28 8 36 0 7.2

45. 7.43 7.38 .05 50.5 64.0 13 .5 34 5 39 0 3.5

46. 7.41 7.34 .07 37.1 53.9 16 .8 24 0 30 6 6.6

47. 7.38 7.25 .13 23.3 35.7 12 .4 14 2 16 0 1 .8 48. 7.29 7.29 .00 87.0 98.0 11 .0 41 3 48 0 6.7

49. 7.34 7.31 .03 29.1 41.1 12 .0 16 2 22 0 5.8

50. 7.37 7.24 .13 28.5 51.4 22 .9 17 0 23 3 6.0

169
On the basis of the small amount of data presented here, clinical experience, and

information derived from the literature, I would generally favor arterial blood as a reference

point for all acid-base disturbances, even those of metabolic origin. This is not to exclude

the use of serial venous pH and PCO2 measurements made with the reservations mentioned

above. I believe that the door should be kept open on this point, and other experienced

opinions obtained. Again, this may be a matter of judgement--reinforcing my original

concept of the field of pH and blood gas measurement.

REFERENCES

[1] Forster, H. V., Dempsey, J. A., Thomson, J., Virduk, E., and doPico, 6. A., Estimation

of arterial PO2, PCO2, pH and lactate from arterialized venous blood, J. Appl. Physiol.

32, 134 (1972).

[2] Samet, P., Linhart, J., Barold, S. , and Hildner, F. , Reliability of mixed venous blood

for the measurement of blood gas parameters, J. Thor. Card. Surg. 58, 131 (1969).

[3] Sutton, R. N., Wilson, R. F., and Walt, A. J., Differences in acid-base levels and

oxygen-saturation between central venous and arterial blood. Lancet, 2^, 748 (1967).

[4] Paine, E. G. , Boutwell, J. H., and Soloff, L. A., The reliability of "arterialized" venous blood for measuring arterial pH and PCO2, Am. J. Med. Soi. 2M, 431 (1961).

[5] Fleischer, W. R. and Gambino, R. S., eds., Blood pH, PO2 and Oxygen Saturation, 223 pp. (Am. Soc. Clin. Path., Chicago, 1972).

[6] Gambino, S. R. , Comparisons of pH in human arterial, venous, and capillary blood.

Am. J. Clin. Path. 12, 298 (1959).

[7] Nahas, G. G., ed. , Report of The Ad Hoc Committee on Methodology, Current Concepts

of Acid-Base Measurement, Annals of the New York Academy of Science, New York City

274 pp. (Edward Bros., Inc., Ann Arbor, Mich., 1966).

[8] Austin, W. H., The use of arterial or venous blood in acid-base balance, J. Me. Med. Assoc. 236 (1970).

170
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Gaithersburg, Maryland, July 7-8, 1975. Issued June 1977.

BLOOD SAMPLING AND HANDLING IN THE DETERMINATION OF BLOOD

pH AND BLOOD GASES

Arthur H. Richards

Division of Laboratories and Research

New York State Department of Health

Albany, New York 12001 , USA

The collection and handling of blood for the determination of blood pH and blood gases

require that the source (arterial, capillary or venous), sampling site, collection materials

and storage conditions be precisely and clearly defined in order to ensure that a proper

sample is presented to the analytical unit.

Although the source of the blood must sometimes be selected on a purely pragmatic

basis, arterial blood is the preferred type of sample and will reflect the true status of

the pulmonary system with respect to pH and blood gas content. Collection of the sample,

usually from the brachial, femoral or radial arteries, requires a high degree of expertise.

The patient must be relaxed and the blood drawn with a minimum of trauma. Multiple arterial

samples are not recommended, but, when several samples might be required, a Cournand needle

or Seddinger catheter can be employed.

Blood collected via a vasodilated capillary bed has been reported to be virtually

equivalent to arterial blood [1-3]^. The proper technique for collection of a capillary

sample (usually from the earlobe, finger or heel) is to warm the site thoroughly by immersion

in warm water, then allow the blood to flow into a capillary tube without impairing or forcing the blood flow. The results on arterialized capillary blood have been reported to show good agreement with samples taken from a true arterial source, although a notable

exception has been reported: P02 arterial versus capillary measurements were reported to be

in poor agreement [5].

Venous blood can be employed for the routine measurement of CO2 content. However, it

is well documented that oxygen tension can be considerably different than in arterial blood

[10-12,14]. Some reports have stated that reasonably satisfactory samples can be obtained

by the application of heat to the source to create "arterialization. " One monograph recommends

that blood be taken from a vein on the back of a hand that has been subjected to 46-47 °C

heat for 10-15 minutes [11]. The use of a light tourniquet is acceptable, but it should be- released during the drawing period. It has been reported that, when these conditions are closely followed, good approximation of arterial blood is achieved for pH and PCO2, although

P02 results can be significantly different.

The collection device most frequently recommended is the glass- type syringe with matching

barrel and plunger. A light coating of heparin anticoagulant solution should be used for

lubrication as well as for preventing clot formation. Samples are always drawn under

anaerobic conditions. Plastic syringes cannot be used. as the porosity of the plastic creates

a potential for CO2 and O2 to diffuse into the plastic. Vacutainers containing sodium

heparinate have been shown to be acceptable [8]. Other anticoagulants are unacceptable, and

oil or lubricants must never be used.

It is recommended that the blood sample be placed in ice water if analysis cannot be

performed within 20 minutes [2]. It is reported that blood pH will not change more than

0.01 pH units when kept at 4 °C for up to four hours. PCO2 is a more labile determination.

Figures in brackets indicate the literature references at the end of this paper.

171
for which it is recommended that the sample stored on ice be analyzed within 30 minutes.

The sample must be kept from exposure to room air by inserting the collection needle into a cork or rubber stopper or by using any appropriate covering device. The sample must be

remixed just prior to proceeding with the determinations. Mixing is best accomplished by

adding a few drops of mercury to the sample, both to assist in completely mixing the

anticoagulant and remix to the sample at a later time.

Numerous studies have compared the results of tests in which arterial blood and venous

and capillary bloods, arterialized by a variety of techniques, were used. Paine, et at. [4]

investigated a group of 29 hospitalized patients by making a comparison between arterial

samples, venous samples and arterialized venous samples. The arterialized venous samples were obtained by prewarming a vein on the dorsum of the hand for 15-20 minutes and were collected with a tourniquet applied to the wrist. The pH values were on the average 0.022

units lower than arterial blood, while the PCO2 was 2.86 mm Hg lower. It was concluded that

this arterialized sample is satisfactory for the estimation of arterial pH.

In a separate study, Jung, et al. [14] made comparisons of arterial and venous bloods

collected simultaneously with both syringes and vacutainers. Capillary blood was also used

in some of these studies. The pH, PCO2, bicarbonate, total buffer base and base excess were

all satisfactory on capillary blood, but P02 was not acceptable. The study stated that

blood obtained by Vacutainer was not adequate for any of the determinations, but haparinized

venous blood could be substituted for arterial blood except for the P02 determination.

A study by Spock, et al. [13] showed that P02 could be accurately determined on capillary

samples if a proper technique was employed. An excellent correlation with arterial samples

was shown in this study. These samples were obtained by thumb prick after the thumb had

been warmed at 40 °C. The samples were analyzed promptly and were representative of a number

of disease states.

In yet another study, Koch [11] has shown that capillary blood from the middle finger,

when arterialized by warming at 45 °C, was comparable to arterial blood for measuring pH,

PCO2 and bicarbonate concentration.

One of the favored collection sites for blood pH and blood gas determinations is from

the earlobe. Laughlin, et al. [6] have shown that by warming the earlobe with an electric

heater, P02 values could be obtained quite comparable with P02 values in arterial blood. In the same study samples from finger punctures gave less satisfactory P02 values than those

from earlobe punctures.

Langlands, et al. [7] compared P02, PCO2 and pH values brachial artery and earlobe

blood. Their results showed that values in arterialized earlobe blood corresponded very

closely to arterial blood.

Gambino [8] has utilized the earlobe as a site for collection of capillary blood. He described a rubber-cup collection system and compared his results with arterial values. No significant differences were observed between the capillary blood and arterial blood values

for pH, CO2 content, PCO2 and oxygen saturation.

Siggaard-Anderson, et al. [9] have done studies on capillary blood and state that pH

shows a standard deviation of 0.006 pH units and PCO2 accuracy of ± 2 percent.

In summary, it would appear that although the ideal situation is to obtain arterial

blood, venous or capillary blood can serve as a valid substitute providing the collection

site is carefully chosen and prewarmed and the specimen is carefully collected and properly

cared for prior to analysis.

References

[1] Ibott, F. A., LaSanga, T. S. , Gin, J. B., and Inkpen, J. A., in Clinical Chemistry

Principles and Techniques, R. J. Henry, D. C. Cannon, and J. W. Winkelman, eds. (Harper

and Row, Hagerstown, MD, 1974).

172
[2] Fleischer, W. R. and Gambino, S. R., in Blood pH, P02, and Oxygen Saturation (American

Society of Clinical Pathologists, Commission on Continuing Education, Chicago, IL,

1972).

[3] Gambino, S. R. , Astrup, P., Bates, R. G., Campbell, E. J. M., Chinard, F. P., Nahas, G.

G. , Siggaard-Anderson, 0., and Winters, R. , Report of ad hoc Committee on Acid-Base

Methodology. Amer. J. Clin. Pathol. 46, 376 (1966).

[4] Paine, E. G., Boutwell, J. H., and Soloff, L. A., The reliability of "arterial ized" venous blood for measuring arterial pH and PCO2, Amer. J. Med. Sci. 242 , 431 (1961).

[5] Adams, A. P., Morgan-Hughes, J. 0., and Sykes, M. K. , pH and blood-gas analysis. Anesthesia,

22, 575 (1968).

[6] Laughlin, D., McDonald, J., and Bedell, G., A microtechnique for measurement of P02 in

"arterial ized" earlobe blood, J. Lab. Clin. Med. 64, 330 (1964).

[7] Langlands, J. and Wallace, W., Small blood samples from earlobe puncture. Lancet, 1_, 315 (1965).

[8] Gambino, S., Collection of capillary blood for simultaneous determinations of arterial

pH, CO2 content, PCO2 and oxygen saturation, Amer. J. Clin. Path. 35, 175 (1961).

[9] Siggaard-Anderson, 0., Engle, K. , Jorgensen, K. , and Astrup, P., A micro method for

determination of pH, carbon dioxide tension, base excess and standard bicarbonate in

capillary blood, Scand. J. Clin, and Lab. Invest. 12^, 172 (1960).

[10] Adams, A., Morgan-Hughes, J., and Sykes, M., pH and blood-gas analysis. Anesthesia, 22,

47 (1968).

[11] Kock, G. , Comparison of carbon dioxide tension, pH and standard bicarbonate in capillary

blood and in arterial blood, Scand. J. Clin, and Lab. Invest. 17^, 223 (1965).

[12] Koch, G. , The validity of P02 measurement in capillary blood as a substitute for arterial

P02, Scand. J. Clin, and Lab. Invest. 2]_, 9 (1968).

[13] Spock, A., Lewis, M., and Albertson, T. , A microtechnique for measurement of oxygen

tension in capillary blood, J. Pediatrics, 68, 987 (1966).

[14] Jung, R. , Balchum, 0., and Massey, F., The accuracy of venous and capillary blood for

the prediction of areterial pH, PCO2 and P02 measurements, Amer. J. Clin. Pathol. 45,

129 (1966).

173

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

NON-ANALYTICAL SOURCES OF LABORATORY ERROR IN pH AND BLOOD GAS ANALYSIS

Jack H. Ladenson

Division of Laboratory Medicine

Departments of Pathology and Medicine

Washington University School of Medicine and Barnes Hospital

St. Louis, Missouri 63110, USA

Laboratory error includes more than just the analytical error associated with instrumental measurement. Properly, it should be defined as any error from the ordering of a test

procedure until the interpretation of results. In our experience, the non-analytical sources

of error are more severe than the instrumental and harder to document and eliminate [1]^.

The purpose of this presentation is to review some of these non-analytical sources of laboratory error in pH/blood gas measurements in an effort to alert both the laboratorian and

clinician to these problems.

1. Choice of Collection Site

A. Arterial versus venous blood

The differences between arterial and venous blood found by various studies are shown

in table 1. Comparisons of arterial and capillary blood obtained without any attempt to arterialize the capillary collection site are also shown. It is evident that venous blood

obtained from an arm vein has a lower pH than arterial blood. The magnitude of this

difference varied from study to study but appears to be '^0.05 pH unit. When capillary

blood is compared to arterial blood, discrepancies between the conclusions of the studies

are found. This is probably due to differences between the patients studied as will be

discussed in the next section. Only small differences between arterial and capillary

blood were found in studies of patients under general anesthesia [6,10,11], and were attributed to vasodilation and high cutaneous blood flow caused by anesthesia. These

phenomena have been reported to help to arterialize the capillary blood [8]. In newborns,

very poor correlations were found between arterial blood and blood obtained from an un- warmed heel [15].

The Pco2 of venous blood is higher than arterial blood. The exact magnitude of the mean difference varied from study to study (table 1) but appears to be between 5 and 10 mm

Hg. Again capillary blood agreed more closely with arterial blood than did blood from an arm vein. An important exception to this finding was the poor correlation of capillary

(unwarmed heel) and arterial blood for PCO2 in newborns [15]. Very large differences in

P02 and O2 saturation between arterial and venous blood have been observed [4,7] with

smaller differences between arterial and capillary blood [12,14].

The data indicates that for most purposes, substitution of venous blood for arterial

blood will not cause clinically significant differences for pH, may cause clinically

significant differences for PCO2, and will invalidate measurements of oxygen status.

B. Arterial versus arterialized blood

Because of the possible risk and discomfort of arterial puncture [16,17], many

attempts to arterialize venous blood by means of warming or vasodilation have been

Figures in brackets indicate the literature references at the end of this paper.

175
Table 1. Differences between arterial and venous blood. ^'

Venous PCO2 P02 O2 Saturation Number and type Year site pH (mmHg) (mmHg) {%) of subjects Reference

1963 Vena Cava 005 +2.4 4 patients with

suspected car- diac defects

2

1974 Central Venous

Central Venous

+ 041

-7.5

501 critically or seriously ill patients on 100 percent O2 328 critically or seriously ill patients on 100 percent O2

3

3

1942 Int. Jug. + 053 -10 31 .7 50 normals 4

1959 Arm + 013 30 resting patients 5

1964 Antecubital + 06 +7.4 30 patients under

anesthes i a

6

1966 Cephal ic + 033 -4.7 +27.3 42-43 patients 7

1959 Hand or Wrist + 051 -8.3 14 patients 8 Hand or Wrist + 01 -1.2 7 patients with warm extremities

8

Hand or Wrist + 002 -1.1 10 patients under

anesthesia

8

1961 Hand or Wrist

Hand or Wrist

044 -3.4

-1.6

14 patients

10 patients under anesthesia

(13 samples)

910

1962 Capi 1 lary , earlobe001 +0.4 31 patients under

anesthesia

11

1964 Capillary

earlobe

Capillary , earlobe

+ 01 -1.0

+9

15 patients under

anesthesia

8 patients

6

12

1962 Capi 1 lary , finger

- 004 +0.2 20 patients under

anesthesia

11

1963 Capillary

finger

- 04 +6 4 patients with

suspected car- diac defects

2

1964 Capillary , finger

+ 03 -7.3 15 patients under

anesthesia

6

1965 Capillary . finger

+ 009 -2.8 13 patients 13

1968 Capillary, +8,5 14 patients 14 thumb

1964 Capillary, .15(max) -48(max) 23 newborns 15 heel very poor very poor correlation correlation

^ All differences are expressed as' arterial -venous

^ All capillary blood was non-arterial i zed

176
attempted. These results are summarized in table 2, Direct comparisons of arterial,

arterial ized earlobe, and arterial ized thumb pulp have been performed for P02 measurements. Good agreement between arterial blood and blood prepared by massaging the earlobe

with nicotinic acid (Trafuril^) was obtained. Blood obtained from the earlobe prepared

by massage without nicotinic acid, the untreated thumb pulp, or the heated hand was found

to differ significantly from arterial blood [14]. Other direct comparisons also found

arterialized blood from the earlobe to be superior to that from the finger [12,27].

Table 2. Differences between arterial and "arterialized" blood.

Year

Sample

site PH PCO2

(mmHg)

P02 O2 Saturation

(mmHg) (%) Number and type

of subjects Referc

1 959 Dorsal hand

vein

+.008 -1.9 14 patients 8

1961 Dorsal hand

vein

+ ,,018 -0.1 14 patients 9

1972 Dorsal hand

vein

+ ,,005 -1.1 16 (P02>70)

c 0

. y (Po2<70)

8

5

patients

and normals

18

1944 Earlobe 0.5 n patients 19

1961 Earlobe -,,002 22 patients 20

Earl obe +0.5 10 patients under

anesthesia (21 samples)

10

1964 Earlobe 3 33 patients 12

L. a [ 1 u uc Cnnri agreement 19 patients 21 Earl obe 0.1 20 pati ents 22

1965 Earlobe 0.3 21 patients 23 Earlobe ,006 + 1.0 0.6 14

2 patients and normals

24

1966 Earlobe r = 0.93 16 patients 25

1967 Earlobe -..001 + 0.3 17 patients (respiratory insufficiency)

26

1968 Earlobe 0.2

(P02, 34-103)

42 patients 27

Earlobe 90 (P02, 222-603)

12 patients 27

Earl obe 0.

(P02, 40-91)

22 patients during

exercise

£.1

1968 Ffl rl nhp 0. 5 14 patients 14

1972 Earlobe ,002 -0.8 0.4 25 normals 28

1973 Earlobe + ,.006 -1.4 0.2 84 patients, pre- operative, opera

tive, and postoperative

29

1961 Ear or finger

+

,003 0.2 0.1 13 patients 30

Ear or finger

,003 8 patients after

one minute

exercise

30

177
Table 2. Differences between arterial and "arterial ized" blood (continued)

Year

Sampl

site pH PCO2

(mmHg)

P02 O2 Saturation

(mmHg) (%) Number and type

of subjects Refere

Ear or finger

-0.2 5 patients after

one minute

exercise

30

1961 Finger 6.4 43 patients 26

1963 Finger -.011 +2.2 47 patients 2

1965 Finger + .034 -0.3 30 patients 13

1966 Finger -.004 0.2 0 20 patients 31

1962 Finger and

heel 1.0 0.2 22 patients 32

1964 Heel + .027 -5.9 20 newborns

0-1 hour

15

+ m -1.0 20 newborns

1-3 hours

15

Heel + .006 -0.7 34 newborns

over 3 hours

15

1967 Heel + .026 -6.2 34 newborns

30 minutes to

24 hours

33

Heel 42 newborns

2-6 f— \j davs \j u y 0 33

Heel 9.5 5 hours to

6 days

33

1969 Heel + .007 -8.3 6.3 45 newborns

3-95 hours

34

1973 Heel r = 0.99 -(10-15 ) r = 0.86 37 npwhnm<i nldpr

than 3.5 days

35

1965 Scalp Smal 1 Smal 1 Smal 1 48 patients 36

^All differences expressed as arterial -"arterial i zed".

As table 2 indicates, in almost all studies of adults the capillary method gave

results equivalent to arterial blood. In newborns, however, differences have been found

and it appears that this technique is not valid in very young children (less than a day)

and questionable even in somewhat older children. The lack of agreement between arterial

and arterial ized capillary blood in newborns has generally been ascribed to poor circulation and warnings about using arterialized capillary blood in any patient with circulatory shock have been presented [12,25,27,37]. Observations by Banister [34], however,

have suggested the presence of edema of the foot, rather than circulatory status, as the

important factor in obtaining valid results from- capillary blood in children.

Prior publications [38,39] have presented conclusions which are in keeping with the

data reviewed here. These conclusions were that arterialized capillary blood is equivalent to arterial blood except for the very young, newborns with respiratory distress

syndrome, and patients in shock. Unfortunately, these categories represent a number of

the patients in whom a valid capillary sample would be highly desirable.

2. Collection Containers

Three types of collection containers for macro pH/blood gas determination have been

suggested; glass syringe, plastic syringe, or vacuum tube. In an early study, no dif178
ference between the pH of whole blood obtained from a vacuum tube and that sampled directly from the patient was noted. In the same study, no differences in plasma CO2

content between sedimented blood from a syringe and centrifuged vacuum tubes was observed.

The O2 saturation of arterial blood collected in a syringe or in a vacuum tube was also

found to be the same [40]. Fleisher and Schwartz [41] noted a poor correlation between

the usual vacuum tubes and glass syringes for arterial blood gas determinations and

designed a nitrogen filled vacuum tube. Comparison of this tube and glass syringes,

showed excellent agreement for PCO2 and P02, and only a +0.01 difference for pH [41].

These nitrogen filled tubes were also tested by Lang, et al. [42], who found them to be

severely contaminated with oxygen. For blood samples, good correlations for pH and PCO2 were observed, but significantly higher P02 values were found with the vacuum tube. Further

work by these authors showed the unsuitability of the presently available vacuum tubes for

measurement of'Po2 and oxygen saturation [43].

A. Glass versus plastic

Saline equilibrated to a high P02 (^^700 mm Hg) showed no decline in measured P02

when kept in glass syringes, but showed a decline of 21.6 percent after 3 hours and 67

percent after 48 hours when stored in plastic syringes. Blood tonometered to high P02

values ('^^600 mm Hg) also showed a greater decay in plastic syringes as compared to glass.

The differences in decay rate were apparent in fifteen minutes [14]. The diffusion of

oxygen and carbon dioxide from plastic syringes has been studied. The percent change in

gas content per hour varied considerably (over 100 percent) between the two types of

plastic syringes tested. For the better syringe type, the changes at room temperature (in

percent change of gas content/h) were 0.034 and 0.082 for O2 and CO2, respectively. When

the syringes were kept on ice these changes were 0.128 and 0.154. These authors expected that

the changes would be considerably slower when the syringe content was liquid and did not

feel that the decay would be clinically significant [44]. Measuring the P02 of deoxygenated water gave higher results for samples in plastic syringes than for glass (37 mm Hg

compared to 13 mm Hg) [42]. Laver and Seifen [45] have noted that plastic syringes maintain

a high P02 as well as glass syringes for up to 2 hours, but recommend glass syringes if

delay in analysis is expected.

Scott, et al. [46], found that when water tonometered to 96 percent O2 was sampled in

6 different types of plastic syringes, large decreases were found when the results were

compared to the same sample in glass syringes. Similarly, large increases were found when

sampling water at 0 percent O2. Syringes made of polystyrene showed considerably greater

changes than those made of polypropylene or S.A.N, co-polymer. Tonometered blood samples

showed the following changes (in mm Hg) after 11.5 minutes, -68, 0.4, and 0.0 at an initial

P02 of 680, 102, and 68, respectively. These authors presented evidence that the loss or gain in O2 was due to diffusion into the walls of the syringe rather than diffusion through

the syringe. These authors estimated the magnitude of errors expected under various

conditions. This data is presented in table 3, and extimates the influence of initial P02,

temperature of storage, pH, hemoglobin and barometric pressure on the error introduced by

the use of plastic syringes. It is evident that the error introduced will depend on a number of factors and the acceptibil ity of plastic syringes will depend on the application.

The above study also noted that interchangeable glass syringes were not as effective as glass syringes with matched barrel and plunger.

The above data indicates that under most conditions, the interchange of glass or plastic syringes will not affect the clinical utility of the oxygen measurement [44-46].

The excellent work of Scott (table 3), however, does show that in some circumstances

inaccurate results will be obtained when plastic syringes are utilized. An example of

this situation may be the severely anemic patient with alkalosis who is on a respirator.

There are advantages to the laboratory in using plastic syringes such as cost, safety, and

seal reliability. When mixing the blood, the plunger of more than one glass syringe has

fallen out if care is not taken. The area of collection containers appears to be in need

of some fresh approaches which would retain the accuracy of glass syringes but with the

convenience of plastic ones. Regardless of whether glass or plastic syringes are used,

the addition of mercury to facilitate mixing the blood is not recommended due to the

danger of peripheral embolism [47].

179
Table 3. Computed oxygen exchange of blood stored in plastic syringes [46].

Change of tension

Storage

time

(min)

Ini tial tension

(mmHg) (mmHg) (% initial

Change

of content

{7o initial)

"Normal"

condi tions

Effect of storage time

Effect of ambient

temperature during

storage (Ta)

Effect of Hb

30-60

Effect of pH

Effect of barometric

pressure- (Pb)

30-60

30-60

30-60

30-60

650 -106 3 -16 4 -1 4 300 -21 3 -7 1 -0 3 100 5 8 5 8 0 4 33 0 3 0 9 0 9

650 -30 3 -4 7 -0 4 Ta 650 37 -129. 1 -19. 9 -1. 7 650 4 -83. 3 -12. 8 -1. 1 100 37 3. 3 3. 3 0. 2 100 4 7. 6 7. 6 0. 5

650

650

100

100

300

100

Hb

15

5 15

5 pH7.7 6.8

Pb 780

710

600

-106.3

-107.5 5.8

12.7

-24.3 -10.8

6.1 5.2 3.7

-16.4

-16.5 5.8 12.7

-8.1 -3.6

-1.4

-3.6 0.4

1 .1

-0.4

-0.2

0.4

0.3

0.2

Notes: Syringes: 5 ml or 2 ml , initially in equilibrium with ambient air. Exchange factor: 8 x 10''* (ml/lOO m)/mmHg for 30-60 minutes storage and

2 X lo"** for 2 minutes.

Normal conditions: Hb = 15 g/lOO m, pH = 7.4, Pb = 760 mmHg, initial blood temperature = 37 °C, ambient temperature = 22.5 °C. Reprinted with the permission of Brit. Med. J.

3. Effect of Anticoagulants

While serum and plasma give comparable results for pH [40] and CO2 content [48], ^ whole blood is necessary for oxygen analysis. The use of citrate [40,49], oxalate (NH^

and K"*") [40], or EDTA [7,40] as anticoagulants has been found to cause low results for pH.

The use of potassium oxalate alone as anticoagulant caused an increase in pH [40,50].

This increase in pH was shown to be related to chelation of calcium [50]. Heparin has

almost no effect on pH (1 mg/ml = -0.003). This small effect was the same in lipemic

blood but no data concerning the storage of such specimens was presented [51]. In an effort to inhibit glycolysis, heparin-fluoride mixtures have been utilized. In one study,

no effect of fluoride on oxygen consumption was found [52] while in others a decrease in

decay of P02 of 1/4 to 1/3 was observed [53,54]. Paradoxically, fluoride completely eliminated the rise in PCO2 with storage even though the P02 was only partially affected [53]. No differences in the pH of four blood specimens were found when heparin-fluoride was utilized [49]. The addition of fluoride in amounts sufficient to inhibit glycolysis

lowers the pK' of blood by 0.01 units due to a salt, effect [55]. Fluoride (5 mg/ml)

caused an increase in pH of 0.03 and a fall in PCO2 of 6 mm Hg. This effect was found to be concentration dependent [51]. Upon storage, blood containing fluoride reacted differently

than blood without fluoride. When stored at 38 °C, the change in the pH of blood reached a minimum at 1/2-1 hour (0.01-0.02) and a maximum at 3-4 hours (0.17-0.19). This phenomena

was concentration dependent, delayed at 22 °C, and was not due to leukocytes [51]. The

addition of fluoride causes a rise in potassium which invalidates its use for potassium

measurement [56]. In our laboratory, potassium analysis is being requested on the same sample of blood as pH/blood gas with greater frequency. The rise in potassium and the above

data on changes in pH during storage suggest that fluoride not be used as an additive to samples for pH/blood gas measurement. It is therefore recommended that heparin be used as the sole anticoagulant for pH/blood gas analysis.

180
4. Stability In Vitro

In this section the stability of pH, PCO2, CO2 content, and P02 in blood under varying

conditions is reviewed. Since anticoagulants other than heparin can have a deleterious

effect on the measured values of these quantities, only data obtained with this anticoagulant will be considered. As will be discussed later, the temperature of measurement can affect results and therefore some early data where the temperature control is suspect [57]

will not be considered.

A. pH

The average decay rates for pH (-pH/h) found in various studies are shown in table 4. This decay rate has been shown to be independent of hemoglobin concentration but very

dependent on the number of leukocytes [51]. It is evident that the decay in pH is quite

temperature dependent and this dependence has been shown to follow the Arrhenius relationship [62]. Exposure to air by dropping 50 yl of blood onto a glass slide resulted in a rise of only 0.01 pH after 2 minutes [51]. The small decay rate for pH should not cause problems in routine handling of blood specimens if they are placed in ice water after

being obtained. Such treatment has been stated to result in decay rates of no more than

-0.008 pH/h even in patients with leukemia [63],

B. Pcoa

As shown in table 4, PCO2 in blood has been found to increase during storage. For

blood stored at 4 °C the average increase is less than 0.7 mm Hg/h. The magnitude of this

increase has been found to be related to the number of white blood cells present but not

to the number of reticulocytes [53]. This small increase will not usually be clinically

Table 4. Stability of pH, PCO2, CO2 content and P02.

Storage Rate of Number of Parameter Year temperature change experiments Comments Reference

Data not given 55

Linear for 3 hrs and 51 dependent on leuko- cytes but not hemo- globin

Linear for 3 hrs and dependent on leuko- cytes but not hemo- globin

Linear for 3 hrs and dependent on leuko- cytes but not hemo- globin

Based on 4 hours 13

Linear for 3 hours 51

52

13

1956 0. 060 (-pH/h) ?

1961 38 °C 0 062 (-pH/h) 10

22-24 °C 0 024 (-pH/h) 5

0-4 °C 0 006 (-pH/h) 4

1965 2-4 °c 0 006 (-pH/h) 17

1961 38 °c 4 8 (mmHg/h) 10 22-24 °c 2 5 (mmHg/h) 5 0-4 °c 0 6 (mmHg/h) 4

1965 37 °c 6 6 (mmHg/h) ?

4 °c 0 7 (mmHg/h) ? 2-4 °c . 0 2 (mmHg/h) 5

181
Table 4. Stability of pH, Pcoa, CO2 content and P02 (continued)

Storage Rate of Number of Parameter Year temperature change experiments Comments Reference

PO;

1959 room? 0 ? No effect after 2 hours 40

1971 room? 0 1 No effect after 2 hours

58

1961 37 °C 175 (-mmHg/h) 5 Initial P02 , over 400 52

20 (-mmHg/h) 5 Initial P02 , 'v, 100

1965 37 °C 3. 4 (-mmHg/h) ? Initial P02 , 51-93 59 22-24 °C 1. 8 (-mmHg/h) ? Rate was not constant,

and was calculated

from 1 hr data

1 °C 1

.

7 (-mmHg/h) ? Rate was not constant,

and was calculated

from 1 hr data

37 °C 121 (-mmHg/h) ? Initial P02, 360-670,

rate was not constant

and was calculated from

1 hr data 22-24 °C 16 (-mmHg/h) ? Initial P02, 360-670,

rate was not constant

and was calculated

from 1 hr data

1 °C 1

.

7 (-mmHg/h) 7 Initial P02, 360-670,

rate was not constant

and was calculated

from 1 hr data

37 °C 156 (-mmHg/h) 55 High initial P02 53

4 °C 30 (-mmHg/h) 10 High initial P02

07 op 24 ( -mmHn / h ^ 5 Initial P02, 60-100 53

37 °C 48 (-mmHg/h) 16 Initial P02, 125-150

37 °C 162 (-mmHg/h) 55 Initial P02, 440-610

1966 4 °C 10 (-mmHg/h) 7 31 room 450 (-mmHg/h) 89 samples

5 normals Breathing 100% O2. 60 Rate is projected based

on first 10 minutes

1968 37 °C 180 (-mmHg/h) 7 Initial P02, 680 14

24 °C 120 (-mmHg/h) 7 Initial P02 , 680

4 °C 30 (-mmHg/h) 7 Initial P02 , 680

37 °C 50 (-mmHg/h) 7 Initial P02, 143

24 °C 20 ( -mmHg/h) 7 Initial P02 , 143

4 °C 10 (-mmHg/h) 7 Initial P02, 143

7 180 (-mmHg/h) 7 Higher if leukocytosis present

61

1968 0-4 °C -0 5 (-mmHg/h) 10 P02, 52-92 27

significant but extensive information on patients with high levels of white blood cells,

e.g., leukemia, is not available.

C. CO2 content

As expected from the data concerning pH and PCO2, little or no change in blood CO2

content has been found when the blood is kept unopened to air (table 4). The temperature

182
of centrif ugation has been suggested as a possible small cause of error in CO2 content

measurements because of the different temperature coefficients of plasma and whole blood

[64]. In two studies, however, no effect of the temperature of centrifugation (5 °C or 37.5 °C) was found [40,55]. Since the temperature of a centrifuge operated in an ambient

temperature room is over 30 °C, this potential error source can be ignored.

A more significant stability problem occurs when plasma or serum is placed in the

plastic sampling cups which are so commonly used. The loss of CO2 in sample cups open to the air has been shown to average 3.5 meq/1 for 15 samples after 1 hour. In one patient

with a PCO2 of 75 mm Hg, a 7 meq/1 loss after 3 hours occurred [65]. Four techniques to minimize CO2 loss to the atmosphere have been described. The oldest of these techniques

is the use of oil to eliminate contact of the sample with the atmosphere. This technique

has been criticized for two reasons: (1) since blood collected with and without oil gives

the same results [48], its use during collection is superfluous; and (2) CO2 is quite soluble in oil (see annotated bibliography in reference 66) and significant absorption of CO2 occurs if the sample is mixed with the oil. Such mixing can occur during collection

and/or centrifugation [67]. While covering the sample cups with oil does retard CO2 loss

to the atmosphere [68], the potential CO2 loss to the oil makes its use unadvisable.

The second method consists of placing a plastic disk in the sample cup so that the

sample is isolated from air. The sample probe then offsets the disk at the time of analysis.

This system has resulted in no loss of CO2 in 3 hours [69]. The third method is the use of added alkali (1 drop of 1 N_ NHi+OH) to decrease the PCO2. This technique was found to stabilize the CO2 content for up to 4 hours and was felt to be superior to the plastic

disk method [65]. The fourth technique is the use of a metal or plastic plate to cover

all the cups in a sample tray. Such equipment is commercially available but the author is unaware of any published studies on its effectiveness in retarding CO2 loss.

The loss of CO2 to the atmosphere during handling is certainly capable of affecting

the clinical interpretation of CO2 content measurements. Of the methods noted above for

minimizing this phenomena, the alkali method of Gambino [65] appears to be best but further

validation by other workers would be helpful.

D. P02

The stability of P02 in shed blood under various conditions is shown in table 4. It

is evident that the stability of P02 is quite temperature dependent [14,25,53,54,59,70,71]

and rapid chilling of all blood samples is strongly suggested. It is also evident that

the rate of decay of P02 is highly dependent on the initial P02 [14,25,52, 53,59,71,72].

This has been attributed to the shape of the oxygen dissociation curve [53]. The O2 decay

rate in shed blood is not affected by total hemolysis of the blood sample [53].

The source of the oxygen consumption in blood specimens has been the subject of

study. In one study the oxygen consumption of normal mature erythrocytes was found to be

negligible but blood with abnormal numbers of reticulocytes had an oxygen consumption

proportional to the percentage of reticulocytes present [73]. This increase in oxygen

consumption with increase in percentage of reticulocytes has been confirmed by others

[53,74]. The rate of oxygen consumption has also been related to the white cells. Figure

1 shows the data of Lenfant and Aucutt [53] relating oxygen consumption to the white blood

count. Normal and abnormal white cells are included in this data but no comparison of

their oxygen consumption was made. Higher oxygen consumption in blood from leukemic

subjects has been noted by another worker as well [61]. Because of the dependency of O2

consumption on the cellular components of the blood, correction factors for storage time

are not recommended and prompt analysis of chilled specimens appear the best way of

handling blood for accurate P02 measurements.

Exposure of blood to air for short periods (10-30 seconds) results in significant

changes in P02 toward the P02 of room air [25]. The presence of air bubbles (up to 10

percent of the total syringe volume) has little effect on P02 until the surface area is

increased by mixing [71]. The presence of air bubbles in blood samples has been blamed

for discrepancies in interlaboratory surveys for P02 [75]. Since whole blood is mixed

prior to pH/blood gas analysis to avoid suspension of the cells, great care should be

exerted to exclude air bubbles at the time the sample is obtained.

183
30

25

20 •

15

10"

5

-

PO2 decay, mm Hg/min

\ •• ••. • • • • White Blood Cells Count

per mm^

10,000 20,000 50,000 100,000 200,000 500,000 1,000,000

Figure 1. P02 decay as a function of the white cell count. Black circles are for samples

containing NaF. Arrows indicate the P02 decay of samples with normal white cell counts

without f-^) and with (—>) NaF. (From Lenfant and Aucutt [53] with permission of

J. Appl. Physiol. )

5. Effect of Measurement Temperature

pH, P02, and PCO2 values are all dependent on the measurement temperature. Since

virtually all commercial pH/blood gas equipment measures these quantities at 37 °C, an error is introduced when the temperature of the patient at the time the sample is obtained

deviates from this value. The work of various authors has been reviewed and a simplified

table derived for temperature correction of pH, P02, and PCO2 [76], This table is shown

as table 5 and is recommended as a guide to temperature effects on pH and blood gas measurements.

6. Effects of Red Cells on pH Measurement

No effects of hematocrit on the measurement of P02 or PCO2 have been reported. However,

differences in measured pH between plasma and whole blood have been observed (plasma-whole

blood =0.01) [55,77]. The relationship between pH and hematocrit (0-100) in 5 blood

samples has been elucidated and the cause of the small pH discrepancies was felt to be due

to differences in residual liquid junction potentials between plasma and whole blood [77].

The magnitude of this error is quite small and except in highly unusual circumstances will

not be clinically significant.

7. Effects of Dilution

A 12-13 percent dilution of blood with physiological saline produced no significant

change in pH but resulted in a fall in PCO2 of 16 percent and a fall in plasma bicarbonate

of 15 percent [51]. Another study showed an increase in pH of 0.023 pH unit after a 20

percent dilution with physiological saline. The addition of calcium (5 meq/1) to the

saline eliminated this dilution error [78]. A 33 percent dilution of blood with isotonic

saline caused no change in pH, a decrease of 3.8 mm Hg (4 percent) in P02 and a decrease of

9.2 mm Hg (24 percent) in PCO2 [24]. The above data indicate that pH and P62 exhibit little

184
Table 5. Temperature- correction factors for blood pH and gas measurements [76].^

atient's temperature pH ^C02

{%) (%)

°F °C (Add to observed values)

110 43 -.09 +22 +33

109 42.5 -.08 +21 +32

108 42 -.07 +19 +30

107 41 .5 -.07 +17 +27

106 41 -.06 +16 +25

105 40. 5 -.05 +14 +22

104 40 -.04 +12 +19

103 39.5 -.04 +10 +16

102 39 -.03 +8 +13

101 38. 5 -.02 +6 +10

100 38 -.01 +4 +7 98-99 37 None None None

97 36 + .01 -4 -7

96 35. 5 + .02 -6 -10

95 35 + .03 -8 -13

94 34. 5 + .04 -10 -16

93 34 + .04 -12 -19

91 33 + .06 -16 -25

90 32 + .07 -19 -30 88 31 + .09 -22 -35

86 30 + .10 -26 -39

84 29 + .12 -29 -43

82 28 + .13 -32 -47

81 27 + .15 -34 -51

79 26 + .16 -37 -54

77 25 + .18 -40 -57

75 24 + .19 -43 -60

73 23 + .21 -4S -63

72 22 + .22 -48 -65

70 21 + .24 -50 -67

68 20 + .25 -53 -70

^Reprinted with permission of Cliniaal Chemietyy.

change, even with high dilution, but PCO2 and CO2 content change in proportion to the degree

of dilution. Unless a large volume of anticoagulant is used, this should not introduce a significant error. In sampling from indwelling catheters, care should be taken to remove

the fluid filling the catheter before obtaining the blood specimen if accurate assessment of

CO2 status is desired.

8. Effect of Stasis

No difference in pH between venous blood obtained without a tourniquet and blood

obtained with the tourniquet in place for 2 minutes was found in 5 subjects. It was emphasized that the tourniquet should be left in place since the stasis produced is at the

capillary level and is not reflected in the vein until the tourniquet is released [79].

These results have been confirmed and it was noted that the tourniquet can be left in

place for 5 minutes without exercise (hand pumping) but that after tourniquet release the

pH decreased precipitously [5].

185
9. Miscellaneous Effects

Arterial pH was found to decrease by over 0.2 pH units after exhausive exercise and

took over 30 minutes to recover. Pcoa decreased by approximately 5 mm Hg as well [80].

Continuous monitoring of blood pH has shown small fluctuations with each group of respirations and has shown that hyperventilation results in a large increase in pH [81]. Warnings

concerning increases in pH and decreases in PCO2 during hyperventilation due to the pain or anxiety of arterial puncture have been presented [82]. PCO2 values obtained in the

sitting or standing position are 3-4 mm Hg lower than those obtained while recumbent [83].

The above data suggest that pH/blood gas studies should be performed on subjects known to be at rest unless special investigations are being performed. The injestion of food has

been shown to cause an increase in arterial pH of 0.015 to 0.03 with no change in PCO2.

The PCO2 was found to rise with sleep (5 mm Hg maximum) and the pH to decrease [84].

These changes associated with sleep and meals are small but should be thought of when it

is decided when to obtain a sample for pH and blood gas analysis. Capnography (recording

of expiratory COa-curve) did not alter the measured values for P02 or PCO2 in patients

with chronic lung diseases [85]. Halothane caused a slow response of some P02 electrodes

and could cause erroneous P02 results with such electrodes [86]. Further work on this

problem with more modern P02 electrodes has not yet been reported.

10. Conclusion

Advances in instrumentation have resulted in reliable, commercially available systems

for the analysis of pH and blood gases. The future will probably see even more reliable

instruments and these determinations will become more commonly performed than they are today. As the analytical component of laboratory error decreases it will be even more important for the clinician and laboratorian to appreciate the myriad of other factors

which can lead to "laboratory error". In reviewing this information, it was obvious that

for proper interpretation of laboratory data more knowledge needs to be obtained regarding

these factors. Factors such as composition and design of collection devices, changes in values when storing blood with abnormal cells, and laboratory sample handling for CO2

content require further careful investigation.

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Lang, G. E., Mueller, R. G., and Hunt, P. K. , Possible error resulting from use of

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water samples stored in plastic and glass syringes, Brit. Med. J. 3_, 512 (1971).

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Lenfant, C. and Aucutt, C. , Oxygen uptake and change in carbon dioxide tension in

human blood stored at 37 °C, J. Appl. Physiol. 20, 503 (1965).

Mellemgaard, K. , Disappearance of oxygen from shed blood at various temperatures,

Saand. J. Clin. Lab. Invest. 18, 380 (1966).

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determinations, J. Appl. Physiol. 9^, 189 (1956).

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versus heparinized plasma for eighteen common chemistry tests. Is serum the appropriate specimen? AmeT. J. Clin. Path. 62, 545 (1974).

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Med. 22, 1115 (1937).

Spring, B. L., Simultaneous determination of carbon dioxide and chloride by the

autoanalyzer, Amer. J. Med. Teoh. 37_, 361 (1971).

Eldridge, F. and Fretwell, L. K. , Change in oxygen tension of shed blood at various

temperatures, J. Appl. Physiol. 20^, 790 (1965).

Fletcher, G. and Barber, J. L., Effect of sampling technique on the determination of

Pa^^ during oxygen breathing, J. Appl. Physiol. 21, 463 (1966).

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Soi. 148, 115 (1968).

Kelman, G. R. and Nunn, J. F., Nomograms for correction of blood P02, PCO2, pH, and

base excess for time and temperature, J. Appl. Physiol. 21, 1484 (1966).

Gambino, S. R. , Astrup, P., Bates, R. G., Campbell, E. J. M. , Chinard, F, P., Nahas,

G. G. , Siggaard-Andersen, 0., and Winters, R. , Report of the Ad Hoc Committee on acidbase methodology, Amer. J. Clin. Path. 46^, 376 (1966).

Astrup, P. and Siggaard-Andersen, 0., Micromethods for measuring acid-base values of

blood. Adv. Clin. Chem. 6^, 1 (1963).

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(alkalinization) for preventing loss of CO2 from open cups, Amer. J. Clin. Path. 45,

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Chez le nourrison, Clin. Chim. Acta 2_, 442 (1957).

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189
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Ishikawa, S., Fornier, A., Borst, C, and Segal, M. S., The effects of air bubbles and

time delay on blood gas analysis, Ann. Allergy 33^, 72 (1974).

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Clin. Med. 69, 853 (1967).

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745 (1919).

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Runck, A. H., Proper technic in measurement of blood gases, n.e.j.m. 291_, 476 (1974).

Burnett, R. W. and Noonan, D. C, Calculations and correction factors used in determination of blood pH and blood gases, Clin. Chem. 20, 1499 (1974).

Maas, A. H. J., pH determination of body fluids with micro glass electrode and a saturated KCl bridge in the cell, CUn. Chim. Acta 28, 373 (1970).

Ware, A. G. , Nowack, J., and Westover, L., Capillary blood pH by a dilution technic,

Clin. Chem. 9, 340 (1963).

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and after exhaustive exercise, J. Appl. Physiol. 34^, 677 (1973).

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blood gases: A concise guide for clinicians, South. Med. J. 66, 1051 (1973).

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and Wilkins Co., Baltimore, 1964).

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Saand. J. Clin. Lab. Invest. 22, 151 (1968).

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Soand. J. Clin. Lab. Invest. 34, 43 (1974).

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electrode errors due to polarographic reduction of halothane, J. Appl. Physiol. 31_, 640 (1971).

190
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

INSTRUMENT SPECIFICATIONS

S. Raymond Gambino

Columbia Presbyterian Medical Center

New York, New York 10032, USA

Before I give you my specifications for laboratory instruments I want to show you what

has happened to the number of blood gas assays performed at our hospital from 1970 through

1974 (table 1).

Table 1. Blood gas tallies.

Central Peripheral

Day Evening Night

1970 5,639 9,824

1971 7,034 13,593

1972 8,464 11,794 3,083

1973 7,505 8,812 4,756

1974 8,292 9,732 5,053 13,356

The data in table 1 show that in 1970, we performed a little over 15,000 blood gas

analyses whereas in 1974 we performed 36,433 such assays. More important than the short

doubling rate is the distribution of the assays. They do not concentrate on the day shift

or on weekdays. Instead, the load is rather uniform throughout a 24-hour period, 7 days a week. In addition, table 1 shows preliminary data for assays performed outside of the main

laboratory. In 1974, we had 2 peripheral blood gas machines located in intensive care units.

Recorded assays performed on these 2 instruments by non-laboratory staff came to 13,356 in

1974. However, the actual number of such assays was probably more than 3 times that number

since most emergency assays are not logged. In 1975, with the availability of instruments

which incorporate automatic washout. Dr. Philip Altman and I have expanded peripheral cov- erage on the adult service to four locations utilizing Instrumentation Laboratory's 513

series

Given the above data, and the obvious need for "bedside" measurements, the specifications for a practical clinical blood gas instrument become more obvious. Such an instrument

should:

- Have automatic calibration.

- Have one simple sampling step.

- Be self cleaning.

- Require as small a sample size as possible, preferably less than 200 yl

- Be as small as possible in physical size.

- Consume as little gas and reagent as possible, yet maintain constant calibration and

be ready for instant use at any time.

- Be operable by non- laboratory personnel, including non-technical personnel such as clerks and nurses' aides.

191
I have outlined the functional characteristics of my ideal instrument. What about the

required accuracy? What are the maximum error limits tolerable in a peripheral blood gas unit serving critically ill patients? The following "rough cut" figures (table 2) are based

on my observations of what doctors actually do with the data provided.

Table 2. Tolerable error limits (± 2 standard deviations).

PH + 0.02

PCOg + 3 mmHg at 40 mmHg

POg + 5 mmHg at 100 mmHg

+ 3 mmHg at 60 mmHg

Base + 2 mM/1

Do current instruments meet these minimum standards for accuracy and precision? Do current operators of instruments meet these standards? The following tables show that there

is more lab-to-lab variation in precision, utilizing the same instrument-model, than there

is between models from different manufacturers.

Table 3 shows data for pH. Note the striking differences in precision for the same instrument in different laboratories. These data were obtained by Dr. M. Miller of General

Diagnostics utilizing a new ampouled blood gas control which I describe in greater detail in my other paper at this symposium.

Table 3. pH triplicates on 16 days, different labs different machines.

instrument N Mean %RSD

165 a 48 7.366 5.1

165 b 48 7.389 2.0

BMS 3 48 7.386 2.8

ABLl a 48 7.374 5.5

ABLl b 48 7.336 1 .3

313 a 48 7.361 8.5

313 b 15 7.380 6.0

513 48 7.374 2.2

Labs ' mean 7.376

Manufacturer 7.371 1.4

(RSD calculated on H ion cone. not pH)

192
Table 4 shows the data for PCO2. We see again the striking differences in precision

(expressed as %RSD or percent relative standard deviation, i.e. S.D. expressed as a percent

of the mean)

Table 4. PCO2 triplicates on 16 days, different labs different machines.

nstrument N Mean %RSD

165 a 48 37 6 7.4

165 b 48 39 5 2.8

BMS 3 48 33 8 6.4

ABLl a 48 35 2 11.5

ABLl b 48 43 8 3.5

313 a 48 39 0 13.2

3.3 b 15 40 1 1.0

513 48 41 3 3.2

Labs ' mean 38 6

Manufacturer 39 9 1 .8

Table 5 shows data for P02. The same story is again evident. There is greater variation between labs utilizing the same model of an instrument than there is between different

models made by different manufacturers.

Table 5. P02 triplicates on 16 days, different labs different machines.

Instrument N Mean %RSD

165 a 48 120.4 11.3

165 b 48 102.6 3.0

BMS 3 48 109.3 3.2

ABLl a 48 . 126.8 14.9

ABLl b 12 120.8 4.4

313 a 48 109.9 7.9

313 b 15 102.8 1.6

513 48 100.0 1 .6

Labs ' mean 107.5

Manufacturer 104.9 1.4

These data indicate to me that our problem is more a problem of quality control at individual sites than it is a problem with manufacturers, or with primary standards for pH

and gases.

193

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

EFFECTS OF THE LIQUID JUNCTION ON pH MEASUREMENT IN BLOOD;

THE 0.160 MOL/L SODIUM CHLORIDE BRIDGE

A. H. J. Maas

Department of Cardiology and Thoracic Surgery

University Hospital

Cathari jnesingel 101 Utrecht, The Netherlands

1 . Saturated KC1 Bridge

The practical pH-cell is built up of an electrode sensitive to hydrogen ions (usually

the glass electrode) and a reference electrode independent of the test solution (usually a calomel electrode) with a salt bridge of saturated KC1 , which means that buffers and test

solutions are applied in cells with a diffusion potential. Consequently the pH of a solution

X is defined operationally by the relationship:

pH(X) = pH(S) + ^ ^

(1) RT £n 10

In most cases the liquid junction potential during measurement of the standard solutions

is not equal to that existing during measurement of the test sample, that is, the residual

liquid junction potential will be unequal to zero, with the result that the measured pH of

the sample will hardly ever be exactly on the conventional pH(S) scale of the National

Bureau of Standards (NBS). The degree of error can be calculated semiquantitatively from

the molar conductivity and the concentration of the ions using Henderson's equation for the

liquid junction potential between two solutions of different composition. With this formula

we calculated some diffusion potentials of the junction: test soln. || satd. KCl at 25 °C [1]^.

Table 1. Calculated liquid junction potentials and residual liquid junction

potentials of junctions with the saturated KCl bridge (see text).

Liquid junction Residual liquid

potential junction potential

No. Solution X I mV pH

units

diffe- rencemV pH

units

1. equimolal phosphate buffer 0.100 2.1 0.035 1-2 0.1 0.002

2. blood phosphate buffer 0.100 2.0 0.033 2-3 0.9 0.015

3. 0.160 mol/1 NaCl 0.160 1.1 0.018 2-4 0.8 0.013

4. 0.025 mol/1 NaHCOg + 0.135 mol/1 NaCl 0.160 1.2 0.020 2-5 0.8 0.013

5. 1 iquor cerebrospinal is (CSF) 0.160 1.2 0.020 2-6 0.5 0.008

6. plasma 0.160 1.5 0.025

^Figures in brackets indicate the literature references at the end of this paper.

j 195

1
This table shows that:

a. A saturated KCl bridge provides junction potentials that are small: those of the

phosphate buffers of NBS are nearly equal as well as those of spinal fluid (CSF) and plasma.

b. The residual liquid junction potential in the pH determination of CSF and plasma

after calibration with the phosphate buffers is about 0.01 pH unit.

Various factors influence the liquid junction potential [1,2]:

a. Composition . Saturated KCl solution is mostly used as the bridge solution. Attention ought to be paid to dilution of the bridge solution, which gives lower results. So the

pH of plasma is measured 0.003 pH lower with a 3.5 mol/1 KCl bridge as recommended by Bates

[3]. With a 0.15 mol/1 KCl or 0.16 mol/1 NaCl bridge, as we will discuss further on, values more than 0.1 pH lower than those using a saturated KCl bridge in the cell are obtained.

Therefore, it is important to recommend the pH determination with the saturated KCl bridge as

a reference method.

b. Structure . As demonstrated with the micro glass capillary electrode for over

fifteen years, a stable and reproducible liquid junction is obtained when the test solution

in a small plastic tube is dipped into the saturated KCl bridge solution, -i.e., the high

density KCl solution is below the test solution. The opposite arrangement is unstable as the KCl tends to flow down. Allowance has to be made for this too when permeable membranes,

separating the test solution from the KCl solution, are used at the boundary. As a conse- quence to such a bad construction, it is necessary to refill the reference electrode in the

Corning M 165 blood gas instrument every day or two.

c. Temperature . We determined the change in the residual liquid junction potential

with the temperature from the pH difference: pH (ref. 25 °C) - pH (ref. 38 °C) where pH

(ref. 25 °C) and pH (ref. 38 °C) mean the values obtained with the same micro glass electrode

at a constant temperature of 38 °C and two saturated calomel electrodes with a salt bridge,

each at the indicated temperature. Both cells were calibrated with the equimolal phosphate

buffer (pH = 6.840 at 38 °C). For blood, plasma, cerebrospinal fluid and a 0.025 mol/1

sodium bicarbonate + 0.135 mol/1 sodium chloride solution a difference of 0.011 pH was

found; so, we calculate for the temperature coefficient of the residual liquid junction

potential

:

ApH/At = 0.011/13 = 0.00085 pH/°C.

The fact that this coefficient is equal for the mentioned solutions indicates that differences in activity and mobility of the potassium and chloride ions at 25 °C and at 38 °C are the most important factors. To reduce the influence of this thermal diffusion on the pH

readings, the whole pH cell should be thermostated at one temperature, i.e., 37 °C.

d. Effect of Blood Cells . To be able to understand in which order blood cells influence the diffusion potential of the salt bridge, we determined directly the liquid junction

potential differences:

E. (plasma || satd. KCl) -E. (plasma with cells || satd. KCl),

which effect will from now on be called "cell effect."

We used the following cell:

Cal I, satd. KCl || plasma | suspension of cells in plasma || satd. KCl, Cal II

E . E . E

.

J.P J,pc j,c

If we put ^. equal to zero, the emf difference will give the cell effect:

j,pc ^

^

E = E T TT + E. - E. - E T , (E. E 0). cal II J ,c J ,p cal I ^ J ,pc

196
We determined this effect at 25 °C and 38 °C. Figure 1 shows the cell effect in pH units of

10 experiments plotted as a function of the hematocrit.

0.005+

0.000

0.005-

0 o" > 0

0.005+

0.000

0 20 40 60 80

0.005-

100

Figure 1. The cell effect at, respectively, the junction with saturated KCl (la and lb)

and with 0.160 NaCl (Ila and lib) as a function of the hematocrit. The different

symbols indicate different blood samples.

Graphs la and lb show the cell effect increasing with the cell concentration. At lower

pH, the cell effect is greater than at higher pH. At pH = 7.4 and normal hematocrit value

(45%), the cell effect is 0.008 pH ± 0.002 (38 °C) and 0.012 pH + 0.002 (25 "C). To be able

to establish whether the cell effect could be due to the negative charge of the erythrocytes

(a so-called suspension effect), we determined in an analogous way the difference in diffusion potential of the bridge with an isotonic solution.

Graphs Ila and lib show the cell effect of the 0.160 mol/1 NaCl bridge plotted as a function of the hematocrit. We observe that even with hematocrit values of 90 percent, no

cell effect exists. From the experiments we may conclude that the cell effect of the

saturated KCl bridge is not caused by intact blood cells but most likely due to precipitation

of proteins.

2. 0.160 mol/1 NaCl Bridge

In order to avoid this disturbing effect of blood cells on the diffusion potential at the saturated KCl junction, we investigated whether the saturated KCl calomel electrode

could be substituted by a calomel electrode with an isotonic salt bridge, notably the 0.160

mol/1 NaCl bridge [4]. To that end, we made comparative pH measurements with the glass

electrode-saturated KCl calomel electrode cell and the glass electrode - 0.160 mol/1 NaCl

calomel electrode cell; we calibrated with equimolal phosphate buffer.

197
The differences pH (satd. KCl ) - pH (0.160 mol/1 NaCl ) were as follows:

Table 2. pH difference of values obtained with the glass

electrode - saturated KCl calomel electrode cell and the glass

electrode - 0.160 mol/1 NaCl calomel electrode cell.

Solution X pH range pH (satd. KCl) - pH

(0.160 mol/1 NaCl)

Blood phosphate buffer 7.391 + 0.005

0.025 mol/1 NaHCOs + 0.135 mol/1 NaCl 7.267 + 0.153

Cerebrospinal fluid ("pooled") 7.486 + 0.163

Plasma 7.453-7.481 + 0.125 ± 0.005

Blood 7.311-7.370 + 0.106 ± 0.005

This table shows that the pH of the blood phosphate buffer is influenced little. On the

other hand, the pH values of body fluids measured with a cell with a 0.160 mol/1 NaCl

bridge, dependent on the test solution, are more than 0.1 pH lower than those determined

using a cell with a saturated KCl bridge. Since the glass electrode, calomel electrode and

standardization for both cells were the same, the pH differences must be interpreted as differences in residual diffusion potential.

This result will be more clear when we compare some diffusion potentials of the junction: solution X I 0.160 mol/1 NaCl calculated with Henderson's formula. These values are given in the following table 3:

Table 3. Calculated and measured residual liquid junction

potentials of junctions with 0.160 mol/1 NaCl bridge

Liquid junction Residual liquidpotential junction potential

calculated measured

No. Solution X I mV pH

units

diffe- rencemV pH

units

pH

units

1 equimolal phosphage buffer 0.100 8.1 0.137 1-2 1.1 0.019 0.005

2 blood phosphate buffer 0.100 7.0 0.118 1-3 7.1 0.120 0.153

3 0.025 mol/1 NaHCOs + 0.135 mol/1 NaCl 0.160 1.0 0.017 1-4 6.7 0.114 0.163

4 liquor cerebrospinal is 0.160 1.4 0.023 1-5 4.9 0.085 0.125

5 plasma 0.160 3.2 0.053

This table shows that diffusion potentials of about 8 mV are found with the two phosphate buffers and much smaller values with body fluids and bicarbonate solution of I = 0.160.

Further it appears that the calculated residual liquid junction potentials are of the same size as the measured differences. Just applying a pH cell with an isotonic salt bridge thus means, in the case of the 0.160 mol/1 NaCl bridge, a shift in the pH scale of more than 0.1 pH with regard to the cell with a saturated KCl bridge, depending on the test solution.

Both cells have to have the same pH scale if they will be useful. So new buffers had to be

developed with NaCl added to limit the liquid junction potentials with regard to 0.160 mol/1

NaCl. The pH value of these NaCl containing buffers can be determined in cells with a saturated KCl bridge.

198
Calibration buffers were prepared from the two standard phosphate buffers of the NBS.

So much sodium chloride was added that the pH of solution X, measured with the cell con- taining the 0.160 mol/1 NaCl bridge became identical with the pH measured with the cell

containing the saturated KCl bridge. In our case, plasma was the solution to be measured.

The following buffers appear to meet these conditions:

Table 4. pH(Ss,l.j.) values of phosphate buffers with NaCl at 25 °C and 38 °C.

Composition of buffers pH{Ss,l.j.) at 25 °C pH(Ss,l.j.) at 38 °C

with NaCl

0.025 mol/1 KH2PO4

0.025 mol/1 NazHPO^ 6.686 6.660

0.130 mol/1 NaCl

0.008695 mol/1 KH2PO4

0.03043 mol/1 NaaHPOi^ 7.255 7.229

0.120 mol/1 NaCl

To agree with the pH(S) scale of NBS, we have attached an operational pH value (pH(Ss,l . j . )

to the standard buffers, which was measured with the hydrogen electrode-saturated KCl calomel electrode cell with regard to one of the NBS buffers. Redetermination by Drinker et al.

[5] resulted in 0.01 pH lower values.

ter

pH, CORNING, MAAS MODE

For checking purposes, the pH's of a number of blood samples were measured and compared

by means of the glass electrode-saturated calomel electrode and the glass electrode - 0.160 M

NaCl calomel electrode cells. From the hematocrit value of the samples that pH determination

with the cell with saturated KCl bridge was corrected for the cell effect. The average

difference between the pH values of 21 blood samples obtained with both cells was -0.001 pH

which, experimentally, was not significant. A glass electrode system with an isotonic salt

bridge as used in this investigation, is completely reliable, and because of its special way

of standardization perfectly equivalent to the conventional cells with a saturated KCl

199
bridge. Further, the pH cell with a 0.160 mol/l NaCl bridge offers the advantages that the

liquid junction potential at the boundary between blood and 0.160 mol/l NaCl is independent

of the concentration of blood cells and that it can be used without danger for continuous pH

measurement in patients. Until now, this bridge was successfully applied by Drinker et al.

using the Corning Model 165 instrument [5].

References

[1] Maas, A. H. J., pH determination of body fluids with a micro glass electrode and a saturated KCl bridge in the cell, Clin. Chim. Acta 28, 373 (1970a).

[2] Si ggard-Andersen , 0., The Aaid-Base Status of the Blood, 4th revised edition, p. 155

(Munksgaard, Copenhagen, 1974).

[3] Bates, R. G., Determination of pH^ Theory and Practice, p. 312 (John Wiley and Sons,

New York, 1973)

.

[4] Maas, A. H. J., ph determination in body fluids using a cell with an isotonic sodium

chloride bridge, J. Appl. Physiol. 30, 248 (1971).

[5] Drinker, P. A., Noonan, D. C, Ramanai.ah, N., and Tole, J. R., Use of a sodium chloride

phosphate buffer for pH standardization in a new blood gas analyzer with an isotonic

sodium chloride bridge, Clin. Chem. ^9, 1243 (1973).

200
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

THE CRITICAL CARE LABORATORY: A 10-YEAR PERSPECTIVE

Myron B. Laver and Domenic R. Misiano

Anethesia Laboratories

Harvard Medical School

Massachusetts General Hospital

Boston, Massachusetts 02114, USA

Fifteen years have passed since funds were first made available from the National

Institutes of Health to the Department of Anesthesia at Massachusetts General Hospital for

the evaluation of arterial oxygen partial pressure measurements as a clinical method for

early detection of impaired postoperative blood-gas exchange. Progress achieved after a few years of endless frustration with methodology has been exceedingly rewarding. Although

cause and effect are difficult to establish, one is tempted to conjecture that the birth of

intensive care as a separate entity followed shortly after the arrival of the P02 electrode.

The progress achieved from these early days characterized by endless frustration with

methodology, have been exceedingly rewarding. Acute respiratory failure in the patient

with previously normal lungs once considered a medical curiosity is now recognized as a national health problem and the routine measurement of arterial blood gases, particularly

Pq^, is an indispensable tool for early recognition, prevention, and therapy of abnormal

lung function. An intensive care unit without access to blood gas measurement seems inconceivable today. Certainly, if an attempt is made to support patients with mechanical

ventilation, lack of this facility may prove to be disastrous. An understanding of the

growth and present-day problems attendant upon blood gas analyses can be gained by re- viewing our experience for the past decade at the Massachusetts General Hospital.

In 1965, the importance of access to blood gases and pH for acute care was recognized

by a greater number of our professional colleagues which, combined with the pressures

imposed by growing clinical demands on a research facility, prompted formation of the

Anesthesia Blood Gas Laboratory.

Remarkable foresight was exercised by the hospital administration and director of the

Chemistry Laboratory (Dr. S. Rieder) when they insisted that the nature of the laboratory

demanded supervision and control by physicians and/or personnel participating closely in

the care of critically ill patients.

Figure 1 indicates how the demands upon the laboratory have evolved since its inception. By late 1975, we will have performed a grand total of 700,000 blood gas analyses,

each one including Pq^, PcO?» P'^- ^^^^ additional responsibilities listed in table

1, the laboratory now provides an average of 1049 analyses daily for 82 patients, all

categorized as being critically ill.

These figures are not presented to impress. Rather, they go to the very heart of the

problem we have gathered to discuss: (1) what are valid criteria for quality control, and

(2) how is standardization of sensors and their calibration for every day use most appropriate?

Separation of ours from the main hospital laboratory was considered mandatory if a quick turn-around time was to be achieved. It was our original contention that a maximum

period of 10 minutes be allowed for the time between arrival of the sample to the laboratory and a return telephone call to the patient's bedside with the answers. I regret to say that although we are not far away from the original limit, few requests find their way

back within the ideal period. In the early days, problems with electrode technology were

201
1 20 - B Total Determinations (Blood Gases)

1965 1966 1967 1968 1969 1970 1971 1972 1973 1974

Figure 1. Yearly number of blood gas analyses performed in the Critical Care

Laboratory of the Massachusetts General Hospital since its inception in 1965.

Table 1. Daily load in the Acute Care Laboratory.

Average number of patients studied/day = 82

Average number of individual analyses/day = 1049

Analyses : PO2 + PCO2 + pH

O2 content

Serum [K"^], [Na"^]

Serum [Ca^^]

Urine [rt, [Na"^]

Hematocrit

Serum Osmolarity

Urine Osmolarity

Total Protein

Cardiac Output

Total determinations: 10

Average number of analyses/open heart patients/day = 38

the principle reason. Today, we are overwhelmed by sheer numbers and delay is usually

caused by our administrative inability to process the numerous samples pre- and postanalysis.

V The laboratory is located adjacent to the Operating Room, the Recovery Room, the

Respiratory Unit and the Surgical Intensive Care area. Unfortunately, the demand from

other parts of the hospital is increasing as the incidence of acute respiratory disturbances is being recognized with greater frequency. Personnel coverage is shown in table 2. Twenty percent of the total technician time is spent in handling administrative detail

unrelated to the measurements. The sensors have changed little while improvement in calibration, sample handling, and display of data has been significant (fig. 2).

202
Table 2. Personnel deployment.

Monday

to

Friday

Saturday

and

Sunday

Hours

7-3

3-11

n-7

7-3

3-11

11-7

Number of

Technicians

5

4

3

Total technician-hours/week = 624

Average number samples/week = 1049 x 7 = 7343

Average number samples/technician-hours = 7343 =11

A technician performs 1 analysis every 5.1 minutes.

TechnicianHours/Day

40

32

24

96

24

24

24

72

AUTOMATED SUBSTATION I

AUTOMATED SUBSTATION II

Blood Determinations

Na*, K*, Ca*

Hct

CORE LABORATORY

DATA

COLLECTION

Routine Preventative Maintenance

At 2-4 Hour Intervals

PRINT-OUT AUTOMATED SUBSTATION III

Technology: Mass spectrometer

Pol/electrolyte Membranes

Electrodes

Figure 2. Potential plan for a decentralized automated critical care laboratory.

The substations will be tied in with a core laboratory where the data are collected and stored. Errors and diagnostics are recorded continuously in the core laboratory from the substations.

Present problems include: (1) lesser control over individual technician performance,

(2) prolongation of the turn-around time, and (3) excessive loss of personnel hours when

samples are hand carried from distant parts of the hospital.^ Points 2 and 3 suggest that

^ We have calculated that approximately 6 miles are logged daily to hand carry samples

from the bedside or every labyrinth of the hospital to the laboratory. If one includes

the return trip, then the daily total is closer to 12 miles.

203
decentralization within this system must be considered. Together with a diminishing supply

of motivated personnel these problems have emphasized the need for automation which will

permit a short turn-around time assuming no loss of quality control. In order to improve our perspective for standards of accuracy we will review briefly the logistics that go into

determining performance of an acute care laboratory. They apply equally well when attempting

to improve an existing service (see table 3).

Table 3. Logistics to be considered in establishing a critical care laboratory.

1. Geography

Where is the laboratory to be located in relation to ICU, recovery

room and operating room?

Will the demands of other parts of the hospital be met?

2. Turn-around time

How long after a blood sample is drawn does the answer return to the bedside?

3. Quality control

("Are the results trustworthy?")

Instrumentation:

1. How often are the delays prompted by equipment

fail ure?

2. Is the service available from the manufacturer

adequate?

Personnel:

1. Is training adequate to recognize and check for

"bizarre" values?

2. Is there sufficient back-up to cover sickness

and vacation?

4. Cost

How much and how often should the patient be charged?

1. Geography

Geography, i.e., location in relation to the areas that need the service most. In the

smaller community hospital, where separation of services is less than extreme, a decision

as to location may be made with ease. In the larger institution such as a regional center

where building plans generally perpetuate architectural and administrative obsolescence,

centralization may be undesirable, if not impossible.

2. Turn-Around Time

Turn-around time, i.e. , the time elapsed between withdrawal of the blood sample and

return of the results to the bedside. If the samples are hand carried by non-physician

and/or non-committed personnel, the possibility of breakage and loss must be considered.

If the laboratory facilities are strained {e.g., numerous analyses requested at one time),

then a delay is inevitable even if the sample arrives promptly in the laboratory.

204
3. Quality Control

Quality control, i.e., consistent "trustworthiness" is the most difficult goal of all to achieve. Electrode technology is still an art, not science. In fact, the weak link is

still the membrane-covered O2 cathode. Although we have experienced significant qualitative

improvement in recent years, the overall caliber depends on the experience of the technicians

and their willingness to provide the extra effort necessary for frequent and appropriate

calibration. Unfortunately, the gap between corporate profit motive and desire for ex- cellence is still wide. Opinion is far from unanimous on the manner of calibration and

many laboratories will modify the manufacturer's instructions to suit local needs. There

is little reason left to continue Pq^ electrode calibration with liquids that provide

readings equal to blood. The error introduced by using gas-equilibrated water, and in some systems, gas alone introduces errors that are of little clinical importance. Few will

argue that tor clinical purposes, a error of 5 percent above 150 torr is unacceptable

although improvement on this figure is most desirable. Because of the importance of

obtaining a prompt answer, improvement in performance cannot ignore speed. Any device that

does not allow a technician to process 15 samples per hour is unlikely to prove useful.

4. Cost

Cost has received little attention in the past. Our economy no longer permits such

luxuries. Most laboratories have followed the pattern established for other services: one determination, one charge. If we agree that the acute phase of respiratory failure requires

multiple blood gas analyses at short intervals, then the cost to the critically ill individual

may reach astronomical proportions. One must consider the possibility that excessive cost

will encourage a reduction in tests ordered. In order to prevent this trend, our laboratory

initially instituted the "three unit charge" principle, or a maximum 24-hour charge per

patient calculated on the basis of the cost for 3 analyses. All tests in excess of 3 were

free of charge and physicians were encouraged to take samples as often as required by the

patient's condition. In fact, it is not unusual for a patient in severe respiratory

failure to have arterial blood samples drawn and analyzed at 15 to 30 minute intervals.

Clinical experience suggests that this approach can go a long way toward prevention of

catastrophies.

Because of the growing complexities in accounting, we have now established a fixed per

patient per 24-hour cost with no limitation on the number of analyses performed. It may

appear surprising, but effective administration makes this service quite equitable.

Finally, a comment on in vitro versus the in vivo blood gas monitoring.

When our laboratory first became a hospital service, a commitment had to be made

regarding the type of monitoring most suitable for a large hospital population subject to a high incidence of respiratory failure. The intermittent route was chosen for four reasons. First, it permitted better quality control. Second, it allowed all services to take

advantage of this facility. Thus, we did not restrict our services to the intensive care

areas. Third, maintenance and fail-safe technology remained the responsibility of a small core of trained technicians. Fourth, it was more economical. Nothing that has transpired

so far has given reason for a change in attitude.

5. Future

Reponsibil ities of the Critical Care Laboratory have grown. As a result, its efficiency has suffered. It is fair to say that we now face two problems, both related to stability: the membrane-covered Pq^ electrode and personnel. The latter can be partially

solved by a reduction in numbers, but not without increased automation. In fact, increased

automation and decentralization are the directions for the future. Ultimately, automated

substations will be located at several sites in the hospital. The problems there will be

the same as we face today and faced ten years ago: formation of clot and electrode drift.

Ancillary methodology requires critical' evaluation. Mass spectrometry, modifications of

205
gas chromatography, dielectric dispersion spectrometry^ deserve serious consideration to overcome the electrode drift problem. In the meanwhile, blood gas electrodes will be with

us for years to come. The core laboratory will provide around-the-clock preventive

maintenance service with the option of assuming responsibility for analyses if the substation is down. Details of substation design must incorporate the "building block"

approach to facilitate replacement of malfunctioning parts. Although this may appear as an enormous task, I regard it as most feasible and even deserving of grant support.

It would be helpful to predict that our ability to prevent acute respiratory failure

has advanced sufficiently to obviate an investment into a critical care laboratory. There

is nothing in the natural history of the disease to suggest such a course. In fact, trauma

and advanced heart disease are responsible for the most frequent causes of acute respiratory failure that require the services of a critical care laboratory. Unfortunately,

"inadequate gas exchange" is easy to recognize while "adequate gas exchange" is a difficult

diagnosis to make. Until such time that this statement will be obsolete, it is not likely

that we will be out of business.

Dr. Henry K. Beecher, Professor of Anesthesia, Emeritus, gave the senior author

unflinching support in the early days when blood gas electrodes and respiratory care were

in their infancy; Miss Anna Murphy proved to be a loyal and critical assistant whose

administrative devotion made the Critical Care Laboratory a reality. Finally, a .laboratory

working under pressure can only survive if both interest and morale are maintained at high

level. We have been fortunate to receive superb support from numerous technicians.

Although not usually the recipients of accolades, it is they who have made this endeavor a success.

2 Dr. A. Michaels, Alza Research, Palo Alto, California, personal communication.

206
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

A THEORETICAL AND PRACTICAL ANALYSIS OF MICROELECTRODE BEHAVIOR:

THE THREE-SHELL MODEL

R. G. Buckles^, H. Heitmann^, and M. B. Laver^

A study was undertaken to examine the behavior of membrane-covered microelectrodes as predicted by a three-dimensional diffusion model. The following parameters have been

found to be of importance for accurate prediction of electrode performance: (1) cathode

diameter, (2) electrolyte layer thickness and permeability, (3) membrane thickness and

permeability, (4) temperature, and (5) sample description. The model does predict the

previously observed variations in electrode sensitivity, its dependence upon temperature,

and the discrepancy we have noted in readings between gas, water, and blood samples.

The influence of sample viscosity is also predicted accurately. The changes that occur

within the membrane including their influence on overall electrode stability and sensitivity

have been considered. The signal decay observed with the electrode in a microcuvette has

been discussed as an extension of the information provided by the theoretical model. This

has allowed us to consider also the manner in which cuvette design influences electrode

performance.

1. Introduction

Polarographic oxygen electrodes which incorporate a small metal cathode have found

wide applications for biophysical research and are now used routinely in the measurement

of blood gases. When used in vivo, the electrodes allow for the measurement of PO2

in flowing blood, in different tissues, and even at the cellular level [1-5]'*. In vitro

applications include the analysis of oxygen partial pressure in the gas phase, liquids,

and whole blood [2,3,6-8], the measurement of oxygen reaction rates in the HartridgeRoughton fast-reaction apparatus [9], and finally, evaluation of tissue metabolic rates

[10]. These applications are possible because the microelectrode consumes very little

oxygen and membranes with high diffusivity for oxygen can be used to cover the electrode

surface. The original purpose in using a covering membrane was to obviate cathode poisoning

due to protein deposition [2]; however, an important function of these membranes is to isolate the oxygen diffusion field surrounding the cathode [3,4,11].

Proper performance of microelectrodes is dependent upon a stable and reproducible

diffusion field [3,4]. Such a field is generally more difficult to attain then the

potential fields required for thermodynamic electrodes, making the polarographic electrode

more sensitive to artifact. The sensitivity (amps output/mm Hg Pq ) is subject to constant

drift and is not independent of the sample characteristics [12], i.e., variations in

hydrostatic pressure applied during injection of a sample and temperature, cause rapid and

unpredictable variations in sensitivity. For example, the sensitivity for gas will be

different than for liquid samples with the same oxygen tension (P02) [6,11] stirred samples

read differently from stagnant ones [12], and sample viscosity causes variations in

sensitivity [1,13]. As a result of these complications, frequent calibration is mandatory

and, depending on the accuracy desired, the use of tonometered fluids of the same con-

^Present address: ALZA Research, Palo Alto, CA 94304

^Present address: Pius-hospital, 29 Oldenburg, Federal Republic of Germany.

^Present address: Anesthesia Laboratories, Harvard Medical School, Massachusetts General

Hospital, Boston, MA 02114.

"^Figures in brackets indicate "literature references at the end of this paper.

207
sistency as the sample has been recommended [12]. Because the latter approach is impossible with implanted electrodes, in vivo results are difficult to quantify [3].

Several mathematical models have been proposed to describe microelectrode behavior

[3,4,11,14-16]. However, none describe the behavior of the electrode even in a wellcharacterized in vitro environment. If a model is to be useful, it must predict quantitatively electrode sensitivity, the temperature dependence of sensitivity, and the influence

of different membrane characteristics on cathode performance. Both the simple spherical

[15] and linear diffusion models [3] predict a much lower sensitivity [11,12] and a greater

temperature and membrane thickness dependence than observed experimentally [12].

The literature also indicates the electrode behavior depends on the environment in

which it is used. Some authors have described a linear relationship between temperature

and sensitivity [9]; according to others, this relationship is exponential [12].

There are reports that the sensitivity is greater in water than in air [11] and vice

versa [6]. In small cuvettes, the sensitivity in liquid samples decreases with time and

anomalous signals may be recorded during sample injection [12].

Because of inconsistencies in the literature, we reported earlier our studies with

several commercially available cathodes maintained in a wel 1 -characterized environment in

vitro [12]. We now describe a theoretical analysis of electrode behavior with numerical

solution of the three-dimensional diffusion process. The latter is compared to the idealized

models often used in biophysical modeling. The microelectrode serves as a prototype for a study of transmembrane transport problems and thereby exemplifies the influence of spatial

geometry on the behavior of mathematical models dealing with diffusion.

A. Mathematical analysis

Figure 1 is a schematic presentation of an electrically isolated microelectrode. It consists of a fine wire (diameter < 0.050 mm) imbedded in an inert matrix (glass or epoxy).

SAMPLE

////////// />1EMBRANE// / / V///////A

ELECTROLYTE

1'.

?S<;iNERT<VVV

•\/ MATRIX nAXX^

1

' CATHODE

1

* y/^ }^ y/^

i

Figure 1. Simplified O2 cathode geometry, r^ = cathode radius; r = lateral, radial distance away from cathode center; X = vertical distance from cathode surface; te + t^^ = sum of electrolyte and membrane thicknesses.

while its end remains exposed. The surface of the inert material may or may not be smooth

as shown in the figure; most often it has a finite surface roughness that may reach a depth of 20 to 30 micrometers. This assures a continuous electrolyte bridge. Surface of

208
the cathode is also rounded but its radius of curvature is so much greater than the

cathode diameter (the latter defines the size of the diffusional field) that the surface

can be considered flat as in the vicinity of the cathode. If we accept the geometric

simplification shown in figure 1, we can characterize the diffusion of oxygen from sampl

to cathode in terms of a cylindrical coordinate system. The sample whose is being

evaluated may be either a stagnant liquid or a well-defined segment of tissue. If the

fluid is stirred, diffusion occurs only through a thin layer between sample bulk and

membrane, where it is assumed that the liquid flows in laminar fashion over the membrane

surface.

The diffusion of oxygen to the cathode is defined by the cylindrical coordinate

system, using the center of the cathode surface as the origin [17]. For a stagnant

sample, the differential equations describing the diffusion of oxygen from sample to cathode can be expressed as follows:

1 8 /r^P\ — 92p 9P i ^—± = -A^M

r 8r \ 3r / SX^ 9t

9 /9P \ 92p 9P

s \ , s m

Dm • - • — — + - (t < X < (t + t)

r 9r \ 9r/ X^ 9t ^ " " ^

1 9 /r9P \ 92p 9P

s \ , s e

D ^=-^ (OlX<tJ

r 3r \ 9r / dX^ 9t

^

where D = diffusivity (cm^/sec)

P = P02 (mm Hg) M = O2 metabolic rate divided by the O2 solubility coefficient (mm Hg/sec)

e = electrolyte

m = membrane

s = sample

X = vertical distance from cathode

At a large distance away from the cathode, the oxygen gradient disappears {i.e.,

(9P/9r) = (9P/9X) = 0); the gradient at the other boundaries defines the coupling of eqs

(1) through (3). Therefore,

9X

0 (X ^ -)

9P 9P m e (r -> 00) 9r 9r

9P

I — ] X = t^ + t^ mam V gx

/ m s

P = P

s m

D / m \ _ n / e mam —^ = ^ — X = t^

9X/ \ 9X/ '

P = P m e

209
kP

e

D ( 9X )

(X - 0, r < rc) (8)

e 0 (X = 0, r > r^) (9)

where k = O2 reduction rate at the cathode

r^ = radial distance from cathode center (fig. 4)

a = Bunsen solubility coefficient for O2.

For the initial conditions we may assume that a is a constant concentration of oxygen

present throughout the three regions, and at time zero the electrode is polarized. This

can be analyzed with eqs. (1) through (9) to give the electrode current (i) as a function

of time:

where n = number of electrons generated per O2 molecules reduced

= cathode surface area

F = Faraday's constant

It is commonly assumed that the reaction rate for oxygen at the cathode is infinitely

fast when compared with the diffusion rate (in eq. (8), k «). Under these conditions,

P (x = 0) = 0.0 and the O2 reduction rate is limited by O2 diffusion to the cathode [15].

However, the polarograms shown in figure 2, by the absence of a definite plateau, do not indicate any region of complete diffusional limitation, except perhaps at low oxygen

tensions. Therefore, the rate of reaction is included for completeness. The equations

presented here do not include initial conditions nor do they consider the fate of oxygen

once it has reacted at the cathode.

If the electrode is polarized at time zero, the recorded current must include also

the current due to ion migration in the electric field; the time required to establish a stable ionic electric field is greater than the diffusional response time and involves far

greater currents than are attributable to oxygen reduction. Therefore, it is best if we assume a fully polarized and stabilized electrode whenever eqs. (1) through (9) are used

to characterize cathode performance. The solution to the initial condition of constant

PO2 does indicate the response time to a change in P02 of the polarized electrode.

The "n" term in eq. (10) is normally assumed equal to 4.0. However, there is evidence

that oxygen reduction to hydroxyl ions is not complete on a stoichiometric basis.

Satterfield [18] has shown that H2O2 can diffuse away from the cathode and be converted

back to oxygen. Under these conditions one cannot be certain that exactly 4 electrons

correspond to one mole of oxygen reduction at the cathode (eq. (10)).

Although eqs. (1) through (3) are linear, the analytic solution is not easily obtained.

Therefore, we have utilized the Method of Zones numerical technique for their solution

[19]. Figure 3 presents the results of these computations, where cathode current and

sensitivity are given as a function of time. We have assumed that the electrolyte,

membrane, and sample are in equilibrium at a Pq^ of 100 mm Hg and the electrode reaction

rate is infinitely fast. The sample volume is taken to be 1 x 10-^° ml^ with an oxygen solu

^This sample corresponds to a 20 ym layer (unstirred) with the same radial dimensions as selected for the membrane and electrolyte (see fig. 1). A larger sample volume is more

complex, mathematically speaking, but would not alter materially the results.

(10)

X = 0

210
bility coefficient of 0.0238 ml {STP)/m1 -atm. The membrane is polyethylene, 25 micrometers

thick with the permeation properties shown in figure 3 [20]. The electrolyte layer

thickness is assumed to be 20 ym (see Sec. 2, Results) and with the same O2 permeation

properties as water,

45r

Instrumentation Lab (x 10-12)

1 40 \

1

35-

30

25

^ 20

15

10

Radiometer I (xlQ-'^)

Radiometer n (xlQ-'^)

Radiometer m (xlQ-'^)

Beckman (xlQ-'') Radiometer J2 (xlQ-'z)

0.4 0.6 0.8 \/OLTS

Figure 2. Effect of polarizing voltage on standard sensitivity

(^"'^STAND^ ^0

for ambient air of various commercially available oxygen electrodes. Sample: ^ stirred water at 38 °C. Membrane: polypropylene 0.015 mm.

5.0

o 4.0 -

» 3.0 I

E6

1.0

NUMERICAL SOLUTIONS, S

••STEADY STATE RESULT OF THREE-SHELL MODEL /

rc = IO^ te =20;i

tm = 25/i Pm=5.8 X IO''°m1(STP)-cm

CURRENT, I' CONDITIONS

Pe = Ps = 1.115x10"' ae =08 = 0238 atm -I om= .0502atm-l

cm'^ -cm Hg-sec

10 20 30 40 50

TIME, SEC. Figure 3. Electrode sensitivity (S), or the ratio of current (I) to sample Pq, was plotted

as a function of time. The sample P02 is the average value of the four-sample zones indicated in figure 4. For symbols see Appendix.

211
The current decreases rapidly for about 10 seconds, then levels off to decrease at a constant rate. During the rapidly decreasing portion of the curve, oxygen dissolved in the electrolyte and membrane (both at P02 = 100 mm Hg) is being consumed. Since oxygen

from both sites diffuses a shorter distance than oxygen from the sample, the rate of

arrival at the cathode surface is greatly increased. As this O2 is consumed, the P02 near the cathode decreases, thereby establishing a gradient from the sample to the cathode.

The current. continues to decrease linearly because the sample size is too small to allow

for a sufficient fraction of dissolved O2 to diffuse to the cathode. Since this rate of

decrease is linear, we must infer that the overall electrode response time (a complex

function of the diffusive properties of electrolyte, membrane, and geometry of the system)

is much less than the rate of oxygen depletion from the sample. Thus, although the sample

PO2 is decreasing continuously the concentration profile within the electrolyte and membrane

readjusts so, rapidly that the ratio of P02 remains constant with time.^ Presence of this

condition is supported by the graph which expresses electrode sensitivity (S) as a function

of time (fig. 3). After the initial transient period, sensitivity appears to be independent

of time (and therefore, Po2)- ^""^s constant for the appearance of a stable sensitivity

is 2.8 seconds, while the predicted time constant for linear diffusion through the membrane

(i.e., t2/Dni) is 7.1 seconds. The fact that the sensitivity does pass through a transient

period is contrary to the predictions of Hudson [14]. The steady state isobars for the

electrode which we described in figure 3, are shown in figure 4. The lines are seen to

Figure 4. Pq, isobars surrounding the cathode under steady state conditions. These

values are By numerical solution of eqs. (1) to (9). This consists of dividing each zone (electrolyte, membrane, and sample) into four radially-located zones and several zones in the X direction. 62 flux equations for each zone are derived from eqs. (1) to

(3) and solved simultaneously by using the boundary conditions described in eqs. (4)

and (9).

Hg.

Note in these considerations, the sample P^ has decreased from 100 to 73.8 mm

U2

form an approximately spherical surface around the cathode. P02 in the electrolyte at a radial distance of 90 m, or 9 times the cathode radius (r^ = 10 pm) is less than the

sample Pq while Po„ at the membrane-sample interface is witnin 0.6 percent of the sample

value. computing the P02 gradient in the electrolyte and membrane from these isobars,

one can see that over 98 percent of the oxygen flux comes by diffusion in a radial direction

within the electrolyte {i.e., it arrives by crossing the membrane in the region where

r > r^). If the electrolyte layer were thinner, this flux would not be significant.

^For a more detailed analysis of concentration-relaxation problems, see reference [21],

212
However, because of the large radial flux term, the amount of oxygen reaching the edge of

the cathode is greater than that which reaches its middle. Thus, the electrolyte layer

thickness (ELT) is an important determinant of total flux because an increase in ELT permits oxygen transport across a larger membrane area and results in a higher measured

electrode sensitivity.

From these studies of diffusion patterns in the peri-electrode space, it became apparent

that a geometrically simpler model would provide a useful tool in our attempt to account for

the influence of ELT on electrode performance. In the spherical model shown in figure 5,

P02-'"

< re

' rm

Figure 5. Idealized 3-zone model of the oxygen microelectrode, electrolyte, and membrane

complex. The radial P02 gradient under steady state conditions assumes a significant

resistance in each zone. For symbols see Appendix.

the cathode is treated as a hemisphere surrounded by hemispherical shells of electrolyte

and membrane. The actual cathode radius is used but the current or

is defined in terms of a spherical area on the assumption that oxygen flux density is

uniform over this cathode. An increase in area compensates for the increased flux

density near the outer edge of the cathode as determined in the previous model and, by

increasing the ELT, the effective membrane is also increased, thereby increasing sensitivity. We have used the model only to predict the steady-state electrode sensitivity

because this is most readily shown for a spherical model [15]. The oxygen transfer rate

from sample to membrane surface is characterized by a transfer coefficient which can be

expressed as a function of stirring rate. The same boundary equations apply as given in eqs. (4), (5), (7), and (8). While the boundary conditions expressed by eq. (6) are replaced by:

D (— I = k(P, - P„) r = r^. mam

y ^^1 ^ s m' m

213
The mass transfer coefficient, k, is used to allow for the effect of stirring. Also,

since convection is never absent in the cuvette chamber (particularly when one is using

whole blood), use of eq. (6) provides only an approximation. As there are several empirical correlations available for predicting k, its incorporation into the model does not seriously limit the model's utility. (See Appendix for a description of one method for

predicting k.)

The steady-state solution for sensitivity (S) in the spherical 3-shell model is given

by:

S =

i 2^nF / r 1 t^ t^

22,400 kr^ P r r P r r m m m e e e c J

(12)

where Pm is the product of D and a, the so-called permeability coefficient. Using the same constants as were used in the numerical solution presented in figures 3 and 4, the 3-

shell model predicts a sensitivity (S) of 3.38 x 10"^^ amps/mm Hg, while the steady-state

value from figure 3 is 3.45 x 10"^^ amps/mm Hg. If one uses the simpler spherical model

(without an electrolyte layer), the predicted sensitivity will be 8.8 x 10"^^ amps/mm Hg,

while the linear diffusion model [3] predicts a sensitivity of 3.15 x 10-^^ amps/mm Hg.

The 3-shen model appears to agree favorably with the numerical solution, while the simpler

models predict a much lower cathode sensitivity.

Because of the reasonably good agreement we have found between the numerical solution

and the 3-shell model, we elected to use this more easily manipulated model for our evaluation of electrode behavior. Equation (11) indicates the properties of sample, membrane

and electrolyte required to evaluate the sensitivity of the electrode. Information required

include cathode size, membrane thickness, ELT, temperature, membrane and electrolyte

permeability, and finally, sample characteristics. Since ELT is usually not known, all

other information must be available in order to infer its value. The results to be re- ported indicate that all these parameters are indeed significant.

2. Results

Using data published by others and our own [12], we have calculated: (a) electrode

sensitivity, (b) temperature-dependence of the sensitivity, and (c) dependence of sensitivity on sample type. A technique for evaluation of the influence of electrolyte layer

thickness is also described.

A. Electrode sensitivity

Prediction of electrode sensitivity requires all the information described in eq. (12). Unfortunately, there are no reliable methods at present for determining the electrolyte layer thickness directly, therefore its value must be inferred. Failure to include

the electrolyte layer thickness has resulted in models that do not satisfy the observed

electrode behavior. Figures 6 and 7 indicate the dependence of electrode sensitivity on electrolyte layer thickness for two different membranes and several membrane thicknesses.

The sensitivity increases with increasing electrolyte layer thickness because the high

electrolyte permeability (table 1) does not limit oxygen diffusion so as to offset the

increased area

^'^ - f(^CATH ^

^e ' *m)'^

on the outside of the membrane. For each membrane thickness, several curves are drawn,

the lower one for stagnant water and the upper curve for a gas sample.

In figure 6 we have plotted our experimental data for two values of membrane thickness

on lines calculated according to eq. (12). These lines indicate the values for sensitivity

when the electrodes record P02 in air, stirred water, and stagnant water. For each membrane

214
0.015 0020 0.025

ELECTROLYTE LAYER THICKNESS (mm)

Figure 6. Electrode sensitivity as a function of electrolyte layer thickness for polypropylene membranes of different thicknesses. The curves are calculated with the 3-zone

model, i.e.y eq. (12), based on the following conditions: T = 38 °C; cathode diameter,

0.020 mm; sample, water equilibrated with ambient air.

thickness, the three data points obtained with a microelectrode predict the same electrolyte

layer thickness, irrespective of sample type. The electrolyte thickness so determined for

several electrodes covered with polypropylene fell in the range of 0.020 to 0.030 mm. It is

important to note that a change in sensitivity caused by a change in electrolyte layer thickness does not alter the difference between the readings obtained with gas and stagnant liquid

sample.

Similar data have plotted for two different thicknesses of Teflon FEP membranes at

38 and 25 °C (fig. 7). The derived range of electrolyte thickness is 0.0055 to 0.007 mm.

In figure 8 we have indicated the sensitivity dependence on cathode size for a 0.015 mm

polypropylene membrane at 38 °C using the derived electrolyte thickness of 0.025 mm.

In an attempt to measure directly the influence of electrolyte layer thickness on electrode sensitivity, we constructed the device (i.e., electrode holder) shown in figure 9. The micrometer pe'^mits adjustment of the cathode tip distance (any commercially available

microelectrode can be used) relative to the fixed membrane in steps of several micrometers. The hydrostatic pressure of the electrolyte is maintained constant by an overflow

hole. A sequence of runs were made by suspending the electrode in a thermostatted bottle

filled with water, stirred by a magnetic bar at 300 rpm, and through which we bubbled a gas of known oxygen concentration. The cathode was positioned at about 0.75 mm above the

membrane and a steady state reading obtained. The micrometer was then turned to the next

unit (0.032 mm/mark) and the new sensitivity determined. This procedure was continued

until the membrane appeared to be visibly distended.

Figure 10 presents the results of these measurements. Two extreme cases may be

estimated by simplified models and the theoretical maximum and minimum sensitivity computed. Maximum sensitivity (S|vj/\x) must occur at an electrolyte layer thickness where the

membrane no longer influences the diffusion field. In this case, the electrolyte will be

equilibrated with the sample P02 and the spherical-diffusion-in-a-homogeneous-medium model

215
0 0.005 0.010 0.015 0.020 ELECTROLYTE LAYER THICKNESS (mm)

Figure 7. Electrode sensitivity as a function of electrolyte layer thickness for different thicknesses of Teflon FEP membranes. The curves calculated with the 3-zone model

based on the following conditions: T = 38 or 25 °C, as indicated; cathode diameter,

0.020 mm; sample, water in equilibrium with ambient air, stirred or stagnant; membrane

thicknesses from 0.0125 to 0.050 mm as indicated. Experimental data were obtained from

Heitmann et at. [12]. Note that Teflon FEP (0.050 mm) exhibited the smallest theoretical difference between stirred and stagnant water; experimentally no significant difference was found between these two samples.

I

5

6.0

4.0

2.0 te =0.025 mm

T = sa-c

tn)=O.OI5mm (Polypropylene) / 1

Sample: Gas

1 1 1 0.005 0.010 0.015 0.020

CATHODE RAD/US (r^J./nm

0.025

Figure 8. Dependence of electrode sensitivity on cathode size under the conditions stated.

should apply [15]. This value is plotted in figure 10 and is about the same as the maxi- mum we have obtained experimentally. Minimum sensitivity (S^in) will be obtained when the

spherical diffusion flux is totally restricted to the membrane; i.e., when the electrolyte

216
Figure 9. Cuvette-electrode assembly which permits a change in electrolyte layer thickness. A. Electrode-micrometer assembly which allows for an accurate displacement of the

electrode away or toward the membrane. B. Housing. The electrolyte is contained

within the lower half of the plexiglas chamber. C. Membrane holder plate. D. Assembled

electrode.

0 minimum, m m+160 m+320 m + 480 oo ELECTROLYTE LAYER THICKNESS, fg (microns)

Figure 10. Influence of electrolyte layer thickness (ELT)(tg) on electrode sensitivity

(S). Maximum sensitivity (Sm/\x) is recorded at right (®) at an electrolyte layer

thickness which is large enough so that the membrane no longer influences the diffusion

field. Point at left (®) indicates the theoretical minimum sensitivity (S^in). At

Sm/\x» behavior of the electrode is membrane- independent and ELT dependent; conversely,

at Smii\| electrode behavior is limited principally by membrane diffusivity and solubility

for O2 and less by these properties in the electrolyte.

217
layer thickness is zero. The theoretical value, also plotted in figure 10, is much lower

than the actually determined minimum sensitivity. The asymptotic approach to the experimental minimum sensitivity suggests that one can never reach the level of a zero electrolyte

layer thickness. These results suggest that the electrolyte layer thickness in responsible for the changes in sensitivity of microelectrodes over the span of theoretical

values used.

B. Temperature dependence of electrode sensitivity

According to eq. (12), three temperature-dependent parameters determine electrode

sensitivity. They are: permeability of the electrolyte, membrane, and sample. It can be

shown that one need only consider behavior of the electrolyte layer and membrane in order

to determine the true electrode temperature sensitivity. Both of these parameters are given by the general form:

p T ^ p -AE/RT

i e (13)

where P and Ep are obtained from table 1. Within the limits of 0 to 50 °C, one can also

express the temperature dependence of electrode sensitivity in the following form:

6/RT

(14)

where Pi is obtained from table 1. Within the limits of 0 to 50 °C, one can also

tionality. The rate of change of sT with respect to the temperature, T, is not constant; rather, it decreases with increasing temperature. Thus, it is preferable to report the

temperature sensitivity in terms of the constant B.

Table 1. Oxygen permeability data.^

Membrane

Manufacturer and

Description

Permeabil ity

(Pi) Reference

Union Carbide (50%

crystal 1 ine)

Dupont (branched)

Dupont (linear)

Vorschein Fol ien-Fabrik^

(W. Germany)

Dupont

Dilectrix

Dow Corning

Mylar

S./xlO-Vl^^OO/f^T Polypropyl ene

Polyethyl ene

Polyethylene

Polypropyl ene

Teflon FEP

Teflon TFE

Sil icone rubber

Water

Mylar

^Because of manufacturing variations, one should consult the manufacturer for accurate

permeability data, if not using any of the sources listed.

^Expressed as (cm^ (STP) • cm)/cm2 . ggc • cm Hg).

^Marked with Radiometer P^ electrodes.

U2

8.9xlO-3e-9900/RT

2.3xl0-4e-8S0°/«T

1.7X10-V8090/RT

9.5xl0-^e-5700/RT

14.6x10"^° (38 °C)

4.7xlO-^^e^3140/RT

5.78x10"^^ (37 °C)

26

26

26

21

29

4

21

30

7

218
The influence of electrolyte layer thickness on the ratio of sensitivities at 25 and

38 °C is shown in figure 11. The inferred electrolyte thickness for polpropylene is 0.021

k.

«0

0.9

Teflon FEP

o

I

0.8

0.7

0.6

tm = O.OI25mm /

/

/

Exp. Range*

/

/

Exp. Range

Polypropylene

tfn=0.020 mm

Tcoth " 0 010 mm

0.5- *Heitmann et al, 1967

0 0.010 0.020 0.030

ELECTROLYTE LAYER THICKNESS, mm

Figure 11. Influence of electrolyte layer thickness (ELT) on electrode temperature dependence. The experimental ranges indicated were obtained from Heitmann et at. [12].

The two curves were calcualted from eq. (12) using the data presented in table 1.

to 0.029 mm and that for Teflon FEP is 0.007 to 0.008 mm. These ranges agree remarkably

well with those inferred on the basis of absolute electrode sensitivity.

Some authors have reported a much higher electrode temperature dependence, often

approaching that of the membrane [22]. However, the cathodes used were large (macroelectrodes), in which case diffusion is mainly linear. In such cases, the effect of

electrolyte temperature-dependence is not significant, but does not imply that the electrolyte layer is any thinner.

C. Influence of sample condition and characteristics on electrode sensitivity

As indicated by figures 6 and 7, the influence of sample stirring and the differences

in electrode sensitivity between gas and liquid samples were predicted satisfactorily with

only one value for electrolyte layer thickness. However, this is not a very sensitive

test of the 3-zone model. Measurement of in a stagnant liquid sample within a cuvette

results in the greatest reduction of sensitivity of any sample media; let us examine this case further.

Variations in sample viscosity influence both the diffusion coefficient and the

solubility coefficient of the sample [23]. By taking these influences into account, and a particular set of electrode conditions, one should be able to predict the influence of

219
sample viscosity on electrode sensitivity. Figure 12 presents the results of two such

calculations, one for 0.00625 mm Teflon FEP and the other for 0.020 mm polypropylene.

to

J:

.00

0.96 —

0.92

0.89

0.84

""cath

• 0.010 mm

T =3800

Stagnant Sample

Polypropylene

tm = 0.020 mm

te = 0.025 mm

Teflon FEP

tm = 0.00|2 mm

te =0.0065 mm

O 1.0 2.0 3.0 VISCOSITY OF GL YCEROL - WA TER

SOLUTION, cps

4.0

Figure 12. The influence of sample viscosity on electrode sensitivity. The experimental

data*, taken from Heitman et at. [12], while the two curves (solid and dashed lines) were calculated from the data of table 1.

These two curves agree within 1 percent of the experimentally determined points. When the

sample is whole blood, many other parameters besides viscosity must be considered {e.g.,

variations in "permeability," convective settling, etc.). For this reason, calibration of

the electrode with a solution of viscosity identical to the blood sample may reduce the

discrepancy between blood and gas but not improve the consistency of the results. Rather, one must empirically determine, for a given electrode and cuvette, the solution whose

viscosity results in the same relative reading as blood of a given hematocrit.

3. Discussion

We have presented a theoretical 3-zone model which predicts O2 cathode behavior by

taking into account membrane characteristics, electrolyte layer thickness, and finally,

cathode diameter. Variation in sensitivity with cathode size has been reported to be

nearly linear by Fatt [11], but this author assumed, on the basis of his model, that the

findings represented anomalous electrode behavior. This is not at variance with our data

(fig. 8). At a higher cathode size, the dependence does become of second order as predicted by linear diffusion models. As a matter of fact, any model that does not take into

account the electrolyte layer thickness will predict a second order relationship between

sensitivity and cathode size.

Staub [9] reported a microelectrode sensitivity of about 0.6 x lO"'^^ amp/mm Hg for a 0.050 mm diameter cathode, a value much lower than predicted by the 3-zone model for tlji^ thinnest available Teflon TFE membrane (0.00625 mm; sensitivity approximately 10 x 10"

amp/mm Hg). This implies that when » t^ linear diffusion limits the electrode sensitivity.

The possible influence of other parameters on the behavior of microelectrodes does

require consideration.

220
A. Membrane characteristics

The membrane serves two functions: (a) it is the medium in which the major oxygen

gradient is established, and (b) it is the electrical isolator for the cathode. Polypropylene is the most commonly used membrane for static measurements. Others include

Teflon'' (both FEP and TFE), silicone rubber, and Mylar^. Ideally, the oxygen cathodemembrane system should exhibit no change in sensitivity with time in order to reduce the

requirements for frequent calibration. Unfortunately, none of the membranes presently

available fulfill these specifications, and we must consider their properties in more

detail in order to anticipate their influence on cathode characteristics.

The above mentioned membranes are not truly homogeneous. Silicone rubber is always

reinforced with silica particles (approximately 25 percent by weight). The other membranes have both crystalline and amorphous regions. Michaels and Bixler [24] have developed a model for permeation of gases through such membranes based on the concept that

crystalline regions are impermeable to gases and act as an obstruction in an otherwise

unmodified membrane. On this basis, permeability of the membrane can be expressed by the

following equation:

p

= -^^P/'^T (15)

'^m ^ e

where (t)Pg = P from eq. (13).

The factor cj) is generally a linear function of percent crystal 1 inity. The latter

depends on the manufacturing process, thermal treatment, and degree of stretching. Recent

studies with polypropylene have shown a three-fold variation in permeability induced by

mechanical stress [25]. These same studies indicated that stretched membranes will gradually "relax", resulting in a gradual increase in permeability subsequent to the initial

stretching.

This phenomenon probably explains the cathode "aging" patterns which have been observed.

The initial stretch which results when the membrane is first applied reduces permeability,

but the subsequent stress-relaxation increases permeability, and therefore increases

sensitivity with aging. However, there is an alternative explanation. Teflon-covered

electrodes show a greater increase in sensitivity with "aging" than polypropylene [12].

Since stress-relaxation is accomplished by an extended deformation of the membrane, the

increased sensitivity may be the result of a thicker electrolyte layer. To check this

hypothesis, we examined the electrode temperature coefficient, 3, on three succeeding

days. As the sensitivity increased, g decreased, indicating a thicker electrolyte layer.

Both explanations depend on the stress-relaxation phenomena, common to all the presently

used membranes. At this time, there are insufficient data to decide between these two mechanisms.

The observation that polypropylene membranes produce electrolyte thicknesses of 0.020

to 0.030 mm has two important effects on electrode performance: (1) hydrodynamic drag can certainly be expected to effect the electrolyte thickness, and (2) the temperature sensitivity cannot be accurately predicted for a given electrode. Although this second fact is

of great importance when one is considering the use of thermistors for automatic temperature compensation in oxygen electrodes, it is of little consequence in thermostated systems

used routinely for static blood-gas measurements.

Table 1 presents the available permeability data for membranes commonly used with

oxygen cathodes. The data from Buckles [21] were evaluated with polypropylene membranes

available with one type cathode (Radiometer) and are distinctly different from the data of

Myers, Stannett, and Szwarc [26]. This difference is undoubtedly due to variations in

crystalline content [27].

Stiffness of the membrane determines its fit on the cathode surface. The flexible

materials {i.e.^ Teflon FEP) can be expected to conform closely to the shape of the

cathode tip and to fold more smoothly over the outer cathode edges. The degree of stressTrademarks of E. I. DuPont de Nemour Company.

221
relaxation, or "creep", is a function of the initial stretch; because of its lower "toughness". Teflon will exhibit the greater amount of stress-relaxation (for more detailed

discussion of the mechanical properties of polymers, see ref. [3]).

B. Cathode geometry

Shape of the cathode tip has a pervasive influence on cathode sensitivity and stability.

Since we are considering cathodes with diameters less than 0.1 mm, the radius of curvature

is always much greater than the cathode radius. Yet in a previous section, we have shown

that the surface need not be sharply pointed to result in a near-spherical diffusion

pattern. Attempts at reducing the radius of curvature of the electrode tip to increase

"spherical" diffusion is virtually impossible with microelectrodes. Therefore, even pointed electrodes are still "flat" in the region where diffusion occurs. A pointed

electrode tip will result in an increased stress to the applied membrane; as a result,

cathode instability is prominent.

The microelectrode has been developed to eliminate the mechanical inconvenience of

stirring and because it is desirable to use small blood samples without depleting its

oxygen electrodes and their cuvettes have varied considerably and, as expected, the

results have been inconsistent. According to Fatt [11], water flowing through a cuvette

resulted in a signal deviation of 20 percent as compared with the stagnant sample. On the

other hand, stirring did not result in any variation of the signal from that of the

stagnant sample. Explanation for this disagreement was not given. If a sample flows past

an electrode in one direction, hydrodynamic drag occurs in addition to surface renewal of

sample. The hydrodynamic drag is not present when the sample is stirred symmetrically

within the cuvette. Hydrodynamic drag is mainly dependent on the linear velocity across

the membrane surface and sample viscosity (the linear velocity can be approximated by:

electrode diameter x sample flow rate, ml /sec/cuvette volume in ml); it influences the

electrode signal by modifying the membrane position relative to the cathode and by altering the electrolyte layer thickness.

The cuvette design also influences the long-term stability of the signal from a stagnant sample. Addition of a sample to the cuvette, previously equilibrated with a sample of lower P02 generally results in the current rising swiftly to a peak signal, then

slowly dropping off. If the cuvette is a thin, dish-shaped space, the capillary forces

will tend to hold the sample rigidly so that oxygen transfer from sample to membrane

occurs only by diffusion. Thus, once the oxygen is consumed around the cathode, oxygen

will have to diffuse farther to reach the cathode, resulting in a lower signal. In terms

of the 3-zone model described above, one may say that

where t = time in seconds.

The time constant, y, will depend on the cuvette design, being largest when the

spacing is greatest. When the sample is not contained by capillary forces, natural con- vection will enhance simple diffusion of oxygen.

C. Cuvette design

k = k„e

0 t/Y

(16)

4. Appendix

C(r)

A,-

D,

P,

F

area normal to flux, j^, at a surface i (cm^)

oxygen concentration, [ml (STP)/ml] at a radius r oxygen diffusion coefficient of the material i (cm^/sec)

Faraday units, (96,500 amp-sec/mol)

oxygen permeability of material i, expressed as

222
ml (STP) - cm cm^-sec-mm Hg

(17)

R = gas content (cal/g mol-K)

= electrode sensitivity i.e.,

[current for P^, of ambient air] - [current when P^, = 0] U2 U2

Pn (mm Hg) for ambient air

U2 T = temperature (kelvin)

U^. and V^. = constants of integration

j-j - oxygen flux rate, ml (STP)/sec, across surface i

T = PA U2-I

•^i (18)

"i

k* = sample mass transfer coefficient, [ml (STP)/cm2 - mm Hg - sec]

*The mass transfer coefficient to a spherical particle of radius can be estimated

as a function of the dimensionless velocity in a variety of ways [28]. However, at

low Peclet Numbers (below 1.0 (see ref. [24]), the dimensionless transfer coefficient,

"k"/D, reaches a limiting value of 1.0. Thus,

k = D a /r^ X 760. (19) m m m

When stirring occurs, one generally finds that the ratio of the sensitivity with

stirring to the sensitivity in a stagnant sample for an electrode immersed in electrolyte without a membrane, is 2:1 [12]. Thus, we may safely say that the k to be used

when complete stirring occurs is k = 2P /r„ x 760. ^ ^ m m

"k" = mass transfer coefficient to spherical particles, mol/mol sec

n = number of electrons that react to reduce one mol of O2

n^. = vector normal to surface A^.

r = radius (cm) r„ = r^ + t„

e c e

r = r + t + t m c e m

t = thickness

a = Bunsen solubility coefficient [ml (STP)/ml-atm]

e = effective activation energy for oxygen diffusion in a micro-Clark-type oxygen

electrode (cal/g mol

aE = activation energy of permeation (cal/g mol)

Y

= time constant for oxygen diffusion in a stagnant sample

(J) = permeation factor; reflects dependence of permeation on crystalline content

Subscripts:

m = membrane

e = electrolyte

s = sample

c = cathode

O2 = oxygen

223
This investigation was supported by U.S. Public Health Service Grant GM 15904-08.

References

[I] Charlton, G., A microelectrode for determination of dissolved oxygen in tissue, J. Appl. Physiol. 16_, 729 (1961 ).

[2] Clark, L. C, Jr., Monitor and control of blood and tissue O2 tensions. Trans. Amer.

Soo. Art. Internal Organs 2^, 41 (1956).

[3] Davies, R. W. , The oxygen cathode, in Physical Technique in Biological Research^ Vol

p. 137 (Academic Press, New York, 1962).

[4] Grangsjo, F., and Ulfendahl, H. R., Factors influencing the properties of electrodes

for the continuous measurement of oxygen tension in tissues. Acta Soa. Med. Upsal.

67, 107 (1962).

[5] Kreuzer, F., Harris, E. 0., Jr., and Nessler, C. G., Jr., A method for continuous

recording in vivo of blood oxygen tension, J. Appl. Physiol. ]5^, 77 (1960).

[6] Laver, M. B,, and Seifen, A., Measurement of blood oxygen tension in anesthesia.

Anesthesiology 26 , 73 (1965).

[7] Polgar, G., and Forster, R. E., Measurement of O2 tension in unstirred blood with a platinum electrode, J. Appl. Physiol. jl_5, 706 C1960).

[8] Silver, I. A., A simple micro-cathode for measuring ?q in gas or fluid, Med. Elec.

Biol. Eng. 547 0963).

[9] Staub, N. C, A simple small oxygen electrode, J. Appl. Physiol. 16, 192 0961).

[10] Hospodka, J., Caslansky, Z., Beran, K., and Stross, F., Continued Cultivation of

Microorganisms, 2nd Symposium, Prague, p. 353 (1962).

[II] Fatt, I., An ultramicro oxygen electrode, J. Appl. Physiol. 1^, 326 (1964).

[12] Heitmann, H., Buckles, R. G., and Laver, M. B., Blood Pq^ measurements: performance

of microelectrodes, Resp. Physiol. _3, 380 (1967).

[13] Tsao, M. v., and Vadnay, A., An electrode for continuous measurement of transient

blood P02 in the vessel, J. Appl. Physiol. 15^, 712 C1960).

[14] Hudson, J. A., Measurement of oxygen tension by the oxygen cathode, Med. and Biol.

Engng. 5, 207 (1967).

[15] Kolthoff, I. M., and Lingane, J. J., Polarography, Vols. I and II (New York,

Interscience, 1952).

[16] Schuler, F., and Kreuzer, F., Properties and performance of membrane-covered rapid

polarographic oxygen catheter electrodes for continuous oxygen recording in vivo.

Prog. Resp. Res. 64 (1969).

[17] Sherwood, T. K. , and Satterfield, C, The Role of Diffusion in Catalysis (AddisonWesley Publishing Co., Reading, Mass., 1963).

[18] Satterfield, C, Supersaturation of oxygen in aqueous hydrogen peroxide solution, J. Chem. Eng. Data 504 (1961).

[19] Strong, P. F., and Emslie, A. G., The method of zones for the calculation of tempera

ture distributions, report to JPL, Arthur D. Little, Inc., April 1963.

224
[20] Bixler, H. J., Solution and Flow of Gases in Polyethylene, Sci . D. Thesis, Massachusetts

Institute of Technology, 1959.

[21] Buckles, R. G., Ph.D. Thesis, Department of Chemical Engineering, Massachusetts

Institute of Technology, 1966.

[22] Carey, F. G., and Teal, J. M., Responses of oxygen electrodes to variables in

construction, assembly, and use, J. Appl. Physiol. 20, 1074 (1965).

[23] Reid, R. C, and Sherwood, T. K. , The Properties of Gases and Liquids (McGraw-Hill

Book Co., Inc., New York, 1958).

[24] Michaels, A., S., and Bixler, H. J., Solubility of gases in polyethylene, J. Polymer

Sci. 50, 393 (1961).

[25] Krewinghaus, B., Sc. D. Thesis, Department of Chemical Engineering, Massachusetts

Institute of Technology, 1966.

[26] Myers, A. W., Stannett, V., and Szwarc, M., The permeability of polyethylene to gases

and vapors, J. Polymer Sci. 35^, 285 (1959).

[27] Conner, W. P., and Scherts, 6. L., Effect of density and orientation on the permeability of polypropylene film, SPE Transaations 2^, 186 (1964).

[28] Levitch, V. G. , Physioaoahemiaal Hydrodynamics [Pi^entice-Hal 1 , Inc., Englewood Cliffs,

N.J., 1962).

[29] DuPont Bulletin T-3A, Teflon FER—Chemical Properties.

[30] Hodgman, C. D., Ed., Handbook of Chemistry and Physios (The Chemical Rubber Publishing

Co., Cleveland, 1963).

225

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DYNAMIC RESPONSE OF A pCOa ELECTRODE

S. J. Pace and M. J. D. Brand^

Technicon Instruments Corporation

Tarrytown, NY 10591, USA

1. Introduction

Gas sensing membrane electrodes are almost universally used for the measurement of the

partial pressures of oxygen and carbon dioxide in blood. The basic operating principles

have remained unchanged since the introduction of the PO2 electrode by Clark [l]^ and the

PCO2 by Severinghaus and Bradley [2], In each electrode the sample is separated from an electrolyte by a gas permeable membrane which allows transport of O2 and/or CO2. The diffusing gas is sensed in the electrolyte solution by: (a) reduction at a platinum or other

noble metal cathode for O2, or (b) measuring the change in pH in a buffer solution with a glass electrode for CO2.

It is remarkable that despite the widespread use of these electrodes for nearly 20

years there have been few attempts to describe their response theoretically. Lucero [3] has

considered some factors which affect the design of PO2 electrodes from a practical standpoint. It was not until 1974 that an analysis was made of the oxygen profiles around a membrane covered PO2 electrode [4]. Theory of the PCO2 electrode was limited to the approximate equilibrium model given by Severinghaus [2] until a contemporary treatment by Ross et

al. [5]. This model considered only the mass transport through the membrane. In the particular case of PCO2 electrode, the kinetics of the reaction between CO2 and water in the

electrolyte solution can be rate limiting.

In this paper, a model for a PCO2 electrode is proposed which considers mass transport

in the membrane and in the electrolyte solution as well as chemical kinetics.

2. Mass Transport of CO2

This model applies to a gas sensing electrode having the geometrical configuration

shown in figure 1. The response of this electrode is determined by the concentration of gas

(CO2) as a function of time and distance from the ion detector. Once the concentration

profile is established through both the membrane and the internal electrolyte phases, the

CO2 concentration may then be applied to the rate law governing the reaction:

fast

C02(g) + H20(ii)<=^ H2C03(£)

^ H + HCOi

^a

The homogeneous reaction is assumed to be rate-limiting and even though the molecularity of reaction (1) is second prder, the solvent mole fraction is very nearly one {i.e.,

C^^Q>>C^Q^) resulting in an overall pseudo-first order reaction. The experimental basis for

this assumption is the reported [6] decreased response time of the CO2 gas-sensing electrode

when carbonic anhydrase is incorporated in the internal electrolyte {i.e., the enzyme catalyzes

^Present address: Department of Chemistry, Texas A&M University, College Station, TX 77843^

^Figures in brackets indicate the literature references at the end of this paper.

227
EMF

E — pH Electrode Assembly

R — Internal Electrolyte Compartment

M — Membrane or Air Gap Compartment

S — Sample

pH ElectrodeSample In

0

M

Ref Electrode

Gas Permeable

Membrane

Sample Out

Figure 1. Geometric configuration of gas sensing electrode.

reaction (1)). Furthermore, the high sample concentration limit rarely exceeds 10"^ molar,

representing a solvent to solute mole ratio of several orders of magnitude. The concentration of CO2 in the sample is assumed to be invariant; this is equivalent to assuming that

the sample is buffered in CO2 and is stirred.

Consider the concentration gradient through the membrane and electrolyte media shown

in figure 2. A steady state concentration of gas A (CO2) is maintained in the sample

compartment S, The membrane (M) in this case is simply an air gap of mostly nitrogen gas

(B). The air gap may be maintained by a thin gas-permeable membrane delineating the

sample/air gap interface (although the entire compartment M may be composed of a gaspermeable membrane, e.g., silicone rubber). At the gas-liquid interphase S/M, which is denoted as X = 0, and through the membrane phase M bound at X = L, the gas phase concentration

(C/\) is expressed in moles/liter or as a mole fraction N/\. Na may, in turn, be defined as the partial pressure of A divided by the total pressure, PA/Ptot' provided A and B form an

ideal gas mixture. Gas B is inert to both the sample liquid and gas A and the partial

pressures of both gases are in equilibrium with the liquid S,

m corresponding to the gas-internal electrolyte interphase (M/R), the con-

" " "ntire system is controlled at constant temperature and pressur

L separates the internal electrolyte from the pH glass electrode

At X ... ^ ^ , , centration is Zp^^. The entire system is controlled at constant temperature and pressure

The interface R/E at X surface.

The boundary conditions for this mass transport problem are

X = 0 c, = C,^

X = m

^A " ^A

\3x /x^L

The condition at X = L states that no H3O diffuses through the pH glass membrane. This is

228
(a)

(b)

t Ca

Gas A

t Ca

Ca.

Gas A

Ca;

M

Gas B

m L

M

Gas B

(a) Low-to-high

concentration change

(b) High-to-low

concentration change

S — Sample

M — Membrane and Air Gap

R — Internal Electrolyte E — pH sensing Electrode

Figure 2. Concentration profiles

at electrode-sample interface.

m L X

a reasonable assumption since the diffusion coefficient of HsO"*" through glass is much

smaller than the diffusion coefficients of A in B (Dy^g) and A in R (D^l)-

The electrode is subjected to an instantaneous step concentration change from C^' to . 0 0

The resulting concentration gradient is the driving force for molecular diffusion of A

into M and subsequently into R. Solving the mass balance equation through region M, one obtains the following equations:

\ m 0 ' c (1)

or

(Ptot - PaJ \Ptot - Pa^

x/m

The flux at the gas-liquid interface M/R is:

°AB

(^A)x=m = - IT (\ - ^A^)

Ptot Dab/^^ m

In

1 - P,

1 - P, (2)

As gas A diffuses through medium R, it is chemically transformed to compound Q (H2CO3) by

229
reaction (1). The total mass transport equation including homogeneous chemical kinetics is:

- (ki + - k2 (C^^ - CqJ = 0 (3)

where C„ and Cp, are the initial concentrations of A and Q. The solution to the above equation % ^0

for the above boundary conditions defines the concentration profile through region R:

where

and

- _ m_

ki + k2 k2

a\l ^ a^^ (2L - m) Y

= e + e

The flux at X = m is;

(5)

From eqs. (2) and (5) an expression for C. is derived;

m

C m ( p- ) a ? - 1

where

(2L

(6)

The removal of geometrical dependence on the ion sensing element in the internal electrolyte can be advantageous for practical applications. This could be accomplished by

230
designing a sensor so as to make it responsive to an average ion concentration as, for example,

a cylindrical pH sensor enveloping the internal electrolyte compartment. The concentration

of A may be averaged throughout the entire volume of compartment R resulting in the following

relationship:

avg m m

avg.

dx

(L - m)a^ (L - m)

(7)

3. Time Response

Equations (1), (2), (4), and (5) accurately describe within the confines of the model

the steady state concentration of CO2 (gas A) and the rate of mass transport. However, no information about the response time of the electrode can be derived from the above treatment

thus far. Since the CO2 gas sensing electrode is slow responding due to a rate limiting

chemical reaction, eq, (4) may be applied for the appropriate distance X (at X = L in this

case) to the rate law for reaction (1). An expression for the time response is obtained

by solving the first order rate equation for reaction Cl)-

dC,

(8)

The solution of eq. (8) yields an expression for the time required to reach the equilibrium concentration C. at distance X = L:

r C In

0 e

e

(ki + k2)t

The following mass action law can then be applied:

+

l<i[C02]

kzCHCOg]

(9)

to describe the time to reach a given pH change as;

t = (ki + k2)"^ In X% ' ^ (% ^(HCOP - ^H^ ^(HCOj) J

^b(V - ^Hp,^ ^(HCOs)^ " ^H^ ^(HCOi)^

(10)

where Cg is the concentration of internal electrolyte (usually NaHCOs) and C^^q- is that

231
formed from CO2. Equation (10) reveals that the time response is inversely proportional to the rate constants and logarithmically dependent on ion concentration changes although in

a complex manner. In the high concentration limit of detection, the square terms become

insignificant reducing the logarithmic term to the simple ratio of (Cu+ - Cu^•^ - fC,,+ - C,.+^

The time required to reach equilibrium via steady state mass transport may also be

determined. The flux at, i.e., X = m may be equated to d/C. Wdt;

dC,

'^m

^ m /

Integration of the above equation yields the following expression for time:

t = [DAL(ki k,)]-^Mn

j

j (11 m

where ? is the reduced concentration parameter fC„ - and ^. is defined as the fractional

reduced concentration. In terms of the average concentration change in compartment R:

t = [d^l^^i ' k,)]-"^|^ln

j [(L - m) C^^^ - 3] |^ | - Inj [(L - m) C^^^ - . c^gj^ . (12)

The reverse process considering a high to low step concentration change is similarly

treated, resulting in virtually identical relationships with the exception that

e' = [l<i/(ki + k2)](^CQ - ) substitutes the expression for 6.

4. Discussion

The sensitivity and detection limit of a CO2 gas sensing electrode are determined by

the equilibrium reaction (1) and by eq. (9). The detection limit is also a function of the

resolution of the pH electrode; for a limit of 0.01 pH unit the CO2 concentration detection

limit can be calculated from eq. (9) to be approximately 10" M. The equilibrium response

of the electrode is not of prime consideration here, but rather the time-dependent response

to changes in CO2 concentration.

Equations (1) and (4) give the steady state CO2 concentration profiles through regions

M and R following a step change in sample concentration. Figures 3 and 4 show numerical

solutions to these equations under different conditions. In figure 3, conditions correspond

to a near ideal CO2 sensing electrode. Diffusion of CO2 across the membrane region M is almost unimpeded, i.e., the membrane is an "air gap." However, it must be noted that

diffusion of CC2 into the electrolyte region R is of much greater importance, even though

the thickness of layer R is much less than M. The curve shown in figure 4 corresponds to

a more general case in which layer M represents a diffusion barrier, i.e., it is a gaspermeable plastic membrane.

232
0.10

0.08

0.06 Cone. Molar

0.04

0.02

M—IK

.001 .002 0.010

m — 0.01 cm

L 0.012 cm

Dal 1.1 X 10^ cm^ sec

Dab z.'t X lU cm sec

— 10'^ M

10"^ M

0.1 sec'^

^2 10^ sec"^

0.011 X (cm)0.012

Figure 3. Concentration profile for CO2 electrode where the membrane M is an

gap.

0.1

0.001

Cone. Molar

0.0001

0

0.1 0.2 0.90 1.00 X (cm)

m = 1.0 cm

L = 1.2 cm

Dal = 10"^ cm^ see

Dab = 10"^ em^ sec

10'^ M

^a:

= 10'^ M

a = 10^ (cm^)

1

^ = 10'^ M

cm

1.10 1.20

Figure 4. Concentration profile for a generalized gas sensing CO2 electrode

where the membrane M is a gas permeable plastic.

233
The time response of the electrode is given by eq. (10). Table 1 shows some calculated values for 99 percent response times for various concentration changes. It must be

noted that the time response for a high-low concentration change is greater than for the reverse. The magnitude of the differences in these response times is a function of the

homogeneous chemical kinetic rate constants and the diffusion coefficient of CO2 in the

electrolyte region R. In the case of the diffusion rate limited process, the thickness of

regions R and M effect the response time equation.

Table 1. Time required to achieve 99 percent of concentration change .

Initial Final t [Dn. (ki + ks)]"' eq. (11)

concentration concentration . /. , , s /t^n

(moles/liter) (moles/liter) ^ ^

IQ-^ 10"^ 4.61

10"^ 10"^ 4.61

10"^ 10"^ 4.61

10"^ 10"^ 9.21

10"^ 10"^ 11.51

10"^ 10"5 13.82

The results obtained using this model of a CO2 sensing electrode agree in principle with

the known experimentally determined properties. It is hoped that the model will lead to a clearer understanding of those factors which contribute to the electrode response and to improvements in electrode performance.

References

[1] Clark, L. C, Jr., Monitor and control of blood and tissue oxygen tensions. Trans. Amer. Soo. Artificial Internal Organs, 2^, 41 (1956).

[2] Severinghaus, J. W. and Bradley, A. F., Electrodes for blood PO2 and pCOa determination,

J. Appl. Physiol. U, 515, (1958).

[3] Lucero, D. P., Design of membrane-covered polarographic gas detectors, Anal. Chem. , 41_, 613 (1968).

[4] Gutherman, H. E. and Goldstick, T. K., Numerical analysis of oxygen profiles around a Clark oxygen electrode. Proceedings of the 27th Annual Conference on Engineering in

Medicine and Biology, Vol. 16, p. 399 (The Alliance for Engineering in Medicine and

Biology, Chevy Chase, Maryland, 1974).

[5] Ross, J. W., Riseman, J. H. and Kreuger, J. A., Potentiometric gas sensing electrodes,

in International Symposium on Selective Ion-Sensitive Electrodes, G. J. Moody, ed.,

p. 473, (Butterworth & Co., Ltd., London 1973).

[6] Severinghaus, J. W., Measurement of blood gases: PO2 and pC02j Annals of the New York

Academy of Sciences, 148 , 115 (1968).

234
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

MONITORING OF OXYGEN PRESSURE IN HUMAN AND ANIMAL BLOOD

H. P. Kimmich and F. Kreuzer

Department of Physiology

University of Nijmegen

Nijmegen, The Netherlands

The surveillance of arterial oxygen pressure (Pa02) in patients with pulmonary or cardiac failure by frequent sampling of blood has several obvious disadvantages. Continuous monitoring of Pa02, including recording of variations in Pa02 with respiration and

cardiac activity, provides more useful information, but this sophisticated monitoring by use of catheter electrodes is still affected by numerous problems. The size of the sensor

as well as stability, flow dependency and transient response are crucial factors. Low

flow dependency and fast response oppose each other [1] , but these requirements calling

for an optimal compromise in any individual situation have received relatively little

attention since a fast response has often been considered to be unimportant, particularly

in clinical monitoring. It cannot be ignored, however, that a fast response is often

necessary or at least desirable and may provide interesting accessory information [2-4],

We have been paying particular attention to this aspect for a number of years.

1. General Problems and Requirements of PO2 Electrodes

in Continuous Monitoring

The problems and requirements of PO2 electrodes in continuous monitoring in vivo include

technical as well as biological aspects. For the evaluation and appraisal of any such

system the following check list may be useful:

A. Sensitivity and accuracy including linearity, reproducibility, short-term stability,

long-term stability, shelf life, transient response, zero current (value and drift), flow

dependency, temperature dependency, sensitivity to CO2, sensitivity to other gases, sensitivity to pressure, to acceleration, and to water vapor (water deposition).

B. Design and constructive lay-out including miniaturization, safety, and complexity

of servicing and operation.

C. Sterilization.

D. In vivo parameters including biocompatibil ity (toxicity, blood changes, tissue

reactions, blood coagulation, etc.), mechanical and electrical compatibility (flow distortion,

tissue changes, electroshock, membrane potential changes), influences of the medium on the sensor (protein deposition, blood coagulation at electrode, mechanical stress, galvanic

reactions)

Intensive research during the past years has led to the solution of most problems

connected with the technical parameters sensitivity, accuracy, constructive lay-out, and

sterilization. For human applications, however, improvement of the constructive lay-out

remains desirable, particularly with respect to the size of the transducer. For prolonged

monitoring of human Pa02 the main issue, however, is the solution of the problems connected

with the in vivo parameters, particularly the biocompatibil ity of the catheter electrode and

the influences of the medium on the sensor. But also the mechanical compatibility and the

positioning of the transducer are important in certain circumstances. Biocompatibility of a catheter electrode is never complete but can be improved by a careful construction of the sensor and by use of proper materials including coating the tip of the catheter with silicone. Protein and blood cell deposition as well as blood coagulation at the tip of the

Figures in brackets indicate the literature references at the end of this paper.

235
electrode affect fast electrodes much more than slow ones. Since these phenomena also

involve a certain risk for the patient, it is advantageous that their effects on the fast

electrodes become noticeable quite readily and thus provide a warning at an early stage.

Trials of monitoring PaOz in the human radial artery using our ring electrode [5]

revealed the necessity of a catheter probe with a diameter smaller than 2 mm. First, this

problem was circumvented by measuring Pa02 in an arterio-venous shunt (as described below),

but simultaneously a new electrode with a diameter of 1.2 mm and essentially similar characteristics as previous designs has been developed. These two lines of approach will now be

discussed.

2. Monitoring of Human Pa02

Jank, et at. [6] applied our previously developed type of ring electrode [5] to patients

by inserting it into a specially made shunt between the radial artery and the antecubital

vein. A slight disadvantage of this system is a possible reduction of the peripheral circulation due to the shunt. On the other hand, this system offers several important advantages.

The markedly increased blood flow in the shunt (as compared with normal radial arterial

flow) decreases the risk of blood coagulation (which is further minimized by coating with

silicone and continuous instillation of heparin through the arterial end of the shunt),

provides thermostating of the metal block holding the electrode by the blood itself, and secures a well-defined flow situation which excludes artifacts due to the remaining flow

sensitivity of the electrode. Further advantages of this shunt arrangement are that calibration can be checked at every desired moment, the transducer can be replaced at any moment if

•

necessary, and other parameters may be measured simultaneously. Cathode and Teflon membrane

are sterilized at 55 °C with ethylene oxide, the Silastic shunt tube by germicide. The

system, as tested in vitro, is stable (variation less than 0.5 percent in 24 hours), linear

and precise (± 0.2 percent) in a broad range of oxygen pressures (from about 10 to more than

700 mm Hg); its response time for 95 percent deflection is 0.4 second.

Continuous recording of Pa02 was achieved for periods of 6 to 24 hours in more than 100

patients, mainly with respiratory failure. The PO2 of blood samples taken from the shunt of

each patient was compared with the continuous Pa02 readings. The values tallied well up to 110 mm Hg. A systematically lower PO2 measured in the sample from 110 mm Hg upwards is likely due to leakage from the syringe and oxygen consumption by the erythrocytes during

transport and analysis of the blood samples. This system may be reliably used for monitoring

Pa02 in patients up to 24 hours without occurrence of any cutaneous or vascular reactions.

3. New Miniature PO2 Electrode

Our new electrode has a lay-out similar to our previous design [5] and also a response

time for 95 percent deflection of 0.4 second, but a diameter decreased from 2 to 1.2 mm

[7,8]. The cathode again is a platinum ring of 0.3 mm diameter and approximately 3 ym

thickness which makes possible a very low flow dependency in the presence of a relatively

high current output. In order to guarantee a fast sufficient electrolyte exchange in the

space between the membrane and the cathode area (which is important for the stability and

small CO2 sensitivity of the electrode) the ratio of the unsupported to total membrane area

had to be increased from 4 to 6 percent. Since this might lead to an increased sensitivity

to pressure, the static and dynamic sensitivity to pressure was investigated with particular

care. The most important characteristics of the probe are (with a 6 ym Teflon membrane at

37 °C):

sensitivity = about 0.6 nA/mm Hg,

linearity = ± 0.2% (from 10 to 760 mm Hg),

reproducibility = better than 0.1 mm Hg,

short-term stability (2 h) = ± 0.25%,

long-term stability (24 h) = ± 1%,

response time for 95% deflection = 0.4 s, zero current < 15 mm Hg (drift 0.5 mm Hg/24 h),

flow dependency (7-100 cm/s) = ± 1%,

temperature dependency (nonlinear) = 2-3%/l °C,

sensitivity to physiological CO2 < 2 mm Hg,

sensitivity to pressure: static ^ 0.2%,

dynamic < 1%,

sensitivity to acceleration (0 to 10 g) < 0.1%.

236
This electrode so far has been applied to physiological studies in the dog and cat as well as to continuous monitoring of human Pa02 in the radial artery where it was introduced

through a Teflon needle. Figure 1 shows a tracing in a cat. In the animal experiments

where the catheter tip was protected by silicone oil, no blood coagulation occurred at the

tip for several hours.

Ventilatory

Frequency

APaQ (torr)

40/

[10

Respiratory p500

% (torr) 0

Respiratory

'C02

r60

(torr) -

Integrated

Phrenic

Nerve

Activity

Systemic

Blood

Pressure (torr)

r200

mi n

; 1 i 1

! p

-—

i

1 i.o -h-hT

-H

i

1

• Tt

! -4

-J

1

. .

20/ .

b

^0/min

fmmmmi mMmmm smb?mm

10 s

Figure 1. Recording of arterial oxygen pressure (PaOa; top tracing) in the

carotid artery of an artificially ventilated cat during an experiment con- cerning the regulation of breathing. The respiratory frequency was changed

from 40/min to 10/min, while end-expiratory PO2 and PCO2 were kept constant

by adjusting the inspiratory concentrations. Note the respiratory fluctuations of PaOa (APaOa), increasing with decreasing respiratoryfrequency,

which can be followed faithfully by these fast PO2 electrodes (for proof,

see reference 4)

4. References

[1] Schuler, R. and Kreuzer, F., Rapid polarographic in vivo oxygen catheter electrodes,

Eespir. Physiol. 3, 90 (1967).

[2] Yokota, H., Respiratory fluctuations of oxygen pressure in alveolar gas and blood of the

dog, thesis, Nijmegen, 94 pp. (1973).

237
Kreuzer, F. , Respiratory fluctuations of oxygen pressure in alveolar air and arterial

blood, in Oxygen Measurements in Biology and Medicine, J. P. Payne and D. W. Hill, eds.

p. 139 (Butterworths , London and Boston, 1975).

Kreuzer, F. , Transmission of alveolar oxygen pressure oscillations, in Morphology and

Mechanisms of Chemoreceptors , A. S. Paintal , ed., p. 176 (Navchetan Press, 1976).

Kimmich, H. P., and Kreuzer, F. , Catheter PO2 electrode with low flow dependency and

fast response, in Oxygen Pressure Recording in Gases, Fluids, and Tissues, F. Kreuzer

ed.; Progr. Resp. Res. 3^, H. Herzog, ed., p. 100-110, (Karger, Basel, 1969).

Jank, K. , Hemptinne, J., Swietochowski , A., and Demeester, M., Continuous in vivo

measurement of arterial PO2 in humans, J. Appl. Physiol. 38, 730 (1975).

Kimmich, H. P., Kreuzer, F., Spaan, J. G., Jank, K. , de Hemptinne, J, and Demeester, M.,

Monitoring of PO2 in human blood, in Oxygen Transport to Tissue - II. Adv. Exper. Med.

Biol., Vol. 75, J. Grote, D. Reneau and G. Thews, eds., p. 33 (Plenum Press, New York,

1976).

Kreuzer, F. and Kimmich, H. P., Recent developments in oxygen polarography as applied

to physiology, in Measurement of Oxygen, H. Degn , I. Balslev and R. Brook, eds., p. 123

(Elsevier, Amsterdam, 1976).

238
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Gaithersburg, Maryland, July 7-8, 1975. Issued June 1977.

ALTERNATIVE METHODS OF CO2 MEASUREMENT, WITH PARTICULAR

REFERENCE TO CONTINUOUS RECORDING

L. H. J. van Kempen and F. Kreuzer

Department of Physiology

University of Nijmegen

Geert Grooteplein Noord 21a

Nijmegen, The Netherlands

Continuous in vivo recording of CO2 pressure (PCO2) in the gas phase has long been common but continuous in vivo monitoring in fluids, particularly blood, has been impeded by

the relatively large dimensions and the slow response of the available devices. A method

of measuring PCO2 in liquids accurately and rapidly by a miniaturized instrument analogous

to the oxygen electrodes developed by Kreuzer and his coworkers [1-3]^ presumably might help

to solve many problems in physiology and reaction kinetics, as well as in clinical medicine.

The conventional CO2 electrode according to Stow et al. [4,5] has been widely used but

remains limited to a response time of the order of one minute in most cases, particularly

with miniaturized types which furthermore are apt to be less stable and have a shorter life

time. The recently introduced General Electric catheter CO2 electrode [6], e.g.^ still has

a response time of one minute. A recent calorimetric CO2 electrode [7] also has a response

time of 2 minutes for 99 percent deflection. We therefore had decided to investigate the

possibility of applying other principles to the measurement of CO2. In the present paper, we will discuss the results of our experience with the quinhydrone electrode and with the

conductivity electrode.

1. The Quinhydrone Electrode

The quinhydrone system was previously applied to pH measurement and has now been

modified for the determination of PCO2 [8]. The principle resides on the oxidation and

reduction of hydroquinone and qui none, respectively; the oxidation-reduction reaction and

the corresponding electrode potential are described by:

Q + 2H'^ + 2e"

^

HgQ

2

F -

t - FL„ ++

, ^^H^^ - r . RT , ^ RT , , ^ In -— = E„ + or In -^^^— + —f In an+

0 2F a^^Q 0 2F a^^Q F H

where Q = quinone; aq, s,^ g, a^+ = ionic activities of quinone, hydroquinone and hydrogen

ions, respectively; E = electrode potential; E = standard electrode potential = 700 mV at

25 °C; other symbols as usual.

If the activities of quinone and hydroquinone are the same, the electrode potential

becomes independent of any of these species and is a function of the hydrogen ion activity

only. This condition prevails when using quinhydrone, an equimolar compound of quinone and

hydroquinone.

Figures in brackets indicate the literature references at the end of this paper.

(1)

(2)

239
A membrane- covered electrode was constructed with either a calomel or a silver-silver

chloride electrode as a reference, either of which is suited. Figure 1 shows the con- struction of this electrode. A platinum ring (o.d. 4.9, i.d. 4.5 mm) was mounted around

the liquid junction of a calomel electrode, or a platinum plate (surface area 8 mm^)

together with a silver ring (o.d. 6.8, i.d. 5.8 mm) was melted into glass (not shown here).

A membrane (Silastic 25 ym or Teflon 6 ym) was mounted and the electrode housing was filled

with a solution_of 2 ml quinhydrone (10"^ mol/1) and KCl (0.1 mol/1) with or without sodium

bicarbonate (10'^ mol/1). Lens paper was used as a spacer to stabilize the fluid layer.

1 ucite support

rubber ring

covering KCI inlet

tefl on brass ring

calomel electrode1 uci te stainless steel

glass chamber —

circulating water —

0-ring supporting

the membrane

rubber ring

|-1 ucite

teflon membrane [ P]atin

I

I ring

I lenspaper —

calomel liquid junction

1 ucite housing

supporting the

membrane

.platinum

ring

cal omel

liquid junction

Figure 1. Construction of C02-quinhydrone electrode [8]

(with permission of Respir. Physiol.).

The pK of hydroquinone is 10.32 ± 0.08. The plot of E vs pH is linear between pH 4 and 9, with a slope -AE/ApH = 58.2 mV/pH at 25 °C, the same with or without CO2. The plot

of E vs log PCO2 with 0.1 mol/1 KCl is linear between PCO2I5 and 80 mm Hg, with a slope

-AE/Alog PCO2 = 55.8 mV/log PCO2 at 25 °C in the presence of bicarbonate. This results in

a -ApH/Alog PCO2 = 0.96, in good agreement with the conventional CO2 electrode.

A number of factors such as chloride concentration, oxygen concentration, light, time,

acid-base reactions of hydroquinone and quinone, semiquinones , autoxidation , and salts

influence the performance of this electrode, but it could be shown that their effect in 0.1 mol/1 KCl is negligible in physiological conditions [9].

The most important characteristics of performance of this electrode, to be compared

with the conventional CO2 electrode, are listed in table 1. Calibration lines are shown in

figure 2 in bovine blood at 37 °C with or without bicarbonate added to the electrode solution.

240
Table 1. Synopsis of properties of various CO^ electrodes

for continuous monitoring in -vivo.

General Electric

Property Stow

electrode

catheter probe

(Stow principle)

Quinhydrone electrode

macro micro

Conductivity

electrode

Cal or i metric

electrode

Accuracy 2% ±2-3 mm Hq similar to Stow electrode <2 mm Hg SE = ±5 mm Hg

Cal ibration 1 inear 1 inear 1 inear 1 inear 1 inear 1 inear

Response

time

1 min

for '99. 5%

67 s for 90%

{^'tn vivo)

2 min

for 95% on

3 min

for 95% off

1 min

for 95%

4-7 s for 90%

2 min

for 99%

Stability 0.2 mV/h ±3 mm Hg

during 6 h <0.2 mV/h 4 mV/h

(platinized)

+5 mm Hg

during 1 week

"NaHCOg

(mV)

-IOOt

-80-

60-

40-

20

0.025 10

'NaHCOc

^HoO

57.2

29.7

—r—

10 20

^H20

(mV)

r+20

-+40

—I 1 T

1

40 60 80 100

+60

-+80

h+100

'^C02(mmHg)

Figure 2. Calibration lines of COa-quinhydrone electrode with

bovine blood equilibrated with various CO2 concentrations at

37 °C with or without bicarbonate added to electrode solution

[8] (with permission of Respir. Physiol.).

The performance of this electrode is fully comparable to that of the conventional CO2

electrode. Compared with the latter, it has the advantage of being an inexpensive lowimpedance electrode which may be miniaturized more easily and therefore might be adapted to in vivo PCO2 estimation.

241
2. A Single-Unit Electrode for CO2 and O2

We also miniaturized the C02-quinhydrone electrode and combined it with a polarographic

oxygen electrode of the CI ark- type, yielding a single-unit CO2-O2 sensing microelectrode

system. The design of this electrode is similar to that of the oxygen electrode constructed

by Schuler and Kreuzer [1] and Kimmich and Kreuzer [2]; a slightly larger version is shown

in figure 3 [10].

insul ation

silver-silver chloride electrode

I trovidur

si 1 ver support

catheter

lead

fluid compartment

Araldit

platinum electrode

silver ring

Silastic membrane

8. 7 mm 3.5 mm

Figure 3. Construction of C02-quinhydrone microelectrode system [10]

(with permission of Respiv, Physiol.).

The cylindrical probe has an external diameter of 3.5 mm and a length of 8.7 mm. It consists of a central platinum electrode with a surface area of about 2.5 mm^ sealed with

Araldite into a holder of polyvinyl chloride (Trovidur) and surrounded by a tubular Ag-AgCl

reference electrode. The tip of the probe is covered by a thin membrane (Silastic 25 pm or Teflon 6 ym) fixed to the reference electrode by a silver ring. A quinhydrone solution of

10 3 mol/l in 0.1 mol/1 KCl with or without bicarbonate added, is introduced into the

probe.

The performance data of this probe used as a CO2 electrode are listed in table 1. Figure 4 shows linear calibration lines for blank and platinized platinum electrodes. The

line of the blank probe lies somewhat above that of the platinized probe. The slopes with

0.1 mol/1 KCl are similar and lower than in the macroelectrode described above. The

sensitivity -ApH/Alog PCO2 thus is about 0.42. Stability is better in the platinized probe

but considerably inferior to that of the macroelectrode described above. Response time,

however, is superior to that of the macroelectrode.

The properties of this probe used as an oxygen -electrode are fully comparable with

those of the oxygen sensors described previously by Kreuzer and his group, particularly

also concerning the fast response, the response time for 95 percent deflection being about

0.4 second at 20 °C, independent of the electrolyte composition. The polarogram is not affected by addition of quinhydrone with or without bicarbonate.

The responses of this unit used as a CO2 electrode to oxygen and used as an oxygen

electrode to CO2 are mutually independent in physiological conditions, at least as long as the succession is not too fast. In view of a response time for 95 percent deflection of

about one minute for CO2 and about 0.4 second for oxygen, it is possible to take successive

recordings of oxygen and CO2 in vivo after elapse of the respective response times.

There is a lower limit for the surface area of the electrode concerning CO2, being

reached in the blank platinum electrode at a surface area of about 1 mm^, which is important for further miniaturization. The platinum surface of this electrode is too large

for the determination of oxygen in fluids since the flow dependency is considerable. Any

compromise in this respect may be difficult without membranes combining high permeation

rate for CO2 (in order to get a sufficiently fast response) with . poor permeability to oxygen.

242
3. The Conductivity Electrode

In a further attempt to reduce the response time which is still too long in the types

of quinhydrone electrode just described, an entirely different principle was adopted in the

device to be presented now, in which the principle of establishing a CO2 equilibrium across

the electrode membrane has been abandoned. Electrolyte conductivity has proven to be a fast, sensitive and accurate method for the determination of CO2 in liquids. The microelectrode system presented here is a first attempt to apply conductometry to continuous

monitoring of PCO2 in vivo [11].

A stainless steel tip is attached to one end of an x-ray double-lumen polyethylene

catheter of 60 cm length, 2.7 mm external diameter and 0.8 mm lumen diameter. A Silastic

membrane of 25 pm is mounted on this tip by a device specially designed for this purpose

and fixed by a stainless steel ring. Each lumen is separately connected to a conductivity

cell located at some distance from the membrane. The whole electrode system is flushed

continuously with bidistilled water at a constant flow rate. Electrolyte conductivity of

the water is measured before entering the catheter and after having passed the membrane.

The difference in conductivity between inflowing and outflowing water is related to the

PCO2 of the medium to which the electrode is exposed.

The sensitivity of the electrode system, which is the change in specific conductivity

(in S'cm"^) of the carrier to a change in PCO2 of the medium, is determined, among other

factors, by the efficiency of CO2 transfer from the medium across the membrane to the

carrier. It depends on membrane permeability and surface area, contact time of the carrier

with the membrane (contact length/flow velocity), and carrier volume per unit membrane

surface area (carrier layer thickness). In this electrode prototype, a membrane surface

area of 3.5 mm^ and a layer thickness of 250 m were chosen. Carrier flow rate was within

a range of 1 to 7 ml/min, thus giving contact times between 0.05 and 0.007 second.

The choice of the conductivity cell depends on the specific conductivity to be measured.

Since at a specific conductivity of the order of yS-cm"^, variations of a few percent have

243
to be determined, a conductivity cell with a_1ow cell constant (= distance between the two plates/surface area of one cell plate, in cm"^) is required. The cell volume should be

close to that of the carrier in contact with the membrane in order to maintain the upper

limit of the frequency at which PCO2 fluctuations of the medium can be followed by the

electrode. A series of 6 types of cell were investigated from which the type shown in

figure 5 was most suited. It consists of two platinum wires (0.2 x 0.05 mm) coated by

Araldite and wound at a distance of 0.2 mm to a coil of 0.6 mm i.d. and a length of 8 mm

(the interior of the coil being stripped of insulating Araldite), fitting into the lumen of

the catheter tip. The main advantage of this location is the reduction of the response

time and the convenience of temperature control. Between uses, the cells were filled with

Hibitane 0.2 percent to prevent microbial contamination.

Figure 5'. Stainless steel conductivity electrode tip covered by a C02-permeable membrane provided with a spiral conductivity cell and mounted on

a flexible catheter [11] (with permission of Respir. Physiol.).

Electrolyte conductivity was determined with a commercially available direct-reading

conductivity meter (Radiometer COM 2, frequency 70 Hz, test voltage 0.25 V) connected to a two-channel recorder. Conductivity meter and recorder were calibrated by precision resistance

boxes. In this way, it was possible to estimate variations in specific conductivity of up

to 0.02 yS'cm"^. A further improvement was achieved by applying a conductivity bridge

which includes both measuring and reference cell. This makes it possible to directly

detect relative variations in impedance due to CO2, reduces the difficulties of eliminating

A-C interference (capacitance effects), and minimizes the influence of temperature changes.

A plot of the specific conductivity against PCO2 in water is curvilinear and convex to the abscissa. However, in the low range used in the operation of the CO2 electrode (specific

conduc tivity about 1-3 yS-cm"^) the curve is practically lineajr, as shown in figure 6. Position and slope of this straight line depend on the flow rate. Increasing flow rate

lowers the specific conductivity at a constant PCO2 as well as the sensitivity (slope).

The reduction of specific conductivity at PCO2 = 0 is due to the decreasing conductivity of

the carrier at higher flow rates, probably originating from boundary effects on the surface

of the conductivity cell. The sensitivity for CO2 is diminished at higher flow rates

because both contact time of the carrier with CO2 at the membrane and time available for

hydration of CO2 are shortened. The dependencies of sensitivity and response time on the

flow rate are both curvilinear and convex against the coordinates. Thus, high sensitivity

and fast response oppose each other. The cell type discussed here at a distance of 10 mm

from the membrane provides, when exposed to a step change of PCO2 between 0 and 5 percent

CO2 and with a flow rate of 0.7 ml/min at 18 °C, a response time of 4 seconds on and 7 seconds off. The sensitivity of this cell in gas or in a stirred liquid phase is similar.

The performance of this cell is again summarized in table 1.

Theoretically, the performance of this electrode will depend on the conditions of

membrane diffusion, convection, and kinetics of CO2. In practical terms, the conclusion is important that both flow rate and distance between membrane and cell determine the sensitivity

of the system. At constant flow rate, the sensitivity increases with this distance until

chemical equilibrium is reached. At constant distance and varying flow rate, the sensitivity

is determined by the initial CO2 concentration in the carrier which is inversely proportional

to the flow rate, and by the time required to attain a certain fraction of the equilibrium

244
"o 3.35 ml/min

0.55 ml/min

1 .75 ml/min

Figure 6. Specific conductivity

of the carrier at three different

flow rates (0.55, 1.75, and

1 .80- 3.35 ml/min) after exposure of

the electrode to a gas phase

sure at 10 °C [11] (with permission of Respir. Physiol.)

of varying CO2 partial pres20 40 60 100

(mmHg)

concentration and hence on the flow rate as well. When a fast response is required, the

conductivity cell should be located close to the membrane. If at the same time high

sensitivity and low CO2 consumption are needed (in order to reduce dependency on the flow

rate of the medium), the carrier flow rate may optimally be chosen in the range where

turnover of CO2 into bicarbonate is high. The flow dependency in liquid media might be

minimized by using membranes with smaller surface area and/or lower permeability to CO2.

The resulting reduction in sensitivity should then be improved by other means, e.g., by

replacement of bidistilled water by highly purified water (conductivity water) which would

increase carrier resistance and thus the relative change in specific conductivity. At

these lower conductivities, however, balancing of the bridge becomes more difficult which

might be partially avoided by using cells of a smaller cell constant. Along the lines of

this reasoning, a further miniaturization of the electrode might be expected to increase

sensitivity and reduce CO2 consumption (flow dependency) at the same time.

Finally, it might be mentioned that we have achieved an accurate estimation of total

CO2 in fluids of low CO2 concentration (0.05-0.5 mmol/1) where the conventional Van Slyke

technique with an accuracy of 0.1 mmol/1 is likely to produce errors ranging from 20 to 100

percent. The essence of this method [9,12,13] is the use of an infrared CO2 analyzer (URAS

2, Hartmann and Braun; cell of 250 mm length, corresponding to a full-scale deflection of

0.01 percent CO2) surrounded by nitrogen rather than by air as usual. The inlet and outlet

of the analyzer were connected to a gas extraction chamber provided with a sample inlet

system. All connections were made of CO2 impermeable Viton tubing. The circulating gas

flow of 120 ml/min was regulated by a membrane pump and a flowmeter. In order to expel all CO2 from the sample, 3 ml of 0.1 mol/1 lactic acid with antifoam were added to the extraction chamber. Before the injection of the sample, the circuit was flushed with C02-free

nitrogen and then closed by turning a stopcock. The analyzer was calibrated by mixing

various amounts of air (containing 0.03 percent CO2) with C02-free nitrogen in a gas mixing

pump. The calibration curve thus obtained is not linear. At these low CO2 concentrations,

water vapor in the closed circuit interferes with the CO2 readings. This effect of cross- sensitivity could be eliminated by inserting a suitable filter supplied by the firm. In

4. Accurate Estimation of Total CO2 in Fluids

245
this way, total CO2 concentrations of 0.1 mmol/1 can be estimated in a 0.5 ml sample with an error of less than 2 percent. By changing the closed circuit volume or the length of the

cell, CO2 concentrations ranging from 0.05 to 100 mmol/1 can be estimated without any con- siderable loss of accuracy. Thiele and van Kempen [14,15] applied this method to the mea- surement of CO2 release by human skin and were able to determine CO2 amounts of ppm/10 cm^

mi n

References

[I] Schuler, R. and Kreuzer, F., Rapid polarographic in vivo oxygen catheter electrodes,

Res-piv. Physiol. 3_. 90 (1967).

[2] Kimmich, H. P. and Kreuzer, F., Catheter PO2 electrode with low flow dependency and

fast response, in Oxygen Pressure Recording in Gases^ Fluids^ and Tissues, F. Kreuzer,

ed.; Progr. Resp. Res. 3, H. Herzog, ed., p. 100 (Karger, Basel, 1969).

[3] Kimmich, H. P., Spaan, J. G., Kreuzer, F. Jank, K. , de Hemptinne, J. and Demeester,

M., Monitoring of PO2 in human blood, in Oxygen transport toTissue - II. Adv. Exper

' Med. Biol. , Vol . 75, J. Grote. D. Reneau and G. Thews, eds. , p. 33 (Plenum Press, \ New York, 1976).

r ^

[4] Stow, R. W. and Randall, B. F., Electrical measurement of the PCO2 of blood, Amer.

J. Physio. 179, 678 (1954). {

[5] Stow, R. W., Baer, R. F., and Randall, B. F., Rapid measurement of the tension of I carbon dioxide in blood. Arch. Phys. Med. Rehab. 38, 646 (1957). I

Schulz, v., Erdmann, W., Ulmer, H.-V., Kunke, S. , Baum, P., and Frey, R., Zur kontinuierl icherl Messung des arteriellen Kohlensauredrucks mit Katheterelektroden und Moglichkeiten

ihres klinischen Einsatzes, Anaesthesist , 22, 416 (1972).

| Vurek, G. G. and Kolobow, T., Continuous blood-gas measurement during prolonged I extracorporeal circulation, in Oxygen Measurements in Biology and Medicine, J. P. Payne ! and D. W. Hill, eds., p. 285 (Butterworths, London and Boston, 1975). i

i van Kempen, L. H. J., Deurenberg, H., and Kreuzer, F., The C02-quinhydrone electrode.

A new method to measure partial CO2 pressure in gases and liquids, Respir. Physiol. \ 14, 366 (1972). I

[9] van Kempen, L. H. J., Estimation of Free and Hemoglobin-Bound CO2, Thesis, Nijmegen, i 100 pages (1972). i

i [10] van Kempen, L. H. J. and Kreuzer, F., A single-unit carbon dioxide-oxygen sensing \ microelectrode system, Respir. Physiol. 23_, 371 (1975).

1

I [II] van Kempen, L. H. J. and Kreuzer, F., The CO? conductivity electrode, a fast responding I CO2 microelectrode, Respir. Physiol. 2A_, 89 (1975). '

'

[12] Ortega, F. G., Orie, S. A. M., and Tammeling, G. J., Determination of carbon dioxide

j content of blood by infrared analysis, J. Appl. Physiol. 21_, 1377 (1966). I

[13] van Kempen, L. H. J. and Kreuzer, F., Accurate estimation of low CO2 content of fluids

by infrared analysis, in Oxygen Affinity of Hemoglobin and Red Cell Acid-Base

| Status, M. R(iirth and P. Astrup, eds., p. 247 (Munksgaard, Copenhagen, 1972). I

[14] Thiele, F. A. J. and van Kempen, L. H. J., Die Porositat der menschlichen Haut: CO2- i Abgabe durch Haut oder Schweissdriisen? (Al kal ineutral isation) , Fette, Seifen ' Anstrichmittel, 2_, 190 (1971 ).

r [15] Thiele, F. A. J. and van Kempen, L. H. J., A micro method for measuring the carbon dioxide

release by small skin areas, Brit. J. Dermatol. 86_, 463 (1972). i

I 246

I
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

NBS STANDARDS FOR pH AND ION ACTIVITY

MEASUREMENTS IN BIOLOGICAL FLUIDS

Richard A. Durst

Analytical Chemistry Division

National Bureau of Standards

Washington, D.C., 20234, USA

and

Roger G. Bates

Department of Chemistry

University of Florida

Gainesville, Florida, 32611, USA

1 . pH Standards

It has been about fifteen years since the National Bureau of Standards, in response to the needs of clinical chemists and physiologists, certified a special composition of the

phosphate pH buffer for use in the biological range of interest between pH 7.3 and 7.5 [1]^.

This buffer, with a ph value of 7.392 at 37 °C, has been widely accepted and used in clinical

laboratories as the primary pH standard. The proximity of the pH value of this standard to that of blood minimizes errors due to faulty response of the glass electrode or nonlinearity

in the meter response. Of course, calibration with a second pH standard such as the 1:1 phosphate buffer (pH = 6.839 at 37 °C) is necessary to detect gross deviations from Nernstian

behavior which would necessitate replacement of the glass electrode.

In 1972, NBS reported the availability of a new buffer for use in the physiologic pH

range (pH = 7.382 at 37 °C) [2]. This buffer, tris(hydroxymethyl )aminomethane ("Tris") and

its hydrochloride salt mixed in a ratio of 1:3 at a tris-HCl molality of 0.05, was selected

because it exhibited several advantages over the phosphate buffer. These included stability

and compatibility with biological fluids and a temperature coefficient which more closely

approximates that of whole blood than does the phosphate buffer [2]. Since preliminary

experiments with cells containing a saturated KCl bridge solution and liquid junction

indicated residual liquid junction errors (i.e.j differences between the operational pH and

the corresponding pa^ values) of up to 0.041 pH unit at 25 °C when compared to the 1:1 phosphate buffer [3], the original Tris buffer was issued on a "provisional" certificate

until consistency with the pH scale could be checked more thoroughly. This buffer should be

considered a secondary standard until more extensive studies of the liquid junction problem

are completed.

Ladenson, et at. have also noted an inconsistency in the operational pH values when

Tris buffers are measured on pH/blood gas analyzers which had been standardized with the

phosphate buffers [4]. While their data indicate liquid junction problems with these

instruments, the interpretation of their findings is not unequivocal, since they also find

pH errors when a physiologic phosphate buffer is measured. The pH errors with the Tris

buffers are somewhat larger but also exhibit both positive and negative deviations from the

expected values. When directly compared to the NBS phosphate buffer, the Tris buffer gave

a pH value which was about 0.012 unit below the assigned value. While the Tris buffer may

not be suitable for use as a primary standard, Ladenson, et at. found it is eminently

suited for comparative measurement with blood pH systems. It also proved to be convenient

and effective in monitoring the performance of these analyzers.

Figures in brackets indicate literature references at the end of this paper.

247
In order to better approximate the colligative properties of blood plasma and serum, preliminary investigations have been made into the certification of the Tris buffer in

isotonic saline solution [2]. While it was hoped that this expedient would reduce the

uncompensated residual liquid junction potential, initial studies at NBS [3] again indicate

operational pH values that are low by about 0.05 pH unit at 25 °C when measured in a cell

with a saturated KCl liquid junction. It is obvious that much more extensive research must

be carried out to elucidate the nature and magnitude of this problem.

Whereas the Tris buffer may eventually prove to be unsuitable as a pH standard, there

are many other possible materials for use as clinical pH buffers, such as the materials in

the list published in 1966 by Good, et at. [5]. One of these, tris(hydroxymethyl )methylglycine

("Tricine"), was the subject of a recent study by Bates and coworkers [6] who found that a solution consisting of 0.06 m Tricine + 0.02 m sodium Tricinate (pa^ = 7.407 at 37 °C) had

a residual liquid junction error of about 0.01 pH unit which would not exclude the Tricine

buffer from consideration as a primary standard for pH measurements. Studies in isotonic

saline solution of two other zwitterionic buffers, HEPES [N-2-hydroxyethylpiperazine-N'-2-

ethanesulfonic acid] and TES [N-tris(hydroxymethyl)methyl-2-aminoethanesulfonic acid], with

pH values between 7 and 8 have recently been completed by C. A. Vega at the University of

Florida.

2. Ion Activity Standards

Ion-selective electrodes, like the pH glass electrode, are finding growing use in

various biomedical studies for determining or monitoring the activities of ions such as

Ca'^''', Na+, K"*", Cl~, F", etc. One very important area of development is the use of these

electrodes in automatic, multi -electrolyte clinical analyzers. As in the case of the pH

electrode, standards are also required for the reliable application of these sensors [7].

Investigations at NBS have been primarily concerned with aqueous standards and will progress

in the future to mixed electrolyte systems and synthetic biological fluids. To date, NBS

has certified standards for Na"*", K"^, CI", and F", and a standard for Ca"*"*" is in progress.

In considering the certification of ion activity standards, one important difference

exists in regard to the measurement process. For pH certification, an electrode system of

proven reversibility exists, i.e., the hydrogen gas electrode/silver-silver chloride electrode,

whereas in the case of ion-selective electrodes, the ideality of behavior has not been

unequivocably demonstrated. This consideration requires that more reliance be placed on the

activity scales and theoretical conventions used in calculating the single-ion activities.

The NBS has adopted the convention for single ionic activities proposed by Bates, et at. [8]

which is based on the hydration theory of Stokes and Robinson [9]. The details of this ion

activity scale have been reviewed by Bates [10] and will not be repeated here. Suffice it to say that the single ion activity coefficients are calculated from experimental values of

the mean molal activity coefficient (yMx)» the osmotic coefficient (0), and a hydration

number (h) according to the equations for univalent electrolytes:

log = log

Y[vix ^ 0-00782 (h^ - \\^)m 0

and

log Yx" = log Ymx + 0.00782 (h^ - \)m 0.

As in the case of the pH scale, it is necessary to make certain extra-thermodynamic assumptions.

The "conventional" step in the hydration treatment is the assignment of hydration numbers to individual ionic species. The hydration numbers are calculated from smoothed values of Y|^x

and 0 using the Stokes-Robinson hydration theory. The hydration number for the electrolyte

is then "split" between the cation and anion by assigning a hydration value to one of the

ions. For the present scale, the hydration number for the chloride ion was taken to be zero, and the other ionic hydration numbers are referred to this convention [10]. Ionic

activity for unassociated electrolytes derived on the basis of this hydration convention

have been found to be consistent with the observed responses of ion-selective electrodes.

248
Whereas the ion activity standards certified at NBS to date have been single salts intended for use in relatively simple electrolyte solutions, such standards are only reliable

over a concentration range of three or four decades, due to ion association problems at higher concentrations and contamination and/or adsorption problems at lower concentrations.

One way to avoid this latter difficulty is to prepare ionic activity buffers which can extend the low activity limit by several orders of magnitude.

Early studies at NBS demonstrated that the linear response of the silver ion-selective

electrode could be extended down to approximately 10"^.^ mol/1 with solutions consisting of

mixed silver halides and sulfide [11]. Not only was the response linear and approximately

Nernstian over 25 decades of concentration, but whereas the response time of the electrode

in unbuffered silver ion solutions increased with decreasing concentration, in the silverbuffered solutions, the electrode response was almost immediate, i.e.j less than 15 seconds,

even at the lowest levels of silver ion activity.

While the concept of "metal buffers" is not new, their use in the calibration of ionselective electrodes has been studied primarily by workers in Denmark and Hungary [12-17].

In Denmark, Blum and Fog developed water-soluble buffers for copper which avoid the slowness

of solution equilibration inherent in precipitate- type buffers which involve heterogeneous

equilibria [12]. This concept has been extended by Hansen, Ruzicka and coworkers [13-15].

They calibrated ion-selective electrodes using buffers for CU++, Cd"*"^, and Ca''"'' based on soluble EDTA and NTA complexes.

In Hungary, Havas, Kaszas, and Varsanyi have reported [16] the development of mixed

precipitate buffers for the halides (including fluoride) and silver ion. More recently.

Bailey and Pungor have suggested that, instead of serial dilutions of standards, the standards

in the range from 10"'+ to 10"^ mol/1 should be prepared by electrolytic generation of the

appropriate ion [17]. They demonstrated their technique with iodide and silver and also

indicated that sulfide, fluoride and thiocyanate could be generated electrochemically

Although the solutions are not buffered, an advantage of such a procedure is that it could

be automated and is more reproducible than the serial dilution method.

Great care must be exercised, however, in applying either the "pure" electrolyte or buffered standards to the calibration of ion-selective electrodes below the point where the

pure electrolyte response begins to deviate from linearity. The emf vs. ion activity response

will be entirely different for the two types of standards, and the interpretation of the

sample data will depend upon whether the sample system is buffered or not. It is clear that

much more work is required to develop ion activity standards which will be suitable for most

applications and acceptable to everyone.

3. Standards for the Future

As the requirements of clinical laboratories and available instrumentation become more

refined, data for salt effects and medium effects on electrolytes and gases are needed to provide the basis for the development and certification of the required standards and quality

control materials.

We plan to obtain data for ionic equilibrium processes in saline media, to develop

accurate methods for determining ion concentrations and gas tensions, and to develop multicomponent standards for calibration of the clinical instruments used for the determination

of pH, dissolved gases, and electrolytes in various biological fluids.

Ultimately such standards may contain all of the major ionic constituents, complexing

agents, and non-electrolytes at levels approximating those in the real samples. Thus, a synthetic blood standard would consist of an isotonic electrolyte solution containing, for

example, Na+, K*, Ca"*"'", Mg"*"^, O2, HCO3, CI", HPO^, and polyelectrolytes in physiological

amounts. Of course, the problems encountered in developing such a multi component standard

would be very great, but the availability of a single standard for the calibration of

instrumentation for the determination of multiple parameters is vitally needed for the

future development of this field. Only when such standards are available will it be feasible

to calibrate multi-parameter clinical instruments conveniently and with some assurance of

the consistency and comparability of data obtained in different laboratories. As far as the

major ionic constituents are concerned, the first step in this direction has already been

taken [18].

249
We are optimistic that given a high enough priority and sufficient time, the problems

will be solved. The continued support of organizations such as those sponsoring this workshop will help insure the success of this goal.

Addendum

NBS Facilities for pH Certification Measurements

Since there was no time during the Workshop on pH and Blood Gases for the participants

to tour the facilities used for electrolyte studies and pH certification, a section from a recent NBS Special Publication [19] on this subject is being reprinted here.

In 1972, the NBS emf measurement system was automated in order to increase the reliability

of the data acquisition and to reduce the time required for a pH certification run from

about five 8-hour days to two 24-hour days with minimal operator attention. The primary

purpose of this system is to cause a constant temperature water bath to cycle through a preprogrammed set of temperatures and to record the data when the bath temperature and

potentials read from the measurement cells pass certain stability and control requirements.

The temperature-controlled emf measurement system is diagrammed in figure 1. A centrally

located computer, a Univac Series 60, model 6135, is time-shared with about a dozen other

systems in the Analytical Chemistry Division. The data communications system consists of

four main parts: (1) the digital data bus, (2) the computer interface, (3) the laboratory

logic box interface, and (4) the laboratory control console. The computer interface connects

the computer to the laboratory via the digital data bus and laboratory interface. The

laboratory instrument console provides control and the capability of entering auxiliary

data for use by the computer to service and/or control the interfaced system. Two-party

line digital data busses have been installed in the Chemistry Building with outlets for each

Console Line Typewriter Printer

Temp Set Point

Selector

Teletype

UNIVAC Series 60

Model 6135 Computer

Computer Interface

Digital Data Bus

Lab Interface

Lab Control Console

Digital Barometer

Disk Mag Tape

Archive

Fluke 8400A DVM

Temp

Sensor

Refrigeration Unit

Proportional Temp

Controller

Standard

Resistor

Heater

Unit

Platinum Resistance Thermometer

Input

Multiplexer

Constant Temp

Water Bath

Electrochemical

Cells (Electrodes)

Electrometer Eppley

Standard Cells

Figure 1. Diagram of the temperature-controlled emf measurement system.

250
' laboratory. The communication between the laboratory and the computer interfaces is accomplished

by a multi-level time-sharing scheme. The laboratory interface contains the logic circuitry

j necessary to connect as many as four experimental systems to the digital data bus. Data

that are routinely collected by the system are stored on a random-access disc for a period

: of several days. If space on the disc becomes scarce, the data files are read onto a master archive magnetic tape for storage.

! In performing a pH certification, the first requirement is to prepare the buffer solutions

and measurement cells with utmost care. As specified in the pH certificates, the carefully

I dried and weighed salts are dissolved in distilled water of sufficient purity to have a conductivity of less than 2 x 10" 6 S'cm~i (ohm"i.cm"^) at 25 °C. For buffers in the neutral

and basic regions, carbon dioxide must also be removed from the water prior to dissolution

of the buffer salts. The hydrogen gas is purified of oxygen by passage through a catalytic

reduction tube or through a palladium purifier.

j The specially designed emf measurement cell shown in figure 2 consists of a hydrogen

i gas electrode compartment, a silver/silver chloride electrode compartment, and a series of

i gas-dispersion compartments for humidifying the incoming hydrogen gas by passage through the

I buffer solution. Details of the preparation of the platinized (or palladized) platinum and

the silver/silver chloride electrodes can be found in references [20] and [21].

Figure 2. Emf measurement cell used for the certification of pH buffer solutions.

251
The prepared measurement cells, usually consisting of two cells containing each of the

three levels of added chloride, are placed in the controlled temperature water bath as shown

in figure 3. Also shown in this figure are the coiled copper tubing hydrogen gas inlets and

the two platinum resistance thermometers.

Figure 3. Emf measurement cells in position in the control led-temperature water bath.

Figure 4. View of the measurement facility including the water bath, instrument rack and

laboratory logic box interface (open on rear wall).

252
The complete measurement system, with the laboratory logic box interface (open) on the

rear wall, is shown in figure 4. Another view, figure 5, shows the entire electronics rack

and the thermostated standard cells on the laboratory bench top. Briefly, the equipment in

the electronics rack consists of (from the top): A digital barometer which is automatically

Figure 5. Close-up view of the

instrument rack with thermostated standard cells on the

laboratory bench to the right.

monitored during the certification measurements is used to correct the partial pressure of

hydrogen for variations in the atmospheric pressure; the autoranging and autofunctioning

digital voltmeter which serves as the analog-to-digital converter for both resistance and

emf measurements; the proportional temperature controller which, in conjunction with the

temperature set point selector (next panel down) and the bath refrigeration unit, maintains

preset bath temperatures to better than ± 0.01 °C by means of immersion heaters; and the DVM

input multiplexer which, under computer control, switches the various measurement parameters

into the DVM for A-to-D conversion prior to transmission to the computer for acquisition and

storage. The functions multiplexed include: temperature (platinum resistance), standard

resistor (calibration), standard emf cells (calibration), and up to ten measurement cells.

The control console permits communication with the computer via a series of thumbwheel

switches which enter all of the required input parameters, e.g., temperature sequence,

number of cells, time delays, and program identification numbers. The control console also

includes a set of pushbuttons which are used to set up, start (send), and terminate the

experiment, as well as indicator lights which signal operations (data acquisition, operate,

auto, ete.) or problems (error, reject). A vibrating-reed electrometer is located below the

control console and, in conjunction with the three-position pH switch, provides a highimpedance input to the DVM for measurements with glass and high-resistance ion-selective

electrodes. The remainder of the equipment in the rack constitutes the Mueller bridge

system for manual checking of the water bath temperature via the second platinum resistance

thermometer shown in figure 3.

253
In operation, computer control begins by automaticany positioning the temperature set point selector at the starting sequence number. When the bath temperature achieves the

nominal value, within certain preset control limits as indicated by the platinum resistance

thermometer, the input multiplexer automatically switches through the readout positions.

After a programmed time delay, the multiplexer recycles through the positions and compares

the values to the previous sets of data. After three cycles, if the data agree within the

requisite control limits, the computer sets the temperature control to the next sequence

value and the operation is repeated. If any of the values exceed the control limits, an out-of-range message is printed next to the incorrect value and, following two more cycles,

the computer proceeds to the next sequence temperature after again flagging the out-of-range

value. Depending on the temperature value or the direction of the temperature change, the

computer turns on the refrigeration unit or an auxiliary heater. After the temperature

sequence is completed, the routine is automatically terminated.

Reduction of the acquired data is performed by batch operation on the NBS central

computer, a Univac 1108. The data reduction consists of a series of operations including

correcting the experimental emf values to the standard partial pressure of hydrogen, calculating

the acidity functions, extrapolating to zero molality of added chloride to obtain the limiting

acidity functions, evaluating the chloride ion activity coefficient using the Bates-Guggenheim

pH convention, and finally calculating the pa^ values. These experimental values are then

smoothed with respect to temperature by the method of least squares to give the certified

pH(S) values.

References

[I] Bower, V. E. , Paabo, M., and Bates, R. G., Standard for the measurement of the pH of

blood and other physiological media, J. Res. Nat. Bur. Stand. 65A, 267 (1961).

[2] Durst, R. A. and Staples, B. R. , Tris/Tris -HCl : A standard buffer for use in the

physiologic pH range, Clin. Chem. 18^, 206 (1972).

[3] Paabo, M., Acidimetric Measurements, Chapter 2 in Eleatroohemioal Analysis Section:

Summary of Activities July 1969 to June 1970, R. A. Durst, ed. , NBS Technical Note

543 (U. S. Government Printing Office, Washington, D. C, Nov. 1970).

[4] Ladenson, J. H., Smith, C. H. Dietzler, D. N., and Davis, J. E., Use of tris buffers

for quality control of blood pH, Clin. Chem. 20, 1337 (1974).

[5] Good, N. E., Winget, G. D. , Winter, W. , Connolly, T. N., Izawa, S., and Singh, R. M. M.,

Hydrogen ion buffers for biological research, Biochem. 5^, 467 (1966).

[6] Bates, R. G., Roy, R. N., and Robinson, R. A., Buffer standards of tris(hydroxymethyl

)

methyl gyl cine ("Tricine") for the physiological range pH 7.2 to 8.5, Anal. Chem. 45,

1663 (1973).

[7] Durst, R. A., Staples, B. R. , and Paabo, M., Activity standards for ion-selective

electrodes, in Biological Aspects of Electrochemistry, Experientia Suppl. 1_8, 275

(1971).

[8] Bates, R. G., Staples, B. R. , and Robinson, R. A., Ionic hydration and single ion

activities in unassociated chlorides at high ionic strengths. Anal. chem. 4^, 867

(1970).

[9] Stokes, R. H. and Robinson, R. A., J. Amer. Chem. Soc. 70, 1870 (1948).

[10] Bates, R. G. , Ion activity scales for use with selective ion-sensitive electrodes. Pure

Appl. Chem. 36^, 407 (1973).

[II] Durst, R. A., Analytical techniques and applications. Chapter 11 in lon-Seleotive

Electrodes, R. A. Durst ed., NBS Special Publication 314 (U. S. Government Printing

Office, Washington, D. C, 1969).

254
[12] Blum, R. and Fog, H. M., Metal buffers as standards in direct potentiometric determination

of metal ion activities, J. Eleotroanal. Chem. 34, 485 (1972).

[13] Hansen, E. H., Lamm, C. G., and Ruzicka, J., Selectrode - the universal ion-selective

solid-state electrode: Part II, Anal. Chim. Acta, 59, 403 (1972),

[14] Ruzicka, J, and Hansen, E, H., Selectrode - the universal ion-selective electrode: Part

IV, Anal. Chim. Acta, 63, 115 (1973).

[15] Ruzicka, J., Hansen, E. H. and Tjell , J. C. , Selectrode - the universal ion-selective

electrode: Part VI, Anal. Chim. Acta, 67, 155 (1973).

[16] Havas, J., Kaszas, M. , and Varsanyi , M. , Radelkis ion buffer solutions. Bung. Soi. Instr. 25, 23 (1972).

[17] Bailey, P. L. and Pungor, E., The calibration and response of ion-selective electrodes

at low concentrations of primary ions. Anal. Chim. Acta, 64, 423 (1973).

[18] Mohan, M. S. and Bates, R. G. , Calibration of ion-selective electrodes for use in

biological fluids, Clin. Chem. 21_, 864 (1975).

[19] Durst, R. A., Standard Beferenoe Materials: Standardization of pH Measurements, NBS

Special Publication 260-53 (U. S. Government Printing Office, Washington, D. C, Dec. 1975).

[20] Bates, R. G. , Determination of pH, Theory and Practice, 2nd edition (John Wiley and

Sons, New York, 1973).

[21] Ives, D. J. G. and Janz, G. J. eds.. Reference Electrodes (Academic Press, New York,

1961).

255

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

USE OF CARBON DIOXIDE- AND OXYGEN TONOMETERED PHOSPHATE-BICARBONATECHLORIDE-GLYCEROL-WATER MIXTURES FOR CALIBRATION AND CONTROL

OF pH, pCOz, AND pOz ELECTRODE SYSTEMS

A. H. J. Maas, A. H. Veefkind,! R. A. M. Van den Camp,

'A. J. Teunissen, A. B. T. J. Boink, and T. J. C. Ruigrok

Department of Cardiology and Thoracic Surgery

University Hospital

Cathari jnesingel 101 Utrecht, The Netherlands

The desire for simultaneous calibration of pH, pC02, and p02 electrode systems of

modern blood gas analyzers and the need for adequate quality control for this equipment

stimulated this investigation. The most important difficulties to overcome were: (1) establishing the pH, pCOa-relationship of the calibration solution, and (2) systematic bias between

the response of gas electrodes, in different media.

1. Calibration Solutions

In this study (for details see reference [1]^), we started from phosphate-bicarbonatechloride solutions equilibrated with carbon dioxide gas for which we were able to calculate

a pH, log pCOa-relationship based on the law of electroneutral ity and the definition of the

ionic strength, using the practical ionization coefficients of phosphoric acid from Bates [2],

and of carbonic acid from Maas [3]. For the NBS equimolal phosphate buffer (0.025 mol/1) to which NaHCOa and NaCl were added in a total concentration of 0.060 mol/1, we found in the

pH range 6-8, at 37 °C the following equation set:

0.185 - I

pH = 7.1834 - 1.34207 /f, + 0.84361^ + a - log ^ _ q 135

a

pH = 10.806 - 0.4662 /f^ + a - log pCOa + log (0.160 - + y) (2)

in which is the ionic strength on a concentration basis, y the added amount of NaHCOa

and a the factor wherein the deviation of the practical ionization coefficients defined

in other ionic and medium systems is discounted.

Subtracting eq (2) from eq (1) results, after rearrangement, in a relationship between

pC02 and I^:

.0 + (0-185 - I^)(0.160 - + y) (3.6226 + 0.8758/f - 0.84361 ). (3) pCU2 + X 10 a o (I - 0.135)

a

This means: by equilibration of buffer solutions of this kind, with gas mixtures of known

composition, that is to say a given pC02, the ionic strength I is fixed, and could be

estimated by eq. (3). In this manner, pH-log pC02 equil ibratiSn lines of buffer solutions

^Department of Medical Physics, Free University, Amsterdam, The Netherlands.

^Figures in brackets indicate literature references at the end of this paper.

257
with different sodium bicarbonate concentration can be predicted from eqs. (1) and (2), on the condition that a is known. The factor a could be determined from eq. (1) or (2)

by pH measurements of equilibrated buffer solutions at a given pC02. From figures 1 and

2, it is demonstrated evidently that introduction of an empirical factor a leads to a good

agreement.

6.6 6.8 7.0 7.2 7.4 7.6 7.8 8.0; pH

Figure 1. pH-log pC02 equilibration lines of four buffer solutions with 0, 15, 30 and 45

mmol/1 sodium bicarbonate, respectively, calculated by means of eqs. (1) and (2) without

correction factor a; the dots are the measured pH values of the same buffer solutions

equilibrated with gas mixtures of 2, 4, 8 and 15 percent CO2, respectively.

Figure 2. pH-log pC02 equilibration lines, differing from each other by 5 mmol of sodium

bicarbonate per liter, calculated from eqs. (1) and (2) by substituting a. The dots are

the same measuring points as in figure 1.

It is known from the literature [4] that whole blood and glycerol /water (3/7 by

volume) mixtures at the same oxygen tension produce nearly equal diffusion currents as a consequence of their similar viscosity. In order to minimize differences between blood

samples and calibration solutions, we added glycerol to the buffer solutions. In aqueous

solutions the correction term (a) was dependent on y:

a = 0.0331 + 1.513y - 8.44y2 - 118. 5y3. (4)

In the glycerol /water (3/7 by volume) solutions, the correction term (b) was dependent on

258
I as well, caused by the medium effect:

6 = (-5.667 + 290.999y - 3113. 55y2 - 4573y3)j

(5) + 0.8700 - 39.343y + 317.56y2 - 2459. 3y 3.

By means of eq. (5) substituted in eqs. (1) and (2) (read a = 3), pH-log pC02 equilibration

lines of glycerol /water (3/7 by volume) buffer solutions with different sodium bicarbonate

concentration can be predicted.

So we were able to prepare suitable calibration solutions of desired pH, pC02 and

p02 values, by equilibrating a solution of 25 mmol of the phosphates, 30 mmol sodium

bicarbonate and 30 mmol sodium chloride per liter glycerol/water (3/7 by volume) mixture

with 4 percent CO2 in air and 8 percent CO2 in nitrogen, respectively. Various pH, pC02,

and PO2 values, depending on the barometric pressures, are given in the next calibration

table 1.

Table 1. Calibration values obtained by equilibration of above mentioned

buffer solution gas mixtures 1 and 2 at 37 °C over the

barometric pressure (B) range 740-780 mmHg.

Gas mixture 1 Gas mixture 2 4% CO2 and 20.09% O2 8% CO2 and 0% O2

B pH PCO2 PO2 pH PCO2 PO2

mmHg mmHg mmHg mmHg mmHg

740 7 420 27.7 139.2 7 226 55.4 0.0

745 7 418 27.9 140.2 7 224 55.8 0.0

750 7 416 28.1 141.2 7 222 56.2 0.0

755 7 414 28.3 142.2 7 220 56.6 0.0

760 7 412 28.5 143.2 7 218 57.0 0.0

765 7 410 28.7 144.3 7 216 57.4 0.0

770 7 408 28.9 145.3 7 214 57.8 0.0

775 7 406 29.1 146.3 7 212 58.2 0.0

780 7 404 29.3 147.3 7 210 58.6 0.0

The usefulness of this calibration buffer system was checked by measuring pH, pCOa

and PO2 values of tonometered whole blood on the Corning M 165 blood-gas analyzer, both by

the conventional calibration methods (buffer and gases) and by using the equilibrated

solutions. The measurements of pH and pC02 were at least comparable, whereas p02 measurements were substantially improved by minimizing the gas-blood difference using calibration

liquid of broadly the same viscosity as blood (see fig. 3).

The accuracy of this calibration method depends on the accuracy of the phosphate

buffers, against which the pH values of the calibration solution is operationally derived

(±0.005 pH), and the accuracy of the composition of the gas mixtures (±1% of the pC02

and PO2 value).

259
Figure 3. Relationship between measured pOz and calculated pOz- Continuous lines are the

calculated regression lines for y on x and the broken lines are the lines of identity;

(a) gas-calibration method; (b) buffer-calibration method.

2. Approach to Quality Control

The proposed buffer system not only opens new perspectives for simultaneous calibration

of pH, pC02, and pOa electrodes, but it also is a step in the direction of quality control,

as indicated by Noonan and Burnett [5]. Realization can be thought by the following line:

(a) Intralaboratory comparison of blood gas equipment can easily be realized by

offering samples from the same thermostated tonometer setup, to the different apparatus.

(b) Interlaboratory comparison would be possible in principal by sending around

ampoules containing a control solution for simultaneous control of pH, pC02, and p02.

In our routine laboratory, we have at our disposal three blood-gas apparatus: from

Radiometer, Models BMS2-MK2, and ABL-1; from Instrumentation Laboratory, Model IL-413.

For quality control of the BMS2-MK2, we measure a few times a week the pH of buffer solutions, marked A III and C III (composition: 0.025 mol/liter Na2HP0[,, 0.025 mol/liter

KH2PO1+, 0.030 mol/liter NaHCOs and 0.030 mol/liter NaCl using, respectively, water (A III)

and glycerol/water (3/7 by volume) mixture (C III) as solvent) tonometered with gas mixtures

of 4 percent CO2 and 8 percent CO2 in air. In figure 4, the fluctuation of the bias between the readings and the stated values over a two-month period is presented.

We also obtained initial experience with quality control of the ABL-1 and IL-413 with

the same buffer solutions equilibrated in a special setup consisting of one gas mixing

pump, delivering two gas mixtures, and a thermostated tonometer vessel. The results of

twelve measurements performed on various days are summarized in table 2.

The pH differences [ca. 0.012 pH) could be explained partly by a difference of 0.008

pH between the ABL-1 and IL-413 calibration buffers and the Radiometer precision buffer

(pH = 7.383), and partly by the type of salt bridge: ABL-1, 20 percent KCl ; IL-413,

saturated KCl with a membrane instead of an open junction.

The PCO2 differences (table 2) are equal for both instruments and do not deviate

significantly from the stated value for the lower pC02 value in contrast to the measured

high PCO2 values which are ca. 3-5 percent lower. Comparison of results obtained with

both instruments using tonometered blood samples revealed a relationship between the

measured and calculated pC02 values as presented in figures 5a and 5b. The regression

lines cross the identity lines near the lower calibration point. The main cause of deviation from identity seems due to the "memory effect" of pC02 electrodes as recently

described by Berkenbosch [6], Cramption Smith [7], and Hahn [8]. They show also that this

effect is reproducible and corrections can be made for it.

260
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261
BMS 73-131. AOO-5357-3690

0.01- A pH

+ o X / ° \ A HI 4 \

/

A > A /\ / ^A—

a

—A—A^

0.01-

0.01 V

c m 8

—I —I —I —i-i-i —I —I- 24 25 26 27 1 2 7 march I

-I —I —I —I —I —I —I —I — —I 9 11 15 16 17 22 24 25 28

april Ti —I —1 —1 —I —I -

1 6 12 13 U 15 mai

I —I —I —I 21 23 27

Figure 4. Tonometry results for pH, with BMS2-MK2 (see text). The mean differences and

standard deviations are: A III 4: 0.0054 ± 0.0045; A III 8: 0.0030 ± 0.0044; C III 4; 0.0026 ± 0.0046; C III 8: 0.0040 + 0.0049 which means that the precision is very high.

Figure 5. Relationship between pCOs of blood and gas: (a) measurements with ABL-1

;

y = 0.9115X + 2.1091; (b) measurements with IL-413; y = 0.8822x + 4.2134. Each dot

is the mean of ten measurements.

Concerning the p02 measurements (table 2), we found that the zero point could not adequately be established by a single filling of the cuvette probably caused by mixing

with contaminants remaining behind in the chamber after a measurement and oxygen loss from

the electrode. Values of aa. 30 mm Hg were read.

262
Using A III solution, the high p02 values agree with the stated value for the ABL-1

and were significantly lower for the IL-413; using C III solution, the p02 readings are lower for both instruments. Also we evaluated the p02 electrode systems of both instruments

using tonometered blood samples. The results are illustrated in figures 6a and 6b. There

is a tendency for the blood-gas regression lines to have slopes lower than those due to gas for both instruments. In the physiological range the difference between the value for

gas and tonometered blood is a few percent, but at higher oxygen tension readings are significantly too low. We feel that there is a need to improve the method of cali'^ration

of blood-gas analyzers.

7oo-p02 measured

mmHg

600.. HO- p02 measured

mmHg

120

p02 calculated mmHg

100 200 300 iOO 500 600 700 20 40 60 80 100 120 UO

700- p02 measured

mmHg

600-

500-

300

200

100-

UO- p02 measured

mmHg

120-

100

L-413

p02 calculated mmHg

-I- -I 100 200 300 400 500 600 700 20 40 60 80 100 120 140

Figure 6. Relationship between p02 of blood and gas: (a) measurements with ABL-1;

y = 0.9682 + 2.2981; (b) measurements with IL-413; y = 0.9804 - 1.0731. Each dot

is the mean of ten measurements.

263
Introductory in two trials, we have sent around 100 ml buffer solutions in plastic

bottles to approximately ten laboratories with the suggestion to equilibrate these solutions

with gas mixtures of CO2 content and to measure pH. So this method was restricted to setups of the indirect pC02 method according to Astrup and Siggaard-Andersen. In figure 7 the results of these trials are reflected.

0.03 ApH

0.02

0.01

0

"o.oi

0.02-

0.03

0.04

0.05

7.0

0.03- ApH

0.02

0.01

0

0.01

0.02-

0.03-

CIII - TRIAL I

7.1 7.2 7.3 7.4 7.5 pH 7.6

pH

CIII - TRIAL II

4^

7.2 7.3 7.4

0.03- ApH

0.02-

0.01

7.5 7.6

AIII - TRIAL II

.pH

+ ^0

0.01

0.02-

0.03

~' 1 7 ! 2 7 ! 3 7:4 7 7: 6 Figure 7. Two trial results of pH determined in CO2 equilibrated A III and C III buffer

solutions. Different symbols are from different laboratories. ^

264
In reference to these figures, we may make the following remarks:

(1) Rather good agreement {ca. 0.01 pH) in the physiological pH range is found by most

of the participants.

(2) Improvement of the results at the second trial, probably caused by experience and

sending aroun-d a precision phosphate buffer for calibration as well. Dramatic was the discovery of pH deviation of 0.011 unit between the precision buffer sent along, and a buffer

from the same firm and with the same batch number, present in one of the laboratories. Up to now the explanatory suggestions are still unproved.

(3) Differences up to 0.02 pH unit for equipment in the same laboratory. See the

closed and open squares in figure 7a.

(4) Possibility of signaling a wrong analysis of a gas cylinder. See open square at

the bottom of figure 7b.

(50 Possibility of detecting small difference in gas mixing pumps. See open and closed

triangles in figure 7b.

In an attempt to eliminate differences in the equilibration setup, and to simplify

the procedure, we are engaged in an investigation into the possibility of additives to the

buffer solution, which are able to absorb oxygen. To such a solution, with enhanced

capacity, the desired carbon dioxide and oxygen tension can be introduced. The only thing

to be done, prior to pH, pC02 and p02 measurement, is to bring the solution to 37 °C

before opening the sample holder. Preliminary results with a fl uorocarbon , because of its

high dissolving capacity for gases, showed usefulness for some equipment and uselessness

for others.

Most likely there will still be many obstacles on the way to quality control, however,

the purpose is worthwhile continuing this way.

References

[1] Veefkind, A. H., Van den Camp, R. A. M. , and Maas, A. H. J., Use of carbon dioxide - and

oxygen tonometered phosphate-bicarbonate-chloride-glycerol -water mixtures for calibration

and control of pH, pC02 and p02 electrode systems, Clin. Chem. 2J[, 685 (1975).

[2] Bates, R. G., and Acree, S. G. , pH of aqueous mixtures of potassium dihydrogen

phosphate and disodium hydrogen phosphate at 0 °C to 60 °C, J. Res. Nat. Bur. Stand.

34, 373 (1945).

[3] Maas, A. H. J., van Heyst, A. N. P., and Visser, B. F. , The determination of the true

equilibrium constant (pi^ig) and the practical equilibrium coefficient (pA'jg) for the

first ionization of carbonic acid in solutions of sodium bicarbonate, cereBrospinal

fluid, plasma and serum at 25 and 38 °C, Clin. Chim. Acta 33^, 325 (1971 ).

[4] Heitman, H., Buckles, R. G. , and Laver, M. B., Blood p02 measurements: performance of

micro-electrodes, Respir. Physiol. 3_, 380 (1967).

[5] Noonan, D. C. , and Burnett, R. W. , Quality-control system for blood pH and gas measurements, with use of a tonometered bicarbonate-chloride solution and duplicate samples of

whole blood, Clin. Chem. 20, 660 (1974).

[6] Berkenbosch, A., Measurement of carbon dioxide tension in poorly buffered solutions

with the pC02 electrode, Pflugers Arch. 318, 217 (1970).

[7] Crampton Smith, A., and Hahn, C. E. W., Studies with the "Severinghaus" pCOo electrode

I: Electrode stability; memory and S plots, Brit. J. Anaesth 47^, 553 (1975)'.

[8] Hahn, C. E. W. , and Crampton Smith, A., Studies with the "Severinghaus" pC02 electrode II: CO2 measurement using a single control analyzer, Brit. J. Anaesth 47, 559 (1975).

265

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

CALIBRATION OF BLOOD GAS ANALYZERS

Alan H. Runck

Corning Scientific Instruments

Medfield, MA 02052, USA

Blood gas measurements are unique in the field of medical laboratory analysis. They

are frequently requested at a point when the subject patient is in a life-threatening

situation which will be dealt with according to the results obtained from these measurements. The techniques that are used by experienced blood gas technicians for collection

and storage of samples and carrying out of measurements are complex but highly standardized:

anaerobic sample handling, "double insertion" of the sample into the cuvette, monitoring of

cuvette temperature and waiting for the proper end-point to be achieved are good examples

of the precise operator technique required. The equipment used by virtually every medical

laboratory to perform blood gas analysis is basically identical (glass electrode system for

pH, Severinghaus electrode for PCO2 and Clark electrode for PO2), with variations in the

volume of blood sample required for analysis, design of supporting hardware and degree of

automation. Yet, for all the similarities in the methodology for measuring blood gas, the means for calibration of blood gas analyzers are not uniform.

Classically, blood gas measurements have been calibrated using dilute aqueous phosphate

buffers for pH [1]^ and analyzed gases for PCO2 and PO2 [2]. Many modifications of calibrating materials however, have been advocated in the literature and are in limited

current use. For pH measurements, calibrating buffers containing added sodium chloride

(saline buffers) have been described [3], as well as bicarbonate-based buffers [4]. Gases

used for calibration of PCO2 and PO2 measurements are not of standard composition, but vary

from laboratory to laboratory, especially with respect to oxygen; some laboratories even

use a solution such as sodium sulphite [5] in lieu of an oxygen-free gas for zero calibration of PO2. The calibrating gases themselves may not, depending on the method of

formulation, contain the composition of gas that is stated on the accompanying label.

Finally, use of tonometered liquids for calibration of PCO2 and PO2 measurements has been

repeatedly reported [4,6,7] and is in routine use in a growing number of laboratories [3].

Why such diversity in calibration of instruments when basic design of the instrumentation is the same? The common reason in the literature seems to be: (1) a search for more reliable and accurate performance, and (2) a search for convenience. Modification of

pH calibrating buffers has occurred for both reasons. Use of saline buffers coupled with

use of a sodium chloride salt bridge is claimed to improve accuracy by reducing the sodium

response of different pH electrodes, and improve reliability by reducing depletion of the

salt bridge solution, as well as improving the stability of the liquid junction potential

[3]. Use of the bicarbonate buffering system is more convenient, since one aqueous solution

win serve as a pH, PCO2 and PO2 calibration point. Change in the manner of calibrating

PCO2 and PO2 measurements has been attempted for the purpose of improving accuracy and

reliability of these measurements. Differences in the PCO2 and PO2 readings between gases

and liquids having identical gas tensions have been reported; for this reason, use of

tonometered liquids having known values for PCO2 and PO2 has been investigated by others as

an alternative to calibration using gases alone.

A major area of controversy in blood gas has been the question of whether changing the

method of calibration will improve performance. PO2 measurements have been singled out as being most in need of improvement, because of the tendency of the PO2 sensor to read lower

on liquid than on gas samples. Maas and coworkers [8] summarized the concern about the

Figures in brackets indicate the literature references at the end of this paper.

267
accuracy of PO2 measurements by presenting data concerning the average PO2 readings obtained

when blood samples having known PO2 values were read on conventional blood gas instrumentation (table 1). The results showed a bias, the sign and magnitude of which were dependent

Tablp 1 Rlnnri/na*^ di'f'fprpnrp

Clin. Chim. Acta, 28,

inIII Pn^ rijp mpfl IflCUOUi <^ 1 irpmpnl" CMldIL 443 (1970)).

(?\fff^r ^ai Lci M;^;ic; riudO) MA H .1

, n. u. V\kz\ Lcllb 9 r . U • 9

O2 PO2 PO2 Correction

(%) (mm) (mm) (%)

5 JO JO . 0 - 1 . / 10 / u /I.I -1.0

20 140 132.8 5,1

40 280 269 3.9

60 420 399 5.0

80 560 521 7.0

on the level of PO2 in the sample being measured. Prior to Maas' studies, the bias was assumed to be a constant percentage of the sample PO2; it was thought to be caused by the

difference in diffusion coefficient in gas calibrant vs. liquid samples and, therefore,

able to be corrected for in the final result. Maas' data, however, refuted this.

Data by Hulands and coworkers [9] showed another aspect of PO2 performance that was equally disturbing (fig. 1). They showed that the magnitude of the PO2 bias varied not

1.20

_ 1.15 d

CO

+11 1. 10 J

T3O

m 1.05

1.00

14 18 22 26 30 2 6 10 14 18 22

June July

Figure 1. Day-to-day variation in PO2 bias (from Hulands et at. [9]).

only according to sample PO2, but also varied on a day-to-day basis that was unpredictable.

Hulands felt that this was caused by the gas/liquid correction factor for the PO2 electrode

changing from day to day; he proposed to remedy this by calibrating the electrode directly

with tonometered 30 percent glycerol solutions (fig. 2).

Our studies have centered on evaluating the clinical results obtained when gases vs. liquids are used for calibration of commercial blood gas analyzers. It seemed that one variable that was not able to be controlled in the studies of Maas and Hulands was that of

sample preparation; blood tonometry and sample handling at high levels of oxygen tension is

268
1.20

1.15

CD

O

1.05

1.00

t \ ^

\

^ r

Glycerol <^ = 0.897 (blood <^) + 0.115

r = 0.902

• Air o High PO2

X Low PO2

I

Figure 2. Paired measurements

of 30 percent glycerol

correction factor and whole

blood correction factor (from

Hulands et al. [9]).

.00 1.05 1. 10 1.15 Whole blood

.20

difficult. Therefore, we decided to look directly at differences between instruments

calibrated with gases and those calibrated with liquids. The two types of systems were set

up side-by-side in the Blood Gas Laboratory at Peter Bent Brigham Hospital and used to measure blood gas values on patients' samples. The results, shown in figure 3, indicate

that the variable bias described by both Maas and Hulands for PO2 is not observed. Comparison data, where both instruments were calibrated with tonometered liquids, are shown in

figure 4. The PCO2 measurements that were observed in this study are shown in figure 5. 500

X

E

£=a

8

COa

400

P 300

13

CD£

13

COO

cuE

cvj o

Q_

200

100

P02(gas col.) = 0.98xP02( liq.col.) S.D. = 5.7mnnHg

n = 55 samples

/

/

Instrunnent

Jo A

• B

A C

(June 73)

( Feb 74)

( Feb 74)

Figure 3. Comparison of gas

vs. liquid calibration for

blood PO2 measurement.

100 200 300 400 500

PO2 measurements using liquid calibration (mmHg)

269
500

X

E

E

<

CDE

3

c

X5O

q;CM o

a.

400

300

200

100

P02.A = 1.01xP02,B

S.D. = 12.5mnnHg

n = 99 samples

. /

100 200 300 400

PO2 reading, instrument B (mmHg)

Figure 4. Comparison of liquid

vs. liquid calibration for

blood PO2 measurement.

500

PCO2 (GAS CAL)-: 1,01 XPCO2 (LIQ CAU

REGRESSION STD. DEV. • 1.6mm

r = 0.99 2

n' 45

Figure 5. Comparison of gas vs. liquid calibration for blood

PCO2 measurement.

30 40 50 60 70 80

PCO^ reading, using liquid calibration

(mm Hg)

The data indicate that gas calibration of blood gas analyzers will yield PO2 and PCO2

values that are as accurate and reliable as those obtained when liquids are used for calibration. These results have been confirmed by more recent data in our laboratories on different types of equipment, and by Maas, who recently compared results obtained with

instruments calibrated with gases and instruments calibrated with tonometered glycerol

water solutions [4].

270
It is intriguing to ponder why earlier data showed such a pronounced variation in

results, both in terms of PO2 dependent bias and in terms of a day-to-day variation in

accuracy. We addressed this question by measuring the PO2 of tonometered water solutions

using a gas calibrated analyzer, and comparing these measurements with those obtained

indirectly from oxygen content determinations on the same samples, using the Lex-Oa-Con

oxygen analyzer. The results, shown in figure 6, are interesting: the oxygen content

500

400

X 300

E

E

g 200

100

0

• PO2 calculated from O2 content o PO2 measured

n 71

/^ § 7.6%

P 4.1 %

/

8.6%

1.0%

1

0 2 3 4

O2 content (m£/dl)

Figure 6. Comparison of measured PO2 of tonometered water vs. PO2 derived from oxygen

content measurements on the same samples.

measurements on the samples at all PO2 levels show that the tonometer performed as expected,

with each sample giving the anticipated PO2 value. However, PO2 measurments on the same samples, using a blood gas analyzer, show the very type of P02-dependent bias as was reported

by Maas. The conclusion from these data is that the variable bias reported by Maas is a function of atmospheric contamination, either during sample handling or instrument manipulation. Such contamination has been observed for every instrument studied and is manifested when recovery studies on tonometered blood or water are attempted. These conclusions

should not imply that a gas/liquid correction for PO2 measurements does not exist, but

rather, that the correction factor remains a constant over a period of time. In our experience,

the correction factor for present day blood gas analyzers is 5 percent, and varies no more

than ± 1 percent. What is most important, is for manufacturers of blood gas instrumentation

to explicitly state the accuracy of their instrumentation at all levels of PO2 and PCO2 so that users can effectively evaluate the performance of the total system without placing all

the blame for poor recovery of PO2 on a "high gas/liquid correction."

Calibration of blood gas instrumentation should be as simple and error free as possible,

without compromising performance. Our studies, and those reported by Maas [4], indicate

that gas calibration for PCO2 and PO2 measurements provides results which are comparable to those obtained when liquid calibration is used, when accuracy is used as the guideline.

References

[1] Bower, V. E., Paabo, M., and Bates, R. G., A standard for the measurement of the pH

of blood and other physiological media, J. Res. NBS, 65A , 267 (1961).

[2] Severinghaus, J. W., and Bradley, A. F., Electrodes for blood PO2 and PCO2 determinations, J. App. Physiol. 1_3, 515 (1958).

271
Drinker, P. A., Noonan, D. C, Ramanaiah, N., and Tole, J. R., Use of a sodium

chloride-phosphate buffer for pH standardization in a new blood gas analyzer with

an isotonic sodium chloride bridge, Clin. Chem. 1_9, 1243 (1973).

Veefkind, A. H., Van den Camp, R. A. M., and Maas, A. H. J., Use of carbon dioxide

and oxygen tonometered phosphate-bicarbonate-chloride-glycerol -water mixtures for

calibration and control of pH, PCO2 and PO2 electrode systems, Clin, Chem. 21_, 685 (1975).

Heitmann, H., Buckles, R. G., and Laver, M. B., Blood PO2 measurements: performance

of microelectrodes, Resp. Physiol. 3^, 380 (1967).

LeFevre, M. E., Calibration of dark oxygen for use in aqueous solutions, J. App.

Physiol. 26, 844 (1969).

Gleder, R. L., and Neville, J. F., Commentary on PO? and PCO2 electrode calibrations,

Amer. J. Clin. Path. 55, 325 (1971).

Maas, A. H. J., and Mertens, P. J., The measurement of the PCO2 and PO2 of blood with

electrodes in an open cuvette system, Clin. Chim. Acta, 28, 443 (1970).

Hulands, G. H., Nunn, J. F., and Paterson, G. M., Calibration of polarographic

electrodes with glycerol water mixtures, Brit. J. Anaesth. 42, 9 (1970).

272
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

QUALITY CONTROL

S. Raymond Gambino

Columbia Presbyterian Medical Center

New York, New York 10032, USA

In my discussion of instrument specifications I indicated that better quality control

(Q.C.) at individual sites was needed more than quality control of manufacturers. In fact,

I can state categorically that without some type of on-site Q.C. program it is impossible

to provide physicians with reliable blood gas measurements. Too many things can go wrong.

Attending physicians provide some degree of Q.C, but it is insufficient. If the data

agree with what the attending physician expects, then he or she is satisfied. If the data

provided by the blood gas instrument is not consi-tent with his or her expectations, then

the physician does not believe the data. In this situation, some type of readily available

referee sample is essential. Since blood gas measurements are made 24 hours a day, 7 days a week, the ideal Q.C. material would be available for immediate use at all times. Tonometered

blood is, therefore, inadequate as an "instant" control for peripheral and central blood

gas instruments around the clock. Tonometered blood, on the other hand, is the proper

sample for use as a secondary standard.

More than 10 years ago. Dr. Arthur Babson of General Diagnostics developed a formula

for a lyophilized serum control for acid-base assays. This lyophilized control is marketed

as "Versatol Acid-Base." In 1970, when we first began our high volume blood gas service, I received many complaints from physicians during the first month of operation. At that time we did not use controls on every shift. The supervisor of the chemistry service, Ms. I. Fonseca, suggested that we run the lyophilized controls on every shift Qvery day of the

week. Once we began to utilize controls, at 3 different levels on every shift, the number

of complaints fell to nearly zero. Now problems are discovered in the laboratory before

results were reported. But the lyophilized controls are not practical for the distant

units operated by non- laboratory personnel. Furthermore, the lyophilized controls lack

assigned values for P02, although a single, but somewhat variable, P02 value for all 3 controls can be obtained by assay.

Following the introduction of our first peripheral blood gas unit it became obvious

that a lyophilized control, or a whole blood tonometered control, was not adequate. What

was required was an ampouled single-use control. Such a control has just been developed by

Dr. James Turner and Dr. Arthur Babson of General Diagnostics. It is marketed as "Blood

G.A.S. Control." Pilot lots of the ampouled controls have been available to investigators

for more than a year. In fact, it was the availability of these unit-use controls for pH,

PCO2, and P02, at 3 different levels, that permitted us to expand our peripheral blood gas

units in 1975.

The new ampouled controls have revolutionized our operation. We are able to leave a peripheral unit unattended and unvisited by any laboratory personnel for several days as long as ampouled controls are available at the peripheral site.

What happens when we open an ampoule? Table 1 shows that in 4 minutes there is little

change in pH or PCO2. The P02 data in this experiment cannot be interpreted since the P02

in the ampoule is so close to the P02 of room air.

How reproducible are the values for pH, PCO2 and P02 when repeat measurements are made

on 2 different instruments that are calibrated with the same gases and buffers? Table 2 shows the data for pH. The pH values are very reproducible. Table 3 shows the data for

PCO2. The PCO2 is reproducible, but there is an obvious systematic difference between the

two instruments even though they were calibrated with the same gases. Table 4 shows the data

for P02. P02 is the least reproducible of the measurements at low values. This has to do

with the particular instrument used. The IL 313 introduces a bolus of room air in front of

273
Table 1. Effect of exposure to air, mean of 6 ampoules in each

of 3 different laboratories.

Table 2. pH comparison of two IL 313'

s

(same buffer)

.

Time: zero 30 s 60 s 240 s

pH 7.628 7,,629 7.,630 7.634

PC02 66.0 66..2 65.,9 65.7

PO2 163.5 163.,8 164,.5 163.2

Ampoule No. 1 No. 2

1

1

-7/

.

/1 00 -7/

.

/inn

, 4 1 y

2 7.,421 7.,416

3 7,,419 7..420

4 7..417 7..419

5 7,.417 7,.424

Mean 7,.419 7,.420

Table 3. PCO2 comparison of two IL 313'

s

(same gas).

Ampoule No. 1 No. 2

1 61 58

2 62 58

3 61 58

4 61 ' 58

5 61 59

Mean 61.2 58.2

Table 4. P02 comparison of two (same gas)

.

IL 313'

Ampoule No. 1 No. 2

1 60 62

2 59 61

3 59 63

4 59 60

5 58 59

Mean 59 61

each sample. The controls do not have any buffering capacity for P02, therefore they are quite sensitive to the presence of oxygen in the instrument at any pressure other than the

pressure in the control.

Are these controls useful when used routinely? Decidedly yes. Table 5 shows typical

actual problems revealed by the control and not revealed by standard gas and buffer calibration.

Table 5. Typical problems revealed by controls.

Expected Found Problem

PCO2 42 46 Needed new membrane

PCO2 58 62 Needed slope adjustment

P02 100 80 Mold in wash solution

P02 100 Slow Bubble behind membrane

PCO2 58 Slow Space between electrode

and membrane too thick

274
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

QUALITY CONTROL AND STANDARDS

Daniel C. Noonan

Port Huron Hospital

Port Huron, Michigan, 48060, USA

We have heard many different points of view on many aspects of blood gas measurements

at this workshop. There is one point, however, upon which we all probably agree. We all

feel that our own respective laboratories provide blood gas results which are of the highest

quality available with modern technology. While this is undoubtedly true, if we were to survey our laboratories, we would find a wide range of methods for the basic standardization

and quality control of these measurements in daily use.

This workshop should have several objectives in this area of standards and controls.

First, we should strive to identify the methods and approaches which seem to be commonly

agreed upon to produce accurate standardization and useful control. Secondly, .for those

areas where disagreement exists, we need to discuss the kind of experimental approaches

which might resolve these questions. Lastly, we need to look into our own crystal ball and

make some predictions on how standardization and control should be done in the future.

With these points in mind, I would like to express some of my own opinions on this

subject and also describe some of niy recent experiences with single phase control material.

Standardization should have optimally several characteristics. First, the material

used should be precisely definable. By that, I mean it should be describable in terms of

exact quantities of materials treated in a specific way to obtain the standard value.

Second, the material should be as close in characteristics to the sample as possible. When

dealing with whole blood as a sample, this is often difficult to achieve while maintaining

definability. Lastly, the standardizing measurement should be done in the same manner as a sample measurement. Often the particular instrument used requires different sample application procedures for sample and standard.

Standardization of pH measurements meets these criteria reasonably well. Gas standardization, however, is another question. By far, the most widely used technique for gas

standardization utilizes mixtures of pure gases flowing through a sample chamber. This

technique violates two of the three basic principles of standardization. First, it is a gas

while the sample is a liquid, and second, it is flowing, whereas the sample is static. The reason for its widespread utilization is that this material is easily handled and the

resulting gas tensions are quite reproducible. As long as the flow rate of the gas is

reasonably constant and humidif ication of the gas with water is accomplished at a constant

temperature, good day-to-day reproducibility is attainable. Biases associated with the

media and flow differences are accepted as the price of reproducibility. Thus, there is a gas-to-liquid correction factor associated with oxygen measurements which stems from the

choice of a gas as the primary calibration medium £1]^. Likewise, there is a difference in

response characteristics of the primary sensors, such that the response time in the flowing

gas is invariably shorter than the response time in a static liquid. Fortunately, both of

these effects tend to be constant and day-to-day reproducibility is generally attainable.

However, there is good evidence to suggest that during the operating life of these electrodes

and prior to obvious malfunction, changes occur in these characteristics which are only

evident by measuring liquids of known gas tensions.

Figures in brackets indicate the literature references at the end of this paper.

275
Methods for correcting for these effects have included calculation of error factors,

variation of sampling techniques [2], and equilibration of liquids with known gas compositions

[3]. This latter effort was equally directed at defining a single phase control material.

Sufficient progress has been made in this area to allow the recommendation of an all liquid

system for pH and blood gas calibration [4,5]. Commercial equipment is now available to perform the equilibration in an accurate and simple manner so as to warrant utilization of

these liquids as a standardization device.

The Radiometer Company has recently made a major advance in instrument design by

incorporation of liquid calibration into their latest blood gas system. The ABLl incorporates

into its design a calibration system which utilizes various buffers equilibrated with

mixtures of carbon dioxide and atmospheric air. It is important that we encourage Radiometer

and other manufacturers to continue in this sound technical direction for future instrumentation.

As modern equipment becomes more automatic, we are often required to use the calibration

scheme chosen by the manufacturer even though better calibration methods might be available.

Moving on to the subject of quality control, I feel methods of control should have two additional characteristics over and above those mentioned already. First, they must be

simple, because a tedious and complicated procedure just does not get utilized in a busy

laboratory. And secondly, they should respond predictably to a developing malfunction

before it becomes obvious without a control.

Many attempts have been made to prepare control systems with a whole blood matrix [6].

While these appear to have usefulness within a given laboratory, specimen variability and

handling difficulties prevent universal application of tonometered whole blood. On the

other hand, protein based controls for pH and PCO2 are commercially available [7]. There

has been a good deal of discussion as to the necessity of a protein based solution for

quality control, the feeling being that a protein solution most clearly acts as whole blood

would act. My attitude is that whatever advantages a protein solution might possess (and I am somewhat dubious that it has any), it is basically a nondefinable medium. I feel the

reproducibility and predictability inherent in a tonometered pure aqueous solution makes it

the media of choice for blood gas control.

Dr. Burnett and I reported on a particularly simple and also accurate device for the

preparation of liquid controls [5]. Prior to the availability of this system, it was very

difficult to get more than a subjective indication of the long term precision and accuracy

of blood gas results. We have had now nearly two years experience with this approach.

Table 1 shows cumulative daily statistics collected over a six month period at Hartford

Table 1. Blood gas quality control with tonometered bicarbonate (Hartford Hospital, 1974). \

1

pH • PCO2 PO2

I

1

(calc. val

.

= 7.35) (calc. val. = 86 mm Hg) (calc. val . =

i 149 mm Hg)

Mean S.D. Mean S.D. Mean S.D. !

Dec 7.34 .016 81 5 142 6 !

Jan 7.35 .012 79 4 145 3

1 Feb 7.34 .011 79 3 144 2

Mar 7.34 .022 84 7 146 3 !

April 7.36 .014 80 5 144 2 ?

May 7.35 .012 83 4 145

Port Huron 7.35 .013 80 3

276
Hospital and similar data for a one month period at Port Huron Hospital. These data show

the consistency of results obtainable over a long period of time with tonometered aqueous

solutions. It also shows that inter! aboratory comparison is also achievable with this

material. Use of material such as this will allow meaningful comparisons of interl aboratory

biases and errors. We have been in contact with several laboratories which have experienced

initial difficulties in matching our expected values. After some consultation, we were able

to identify se'rious unexpected system difficulties in some of the blood gas laboratories.

From the standpoint of commercially available materials, the General Diagnostics Company

has developed a series of aqueous, gas-equilibrated liquids packaged in glass ampoules with

assigned values of pH, PO2, and PCO2. This represents a particularly simple approach to avoid the complexity of tonometry. Table 2 shows some data which we obtained using this

material in our laboratory. Notice that the precision obtainable with this material for pH

and carbon dioxide is as good as that obtained with our own temperature-controlled tonometer.

The oxygen results are somewhat influenced by variations in room temperature. However, a more important factor in the case of oxygen was that we had difficulties establishing a consistent method for transferring our specimens to a syringe for insertion into our instrument.

It is expected that direct sampling will improve this precision. The real beauty of this

product is that it makes blood gas control a very simple experimental process. Undoubtedly,

this technological achievement indicates that we will soon see liquid standards available,

packaged in a similar manner, perhaps even certified by NBS.

Table 2. Blood gas quality control with liquid-filled ampoules

February 1975)

(Port Huron Hospital

pH PCO2 PO2

Mean S.D. Mean S.D. Mean S.D.

GD#1 7.06 .021 23 1.3 172 7.1

GD#2 7.39 .010 40 1.2 121 6.5

GD#3 7.64 .021 65 1.3 82 12.4

In closing my remarks here, I think we now have available the methods and materials to begin to resolve in an objective manner some of the nagging questions in blood gas technology.

In our own laboratory we are beginning a study comparing the methods of quality control over

a long period of time in regard to maintenance factors. I look forward to other activities

directed toward the resolution of the problems in blood gas standardization and control.

i References

[1] Adams, A. P. and Morgan-Hughes, J. 0., Determination of the blood-gas factor of the

oxygen electrode using a new tonometer, Brit. J. Anaesth. 39, 107 (1967).

[2] Instruction Manual for the Model 165 Blood Gas Analyzer (Corning Glass Works, Medfield,

MA, 02052).

'[3] Holmes, P. L. , Green, H. E. , and Majano-Lopez , V. L., Evaluation of methods for calibration

I of O2 and CO2 electrodes, Amer. J. Clin. Path. 54, 566 (1970).

[4] Veefkind, A. H., Van den Camp, A. M. , and Maas, A. H. J., Use of carbon dioxide and

oxygen-tonometered phosphate-bicarbonate-chloride-glycerol-water mixtures for calibration

and control of pH, PCO2 and PO2 electrode systems, Clin. Chem. 21_, 685 (1975).

[5] Noonan, D. C. and Burnett, R. W., Quality control system for blood pH and gas measurements,

with use of a tonometered bicarbonate-chloride solution and duplicate samples of whole

blood, Clin. Chem. 20, 660 (1974).

277
[6] Kenny. M. A., Delaney, C. J., and Raisys, V. A., A P'-oficiency testing and quality

control program for blood gas analysis, Clin. Chem. }3_, b4y 6).

[7] versatol Acid Base (General Diagnostics, Division of Warner Lambert Company, Morris

Plains, NJ, 07950).

278
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DEVELOPMENT OF REFERENCE METHODS: BLOOD GAS ANALYSIS

Arthur L. Malenfant and Kevin D. Fallan

Instrumentation Laboratory Inc. 113 Hartwell Avenue

Lexington, Massachusetts 02173, USA

1. Introduction

The nature of the parameters measured in blood gas analysis leads to unique problems

in attempting to establish reference methods which will serve as an accepted norm against

which manufacturers of instrumentation and practitioners of the method may compare the

results which they obtain. The species are extremely labile both because of changes induced

within the sample and also because of the ease of contamination of the sample by atmosphere. Metabolism of the sample results in lowering of Pq and elevation of ^qq^, while

mixing with air will reverse this situation.

Since the analysis of blood for pH, P^q and Pq must be performed in an environment

which may easily lead to errors resulting from contamination, reference methods must pay

strict attention to both operator dependent factors and instrument dependent factors.

The primary purpose for establishing reference methods is to establish the precision

and accuracy of the procedure used to obtain blood gas information in the clinical laboratory. As in any analysis, accuracy is defined as the expected- or absolute-value. The

nature of the measuring system requires that the reference methods be established to define

the accuracy and precision at the level where the Pq and/or P^q values are being measured.

It has been prevously documented that calibration for P^ should vary with the level

U2

measured [1,2]^. Instruments calibrated to achieve Pn measurements which are within 2 U2

percent of the expected value at 50 mm Hg must be calibrated differently to achieve the same performance at Pq 's of 300 mm Hg. The measurement of Pqq in the clinical laboratory

is over a much narrower range, however, and this requirement is not evident for that

measurement.

While accuracy of results is to be sought as a worthwhile goal, the importance of

precision ought not to be underestimated. Monitoring of the clinical state of a patient

while following the results of sequential measurements over several hours or days provides

a clinically useful tool. Therefore, while results might even be in error in absolute

terms, good precision of results will permit successful care of the patient.

^Figures in brackets indicate literature references at the end of this paper.

279
2. Operational Requirements

The number of factors which affect the accuracy of blood gas is significant. In order

to organize them, they will be described as operator dependent factors and instrument

dependent factors.

Operator dependent factors which affect results begin with the drawing of the blood

sample for the analysis. There will be no attempt to deal with the actual technique used

by the clinician in drawing either arterial or venous samples. The attempt will be made,

however, to delineate aspects which can affect blood gas results by actions taken both

before and after the sample is drawn.

The syringe used to draw the blood will have an effect on the results achieved. There

is literature which indicates that only glass syringes permit accurate results [3] while

other literature indicates that disposable plastic syringes are suitable [4]. Our ex- perience has been that as long as the blood is not stored for extended periods between

drawing and analysis, no significant bias is introduced. If samples are stored under iced

conditions for several hours, changes consistent with increases in Pq and decreases in

Pj,Q caused by diffusion through the walls of the plastic syringe are observed when samples

stored in plastic are compared with samples stored in glass.

The amount and nature of the anticoagulant used also affects the blood gas results

achieved for pH and P^q . The anticoagulant of choice for blood gas analysis is sodium

heparin [5]. At the present time, however, standardization of the number of units of

heparin per ml of sample is not adequately defined. As a result, variation in the volume

of anticoagulant used affects results through dilution of the sample, as well as through pH

changes induced by using other anticoagulants such as ammonium heparin [6].

For the measurement of P^ and for the direct photometric measurement of oxygen

saturation, the choice of anticoagulant is not so critical since the only effect on these

measurements would be through dilution of the sample by the oxygen contained in the anticoagulant.

Once the operator has a properly prepared sample, the next factor under the control of

the operator is the proper calibration of the blood gas analyzer. Accuracy of calibration

of Pq and P^q is dependent on the accuracy of the calibrating media, whether liquid or

gas, and on correction for barometric pressure changes. Manufacturers of calibrating gases

standardize their values through either the Scholander technique or gas chromatography.

Within a given geographic region, day-to-day fluctuations in barometric pressure will be

small, but many, if not most locations will have a standard barometric pressure lower than

the 760 mm Hg standard at sea level. At 3300 feet elevation the barometric pressure would

average about 674 mm Hg.

280
The accuracy of blood gas results, even when each of the necessary calibration steps

is carried out properly, is also dependent on the operator. The instrument must be given

proper maintenance. At present, totally blood compatible materials are not available for

construction of blood gas analyzers. Therefore, regular routine cleaning of the system is necessary to prevent buildup of protein residues from previous analyses. The advent of

automation in blood gas analysis insures that samples will each be handled similarly once

they are inducted into the instrument, but without proper maintenance, even these systems

will not yield results which are accurate in the long term.

Instrument dependent factors present in achieving accurate, precise results in blood

gas analysis deal primarily with the proper calibration of the sensors and introduction of

the blood sample into the measuring chamber of the system.

In the near term, establishment of standards for blood gas calibrators is required.

At the present time, the use of phosphate buffers serves to establish the calibration of

the glass electrode for monitoring pH of blood [7]. Calibration of the gas sensors,

however, continues to be a source of some controversy. The use of standardized, humidified

calibrating gases flowing past the tip of the sensors has been the calibration method of

choice for many systems. However, the question of the gas/liquid membrane correction

factor [8] for oxygen measurements has yet to be adequately resolved. Proposals have been

made for alternatives such as tonometered glycerol -water solutions containing phosphate,

bicarbonate and chloride [9] to serve as the calibration medium. Proposals for other

fluids such as tonometered blood [10] and sodium bicarbonate--sodium chloride solutions

equilibrated with known gas mixtures [11] have also been made, the latter being proposed as

a quality control procedure.

3. Methodology Currently Employed

Currently available control materials for characterizing the reliability of blood gas

results fall into two categories. The first category includes serum-based controls that

are reconstituted by addition of an appropriate solvent. The resulting solution, once reconstituted, is then inducted into a calibrated blood gas analyzer and the results are compared with the nominal values to be expected. These require little time to reconstitute

and are relatively simple to use. Disadvantages in attempting to follow this approach,

however, to develop a standard reference material include the requirement for a high degree

of care in mixing with the solvent to prevent undue contamination with air before sampling,

the fact that the values are established by the same technique as the method being tested,

the temperature coefficient is different from that for blood, and the material must be

handled differently than a blood sample would be handled. This last would lead to incorrect assumption either of problems or lack of them.

Another control material which might be considered is the aqueous based control material

which is supplied in liquid form with known values of pH, P^q and Pq . In this case the

mixing step is completely eliminated. The solution need only be inducted into the blood gas

analyzer and the results recorded. The disadvantages here, however, are similar to the

serum-based material above with respect to a standard reference material. First of all, it

281
is simply an aqueous solution, not blood. Also, the range of values available is limited of

necessity and may not include some of the more extreme ranges where blood gas analysis may

be performed.

In addition, neither of the above approaches tests the operator's ability to properly

handle blood samples under controlled conditions. As noted previously, this is an important

factor in achieving accurate blood gas results.

4. Tonometry of Whole Blood

It is proposed that the method of choice for developing a reference method for blood

gas analysis be the tonometry of whole blood. This approach has a number of advantages.

First, the method being used is the primary method. Second, all controlled parameters are independent of the measuring system, i.e. established without reference to a blood gas analyzer. Equilibration of the whole blood can be made with gases having accurately known

Pq 's and P^q 's. Instrument response will be identical to that when clinical samples are being measured. The operator is handling whole blood in the same way as a clinical sample

once the blood is withdrawn from the tonometer. Any deviations observed will be identical

to those to be experienced with clinical samples. Temperature coefficients are identical,

for example. The pH of whole blood may be adjusted by addition of bicarbonate so that it is both known and predictable under specific tonometry conditions to be used to establish the

Pq and P^q values desired.

Below is shown the results' achieved using the approach described. A volume of whole

blood was standardized to a known pH by addition of bicarbonate. This was then sealed into

containers and kept refrigerated. Over a thirty-day period these blood samples were to- nometered with gases of known composition, and the pH, P^q and Pq determined after to- nometry. The results are tabulated as shown. The analyses were performed both at Instru- mentation Laboratory, as well as at another institution. The gases used for tonometry were: Location No. 1 Location No. 2 Gas No. SD SD

1^ pH: 7 370 0.006 pH: 1 360 0.007

P 49 5 mm Hg 0.6 ^^002 46 5 mm Hg 2.0

v

= 51

.

7 mm Hg 1.0 51. 4 mm Hg 1.95

2^ pH: 7 429 0.007 pH: 7 420 0.010

^C02'

35 8 mm Hg 0.44

''C02" 33 9 nni Hg 0.77

139 4 mm Hg 1.4 137 5 mm Hg 2.1

3^ pH: 7 511 0.006 pH: 7 500 0.013

^C02"

21 7 mm Hg 0.4 P 20 7 mm Hg 0.58

75 1 mm Hg 1.3 75 00 mm Hg 2.8

^Gas No. 1: 7.0 percent O2, 7.0 percent CO2, balance N2

"^Gas No. 2: 20.0 percent O2, 5.0 percent CO2, balance N2 ^Gas No. 3: 10.0 percent O2, 3.0 percent CO2, balance N2

282
Thus, it can be seen that consistent results may be achieved over a period of time

using tonometry with gases of known composition to bring whole blood samples into gas

equilibrium. Although the range shown here is relatively limited, it could easily be

extended to, for instance, high Pp. 's by use of a gas having a P^ of 300 or 400 mm. U2 U2

References

[I] Key, A., Biom. Eng. 9, 1954 (1974).

[2] Reppeto, N. P. and Moser, K. M. , NSCPT Analyzer, 4 (2), 13 (1974).

[3] Sackner, M. A., Med. Times, 95, 79 (1967).

[4] Winkler, J. B., et al.. Chest, 66, 518 (1974).

[5] Gambino, S. R. , et al.. Annals of N.Y. Acad. Science, 133, 259 (1966).

[6] Shapiro, B. A., Chicago Year Book, Med. Pub., 167 (1973).

[7] Bates, R. G., J. Res. Nat. Bur. Stand. 66A, 179 (1962).

[8] Reppetto, N. P. , ihid.

[9] Veerkind, A. H,, et al., Clin. Chem. 21_ (6), 685 (1975).

[10] Adams, A. P. and Morgan Hughes, J. 0., Br. J. Anesth. 39, 107 (1967).

[II] Noonan, D. C. and Burnett, R. W. , Clin. Chem. 20_, 660 (1974).

283
I

I

I

I

I

!

I

1

1

!

i

i

I

i

I

I

i

I

I

1

]

1

i

i

i

I
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975, Issued June 1977.

QUALITY CONTROL AND STANDARDS

S0ren K. S0rensen

Radiometer A/S

Copenhagen, Denmark

The term "quality control" can be interpreted in various ways. In a narrower sense, it can thus be taken to mean no more than a quality check-up of a measuring system. In a wider sense however, it aims at establishing whether the results of tests performed on a sample actually do represent the bodily condition of a patient. It is, of course, this

latter objective of quality control that is the ideal one, but unfortunately it is not

always attainable.

The flow diagram for a sample is seen in figure 1. A proper quality control is performed by letting a known sample run through the whole system— in other words, it is the

reference person we are looking for. It has been attempted [1]^ as a tool for controlling

the "daily mean" to take blood samples from the staff everyday and then use the results as control material. It is quite natural to use blood in quality control as it is the sample

material, but blood is a very special liquid and it is difficult to imitate the physical

and chemical behavior of blood in other media. When trying to simulate blood, it is necessary to get the correct buffer capacity for pH, PCO2, and P02, where the oxygen

buffer is the most difficult to establish. It is necessary to simulate what is called the

suspension effect from the red blood cells when measuring pH. It is necessary to simulate

dissolved protein and the viscosity of the blood. A list of specifications can thus be

made, from which a lot of practical "artificial blood references" can be produced.

the

blood

sampling

the

transport

the measure- ment

the

result

assessment

of

the result

Figure 1. Flow diagram for running a sample.

In some types of analysis, it is possible to get very close to the ideal condition,

e.g., serum reference materials used in quality control of the inorganic salts are excellent

synthetic substitutes for the genuine sample. But as far as pH and blood gases are concerned,

it is not possible to produce the stable reference material that is a necessary condition

for carrying out quality control [2].

Quality control is performed as spot tests, but no test can improve a weak system.

The spot tests are therefore unable to solve quality problems; they can only indicate that

something is going wrong. This means that one cannot merely concentrate on finding the

best of all reference materials. It is equally important to ensure that inbuilt security

in the system is provided. Internationally accepted rules for the application of reference

methods, units, formulas for derived parameters, choice of derived parameters and sample

handling are some of the indispensable conditions on which the quality control must be

based. Internal quality control is much to be the preferred, but technical limitations,

e.g.. Stability problems, and legislation requirements are some of the reasons why external

quality control is utilized. Western European countries, and especially the Medical Society

Figures in brackets indicate the literature references at the end of this paper.

285
of West Germany have shown a great interest in external quality control. Their guidelines

do not deal with blood gases, however, as they have based their quality control on test

sera. External control may be a useful addition to laboratory control, but the time delay

involved is reason enough for rendering it unsuitable as a primary control basis in blood

gas analyses. As a first step to quality control of blood gases, collaborative control

within small geographical regions is a possibility. This could be a valuable, although

expensive, way to build up a control system.

This author feels that it is important to establish that both choice of derived parameters and units are part of a quality control program. One of the most important things in

quality control is the log book. Without recording the trends and the change in scatter of

the results, it is impossible to foresee the coming break-down of the system, which always occurs at the most inconvenient time in the day. To reduce the shut-down period, the

quality control log book is thus an indispensable tool.

It is very difficult to solve the entire quality control problem for pH and blood

gases, but even a partial solution will be a great improvement over the present situation.

Before starting a control program, it is natural to ask what limits of accuracy and precision

are necessary to give the control a meaning. Manufacturers usually refer to the instrumentation specifications. However, these specifications are a result of technical limitations and, therefore, are not necessarily in agreement with the clinical requirements.

The best plan is to let clinical needs, nbt the technical ability, determine the limits of

accuracy and precision. This point is very important.

Some authors [3] are of the opinion that the methodological inaccuracy must not be

greater than one SD of the normal range. The normal range is then 4 SD, which incorporates

both the biological and the methodological variation. For the blood pH, this means an accuracy around 15 mpH, which should be obtainable. In Pco2 measurements, the same range

is 3.5 mm Hg, which can be easily reached.

As regards the separate parameters, pH is the best established. There is a definition

of pH--whether good or bad--and there are buffers which can be used to calibrate the electrode.

Pco2 can be determined directly with a gas-sensing electrode, or it can be determined

indirectly by means of the equilibration function of the sample. In either case, an analyzed

gas mixture is the basis for evaluating the quality of the measurement. P02 is determined

directly with a P02 electrode, and here again gas mixtures form the basis of calibrations.

The lack of a reference material is not the only hindrance to careful control. Bloodgas measurements are typical stat measurements of one to two samples at a time, which means

a Q.C. every second sample is desirable, and that is, in practice, impossible. A complete

calibration of the system is therefore preferable just before the measurement is performed,

but in practice even this is normally not possible. A compromise involving frequent

calibrations during the day can, however, improve the quality of the measurement.

For calibrating purposes, gases are satisfactory, but as a medium in quality control,

a liquid is much to be preferred. A minimum requirement is that quality control shall be

performed on a liquid with known partial pressures of carbon dioxide and oxygen— in other

words--a tonometered test solution. Blood is, of course, the preferred liquid for that

purpose. The preferable control procedures are listed in figure 2.

pH: Buffer, Ton. bicarbonate solution

PC02' Blood, Ton. bicarbonate solution

P02= Blood, (Ton. aqueous solution)

Figure 2. Liquids for quality control.

It is not a sophisticated way of performing quality control, but, on the other hand,

many laboratories do not even reach this level, and those that do are satisfied. That leads

to another problem, namely the gap in the technical standard between the leading laboratories

286
and the ordinary ones, even in relatively homogeneous countries. It is very important that

this gap in technical standard does not grow even wider, thereby creating a situation where

quality control procedures become unattainable for lesser laboratories so that there is a risk that they will refrain from using them. Quite likely, such laboratories would be the ones that stand in greatest need of quality control.

The development problem has other aspects. If the laboratories are unable to prove

their willingness to document the quality of an analysis, some other authority will take

the initiative and create rules which, in the worst case, are impracticable, and in the best

case, impractical. The same demands must be made with respect to the industry that produces

equipment and diagnostic products for the laboratory, as a quality control program will

only be a success if complete confidence exists between authorities, laboratories, and the

industry involved.

The tonometer technique will no doubt give a lot of problems owing to inadequate

training and experience. The only possibility is to promote the tonometer technique and

then to learn from experience whether it is usable for control purposes. The manufacturers

of blood-gas equipment must therefore be prepared to take the great responsibility for

producing tonometers of a reliable and simple technical construction. Today, only a few

companies produce a commercial tonometer. A tonometer, however, is only the tool used to produce gas equilibrium. The standard is the gas mixture which is used to equilibrate the

blood.

Gas mixtures can be obtained everywhere but they have to be analyzed before they are usable as standards in tonometry. That is a point where cooperation on an international

plane could produce fast results. Determination of the compositions of gases is purely a measuring problem, but internationally certified gas mixtures could remove one uncertainty

from tonometry and also from the calibration procedure.

From a technical point of view, tonometered buffers are easier to handle than blood,

and it might be worthwhile to try to evaluate in practice all the differences that can occur between blood and aqueous solutions. This compromise using buffers instead of blood,

can give the unskilled laboratory a chance to start a control program without being completely

lost from the beginning. Because blood-gas analyses are so complicated, the need for good

quality control is immense, but I believe that progress here is more a matter of perspiration

than of inspiration.

References

[1] Whitehead, T. P., Methods of quality control, Ann. Clin. Bioohem. 6_, 94 (1969).

Broughton, P. M. G., The future of quality control, Ann. Clin. Bioohem. 6^, 147

(1969).

[2]

[3] Gowenlock, A. H. , The influence of accuracy and precision on the normal range,

Ann. Clin. Bioohem. 6_, 3 (1969).

287

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

DEVELOPMENT OF REFERENCE METHODS

S0ren K. S0rensen

Radiometer A/S

Copenhagen, Denmark

The precision of a measurement can be determined without having a reference method,

whereas the accuracy may elude the determination without an established and well -documented

reference method. The reference method is the most reliable method irrespective of analysis

speed and cost, which implies that the reference method is not necessarily identical with

the routine method.

A routine analysis is the practical procedure adopted by a particular laboratory. A routine analysis is, of course, tested against the reference method to expose the limitations

of the former. The routine method, which may not be the reference method, is classified in

the group of methods containing bias--known or unknown. It allows one to select just that

method of analysis that meets the prevailing conditions {viz.,, local legislation, technical

training, instrumentation, and future development).

A proposal for categorizing analytical principles has been made by IFCC. According to this definition, a reference method is a method with an accuracy = 0 ± 6, where 6 is

negligible in comparison with interlaboratory imprecision. The ultimate method, called

"definitive method", has no known source of inaccuracy.

In order to obtain the total concept, a definitive method must be available. For pH it

is not possible to set up a definitive method, as I will explain later, but for both Pcoa

and P02 it is--at least in theory--possible to measure the partial pressure in a system by

known, established techniques. What I have in mind is to try to investigate a method for

measuring the total amount of carbon dioxide or oxygen in an infinitely small quantity of

gas. The problem with blood is that the total content of carbon dioxide or oxygen does not

give any information on the partial pressure of the gases in the blood. In other media it

is, however, possible to get such information due to simple dissolving processes in these

media. By equilibrating, e.g.^, a small piece of plastic in the blood and measuring the

total gas content, the partial pressure can be calculated. The only limitation to such a method is the technical problems. With the advanced technical skill of today, it should,

however, be possible to develop a "definitive method" based on this principle, but as far as

I know this has not yet been attempted in practice.

The very restrictive definition of reference methods is difficult to meet especially

for such complex systems as pH and blood gases. The day when such reference methods are available lies somewhere out in the future, and in the light of this definition even pH is a doubtful parameter. pH determination on blood is based on an operational definition, using

a hydrogen electrode and a saturated KCl liquid- junction. Especially the residual liquidjunction potential gives a more or less unknown bias on the measurement. Nevertheless, no one will deny that pH in blood is a valuable and important parameter, and there are, in fact,

no means of eliminating the liquid-junction problem as long as electrochemical cells are used to measure the hydrogen ion activity. The normal way of measuring pH is, of course, to use a glass electrode instead of a hydrogen electrode. The junction potential can be changed

by using other types of reference electrode systems, but it cannot be eliminated and it is

unlikely that anyone has tried to measure blood pH with a hydrogen electrode! What can be

done is to accept the pH definition as it is or, better still, to improve it with a more

precise description of the liquid-junction type of salt bridge and the geometry of the

junction.

289
pH is a good example of a practically defined parameter which has shown its worth, even though several standardizing committees have tried to change the units for pH. It is my

opinion that we still have the best of all proposed solutions to pH in the operational pH

definition, which requires only small changes to lead to what we may call an "assigned

method."

Turning to the PCO2, matters are far more complicated. It may be possible, as already

mentioned, to think out a definitive method which, no doubt, would be useless in practice.

There are two established methods for determining PCO2, one measures the CO2 partial pres- sure with the Severinghaus-type carbon dioxide electrode. The other is the equilibration

method, known as the Astrup technique, developed by Astrup and Siggaard-Andersen. Neither of

these techniques is a direct measurement of CO2 pressure or carbon dioxide activity. It may

seem obvious to favor the electrode technique, as one can argue that an electrode gives a more direct measurement than the equilibration technique.

As all other reference methods, the CO2 reference methods require that certain conditions are defined, for example, fixed and known temperature, oxygen saturation of the sample

and hemoglobin concentration. However, a lot of problems connected with the electrode are known, and can be classified as interferences. Leakage of hydroxide ions from the pH glass

surface, different osmotic pressure in the sample and the electrode electrolyte, and loss of

gas into the bulk electrolyte, are some of the problems which are not easily overcome.

Using the Astrup technique as reference method is not the ideal solution, but for

several reasons I feel that this method offers certain advantages. By tonometer ing the same sample as used for measurement, the equilibration curve is established directly from the

sample. The tonometry process will offer a fixed saturation. When using an electrode,

PCO2 is converted to pH in an electrolyte layer after diffusion across a thin plastic membrane, and a calculation is necessary. In the equilibration technique, on the other hand,

the electrode--indeed, one of the most reliable ones--is used for both calibration and measurement. The Astrup method thus eliminates the Henderson-Hasselbalch equation. One disadvantage of the equilibration is the need for defined or known oxygen saturation of the blood

sample. The Astrup technique will, of course, not oust the gas electrode technique, but it

is my opinion that we have a better theoretical background when using an equilibration technique as reference method.

The P02 analysis is, no doubt, the most complicated one for which to propose a reference

method. Compared with PCO2, for which two well-established methods are available, P02

offers only one method, namely the P02 electrode itself. Admittedly, an electrode which in practice has shown its validity, but whose theoretical basis is very weak, has a lot of

known interferences and side-reactions. To found a reference method on the P02 electrode

itself will hardly lead to success.

Here, too, there is the possibility of using an equilibration technique. If one has a precise knowledge of the oxygen dissociation curve, then the electrode technique can be

converted into an optical determination of the oxygen saturation. Experience with the

optical saturation measurements are fairly good, but the other side of the problem--prior

knowledge of the oxygen dissociation curve--may kill the idea. Several parameters change the

dissociation curve. Examples are pH, PCO2, and enzymes (2,3-DP&-). I feel that optical

analysis based on a carefully selected set of definitions in order to lock the dissociation

curve in position would provide a better reference method than any system founded on the

polarographic oxygen electrode. If the outlined methods are discovered to be useless, only

the definitive methods are left. As a matter of form, it may be recalled that potentiometrically working oxygen electrodes based on redox-potential measurements do exist, but so far,

problems connected with the irreversibility of the material have limited its use. Someday,

a good, reversible, oxygen-binding substance may lead to workable electrodes.

pH none

'C02 determination

in "microbubble"

determination

in "microbubble"

Figure 1. Definitive methods.

pH glass electrode

CO2

'O2

Astrup method

saturation

equilibration

technique

Figure 2. Reference methods.

290
The reason for choosing a method based on equilibration both for carbon dioxide and for

oxygen is that, from a scientific point of view, the processes are better known. From a practical point of view, the suggested methods are more debatable, and a simpler way to solve the reference method problem is not to base the measurements on the gas sensing system,

but to use the tonometer as the reference method. That means to develop a method which is

able to produce a reference material --the tonometered blood sample.

I have now tried to establish a system of methods which could serve as an inspiration

in our endeavors to develop a suitable reference method. No doubt, it will be clear to everybody that, so far, these methods are carried out only as suppositions, but they do at least expose the fields which in my opinion ought to be studied.

291

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

STANDARDIZATION OF ION-SELECTIVE ELECTRODES FOR SERUM ANALYSIS

M. S. Mohan and Roger G. Bates

Department of Chemistry

University of Florida

Gainesville, Florida 32611, USA

Electromotive- force measurements at 37 °C in synthetic electrolyte mixtures simulating

serum provide evidence that ion-selective electrodes respond in a near-Nerstian manner to changes in concentrations of the species to which they are selective. The relatively

minor effects of changes in ionic strength on the electrode potentials are almost entirely

accounted for by changes in the activity coefficients. It is suggested that similar

electrolyte mixtures may serve as calibrating standards in clinical monitoring with ionselective electrodes. The lanthanum fluoride electrode shows promise as a useful reference

electrode in media of this sort.

1. Introduction

In the decade which has passed since their introduction, ion-selective electrodes have

found increasing use in a variety of biological applications [1,2].^ In addition to electrodes

designed to measure electrolytes in body fluids, novel devices based on the "building block"

principle [3] have been constructed to determine other clinically important entities such as urea [4] and glucose [5]. Changes in the configuration of existing electrodes, such as immobilizing the exchanger in a polymer matrix, have resulted in highly reliable and rugged

devices [6].

In the analysis of electrolytes, an important advantage of ion-selective electrodes

over existing clinical methods such as photometry is that they respond to fvee ionic activities or concentrations rather than to total concentrations. If one could obtain free ionic

concentrations (which are the medically significant quantities) from measured ion-selective

electrode potentials, the utility of ion-selective electrodes would be very significantly

enhanced. Such a procedure is dependent upon the development of reliable calibration procedures, particularly in the complex mixtures of electrolytes found in biological fluids. Two

major problems must be faced in establishing suitable standards. First is the experimental

limitation imposed by the imperfect selectivity of the electrode; second is the theoretical

difficulty attending a conversion of activity to concentration.

2. Selectivity

The empirical Nicolsky-Eisenman equation [7] is useful in describing the response of an electrode selective for an ion in the presence of an interferent j:

Eg^ = Constant ± log |^a. + Zk-^.a

.'''^'^j

(1)

where Eg] is the electrode potential, a^ and aj are the activities of the primary ion and

the interferent, respectively, z-j and zj are the corresponding electrical charges, and k^j

1 Figures in brackets indicate the literature references at the end of this paper.

293
is the selectivity ratio. This equation holds fairly well at low interference levels but,

generally speaking, the selectivity ratio is not a constant and depends on the activities of

i and j. Since the relative concentrations of ionic constituents in body fluids are likely

to vary under disease conditions, it becomes necessary to examine the extent to which measured

potential differences are affected by selectivity parameters.

3. Activities vs Concentrations

The development of an operational pH scale was facilitated by the availability of

buffer substances spanning a wide range of pH values at comparatively low ionic strengths

between 0 and 0.1. This not only enables one to separate conventionally the thermodynamically

measurable mean ionic activity coefficients into single ion activity coefficients but assures

minimal residual liquid-junction effects in practical measurements. In the calibration of

ion-selective electrodes, used to determine the clinically important ions, the ionic strengths

of the solutions may exceed 0.1 and may vary rather widely, spanning a wide range of activities

[8]. This would mean not only that the simplest formulas for the conventional separation of mean ionic coefficients into individual ionic contributions would no longer be valid but

that the residual liquid-junction potentials might be large enough to cause appreciable

errors in the measurements.

Metal-ion buffers offer a possible solution. Although buffers for calcium ion, based

on the EDTA and NTA complexes, have been proposed [9], the measured pCa values depend heavily

on the accuracy of the conditional formation constants. Furthermore, ligands complexing

with Na"*", K"*", and other alkali metal ions of sufficient stability in aqueous media are not available.

Despite the success of the convention based on hydration numbers [10] in certain special cases, the evaluation of single ion activity coefficients in multi component electrolyte

mixtures continues to pose considerable difficulties. On the other hand, the ionic strengths

of body fluids are sufficiently low (usually less than 0.17 mol 1"^) that the actual single

ionic activity coefficients are likely to be reasonably close to the values derived from the

conventions that have proved useful at ionic strengths less than 0.1. Taking these factors

into consideration, suitable standards for calibrating ion-selective electrodes for use in

biological fluids are provided by a set of synthetic electrolyte mixtures whose ionic

concentrations span the range normally encountered in clinical analysis. It has been shown

[11] that the electrode potential responds in a near-Nernstian manner to changes in the

aoncentvation Of the ions of interest at the normal ionic strength of serum, although the

slopes S = Eel/log C may depart slightly from the values given by theory. In addition, it

appears possible to apply corrections for changes in ionic strength if these occur due to disease conditions.

4. Emf Measurements in Standards of Constant Ionic Strength

Measurement of electromotive force at 37 °C in cells with and without liquid junction were made with a series of calibration standards in which the concentrations of Na^, K"*", Ca^"*", H"^, and CI" were varied simultaneously between the lower limits encountered under

clinical conditions and their normal values. The indicator electrodes used in the study

were: sodium glass electrode (Corning), potassium electrode (valinomycin exchanger in

diphenyl ether), calcium electrode (Simon lipophilic Ca^"*" exchanger [12] immobilized in

polyvinyl chloride), and pH glass electrode (Corning). Tris(hydroxymethy)aminomethane

("Tris") was used as the buffer substance. The buffer ratios in the reference solutions were varied and pH derived from pK'bh+ = 7.907 in isotonic saline at 37 °C; this apparent

dissociation constant for protonated Tris was calculated from data given by Durst and

Staples [13]. Nitric acid and sodium nitrate were used to vary the buffer ratio and to maintain the ionic strength constant at 0.16 mol 1"^. Details of the preparation of the

standard solutions and other experimental procedures have been given elsewhere [11].

Compositions of the standards of constant ionic strength are given in table 1. As already mentioned, it is important to eliminate or minimize liquid-junction potentials in

order that the measured potentials truly reflect changes in ionic activities or concentrations.

For this reason, particular attention was given to the use of cells without liquid junction,

with reference electrodes immersed directly in the standard solutions. The silver-silver

294
Table 1. Compositions of calibration standards.^ (All solutions have

an ionic strength of 0.16 and contain 1 mmol NaF per liter.)

Standard

Number Sa ^Ca ^Cl PH

1 100 1 .0 0.30 66.6 7.90

2 105 1 .4 0.38 70.0 7.80

3 110 1 .8 0.46 75.0 7.70

4 115 2.2 0.54 80.0 7.60

5 120 2.6 0.62 85.0 7.55

6 125 3.0 0.70 90.0 7.50

7 130 3.4 0.78 95.0 7.45

8 135 3.8 0.90 100.0 7.40

9 140 4.0 1.00 106.0 7.36

^C^ in mmol 1

chloride electrode is of limited usefulness here, since the chloride ion is one of the

constituents being monitored. The reliability of the lanthanum fluoride electrode as a reference was investigated. Accordingly, small concentrations of fluoride ion were added to each of the standard solutions. Possible side reactions such as precipitation of calicum

fluoride or adsorption of the fluoride ion on the walls of the containing vessel were shown

to be of little concern [11]. The emf of three types of cell, namely

SCE II Bridge Solution || Standard ] ISE A

Ag;AgCl | Standard | ISE B and

LaFg I Standard ] ISE C

was measured. Steady potentials (to ±0.1 mV) were usually obtained after about 45 minutes.

In order to obtain the indicator electrode slopes from data for cells of type B, the

contribution to the total cell emf made by the changing chloride ion concentration was allowed for by assuming Nernstian response of the Ag;AgCl electrode. It is evident from the

least squares regression parameters for different electrode pairs given in table 2 that the

electrodes responded in a near-Nernstian manner to changes in the concentrations of the ions

being sensed when the ionic strength was constant.

This conclusion can be confirmed by examining the variation of the activity coefficients

in individual cases. For example, the emf of cells of type B can be written as

E = E°' + S log C^^ + 61.54 log C^^y_ (2)

where Cm and Cq] are the concentrations of the cation being sensed and the chloride ion,

respectively, and y+ and y_ are the corresponding activity coefficients on the molar scale;

S and 61.54 are the Nernst slopes (in mV) at 37 °C for the indicator electrode and the

Ag;AgCl electrode, respectively. On rearrangement, we obtain

E°" E E°' + S log y^ + 61.54 log y_ (3)

= E - S log C„ - 61 .54 log Cp, .

295
Table 2. Parameters for the calibration of ion-selective

electrodes in standards of ionic strength = 0.16.

Indicated ion

(i)

Reference

electrode

E_,,(mV at 37 °C) ce 1

1

S D.

SCE 435.1 - 60.9 pH 0 15

LaFg 437.0 - 61.15 pH 0 32

AgjAgCl 0.68 - 62.20 pH 0 11

SCE 63.3 + 57.3 loq C„ 0 16

LaFg 58.14 + 55.4 log C^ 0 40

Ag;AgCl 0.38 + 58.9 log C^ 0 16

SCE 36.58 + 54.04 loq luy C,y.^^ 0 06 LaFa 37.06 + 54.9 log C..Na

0 27

Ag;AgCl 0.35 + 56.4 log C..

IN a 0 21

Ca^-' SCE 124.9 + 27.8 loq 0 10 LaFa 122.9 + 27.3 log C^^ 0 29

Ag;AgCl 0.25 + 29.8 log C^^ 0 14

^Standard deviation for regression, in mV.

Constancy of the right-hand side of eq. (3) will confirm the invariance of the activity

coefficients. From the low standard deviations for the values of E°" (table 3), it is evident that the activity coefficients remain substantially constant throughout the range of

calibration.

Table 3. Constancy of activity coefficients in calibration

standards of ionic strength = 0.16.

Indicated ion Range of C. E°" S.D.

(i) (mmol r^y (mV) (mV)

Na"^ 100 to 140 49.00 0 18

1.8 to 4.0 77.80 0 20

Ca^-^ 0.3 to 1 .0 140.00 0 14

In measurements of cell A with the SCE as reference, the possible variation of the

junction potentials with the composition of the standards must be considered. Taking as example the cell of type A with a sodium indicator electrode and SCE reference, one can write

E°' + E. = E - S log C^g (4)

where Ej is the algebraic sum of all the junction potentials. The right-hand side of eq.

(4) was found to remain constant throughout the calibration range with a standard deviation

of only 0.2 mV. This result is important in that it attests to the effectiveness of the

constant ionic strength in stabilizing the potentials at the liquid junctions.

296
5. Emf Measurements at Varying Ionic Strength

Under disease conditions, the ionic strength of serum may vary markedly from its normal

value near 0.16 mol 1"^. It thus becomes necessary to examine the effect of this variation

on the measured electrode potentials. It is then possible to identify and evaluate the errors involved when the electrodes calibrated at one ionic strength are used in the widest

range of ionic strengths encountered in the analysis of serum.

The effect of changing ionic strength was studied by preparing three additional sets of

solutions. In all_sets, the concentrations of K+ and Ca^"*" were maintained at their normal

levels of 4.0 x 10"^ and 1.0 x 10"^ mol 1"^, respectively. The compositions are summarized

in table 4. In Set 1, the ionic strength was varied between 0.12 and 0.2 by changing the

buffer concentration while the Na"^ and CI" concentrations were kept constant at 0.1 mol 1"^.

In the solutions of Set 2, the sodium ion concentration was constant at its normal value of

0.14 mol 1"^ and the ionic strength was varied from 0.16 to 0.2 by altering the buffer

concentration. In the third set, the buffer concentration was maintained constant and the

ionic strength was varied from 0.12 to 0.2 by changing the concentration of sodium chloride.

Table 4. Compositions of solutions of ionic strengths

from 0.12 to 0.2a.

Set No. ^Na '^Cl 4 ^Ca

I

1 100 100 4.0 1.0 0 12 - 0 2

2 140 100 4.0 1 .0 0 16 - 0 2

3 100-180 106-186 4.0 1 .0 0 12 - 0 2

C-j in mmol 1'^

; pH = 7.4 at I = 0. 16. All solutions contained NaF

at a concentration of 0.1 mmol 1"-^.

At these ionic strengths, single ion activity coefficients could be calculated with

sufficient accuracy for the present purpose by the Debye-Huckel equation in the form

-log y,

- 'IL^ (5)

l+Ba/r

o with the following values applicable at 37 °C: A = 0.5232, B = 0.3316, and a (the ion-size

parameter) = 4.6A.

The emf of cells of the types

and

Ag;AgCl ] Solution | (ISE) D

Ag;AgCl | Solution | Ca^"^ (ISE) E

LaFg 1 Solution | K"^ (ISE) F

LaFg 1 Solution | Ca^"^ (ISE) G

was measured for the solutions of varying ionic strength and examined in terms of changes

in activity coefficients to be expected from eq. (5). The results indicated that the

changes in emf agreed within 0.3 mV with expectations based on the changes in ionic strength

and activity coefficients.

297
In addition, a comparison of data for solutions of Set 1 with those for similar solutions

of the same ionic strength in Set 2 showed that replacement of part of the Tris-H''' by Na"*" resulted in changes in emf no greater than 0.3 mV. Similarly, a comparison of data for Sets

1 and 3 at the same ionic strengths showed that the emf of cells F and G were likewise

unchanged at all but the highest ionic strength [11]. Thus there is strong evidence that

electrode selectivity parameters have little influence on the potassium and calcium measurements under conditions where a change in ionic strength of serum results from sodium depletion

or excess. This conclusion is important in measurements of serum calcium, for many commercial

calcium electrodes now in common use are subject to a significantly greater sodium error

than is the Simon electrode used in our study.

A further observation is of interest. In the measurements with the SCE reference

electrode it was found that the potential at the junction SCE| jSolution changes with ionic

strength in such a way as to counteract the activity effect. Since most practical measurements are made with calomel reference electrodes, this would tend to reduce the error in the

measured concentrations of ions below what would be expected from activity effects alone.

Probable errors caused by changes in ionic strength are indicated in" table 5. A positive error signifies that the value obtained from the electrode measurement is too high.

Table 5. Probable errors from differences in ionic strength

in cells with SCE standardized at ionic strength = 0.16.

I Na" Ca" error, % error, % error, %

0.12 +1.2 -1.1 +5.2

0.14 +0.8 -0.4 +3.0

0.16 0 0 0 0.18 +0.8 +0.4 -1.3

0.20 +2.0 +1.1 -2.1

6. Conclusions

While the evaluation of a single ion activities in biological fluids remains a problem

of considerable complexity, measurements in synthetic electrolyte mixtures simulating serum indicate that ion-selective electrodes respond in a near-Nernstian manner to changes in the

concentrations of the species for which they are selective. At a constant ionic strength of

0.16 mol 1'^, the activity coefficients and 1 i quid- junction potentials are reasonably constant. Effects of changes in ionic strength on the measured cell emf are relatively minor

and almost entirely accounted for by alterations in the activity coefficients. These results

lead one to expect that standards similiar to the ones described here may be well suited for

the calibration, on a concentration basis, of ion-selective electrodes for use in biological

or clinical media. The lanthanum fluoride electrode appears to be promising as a reference

electrode, provided account is taken of the effect of ionic strength on the activity

coefficient of the fluoride ion.

References

[1] Moore, E. W. , Studies With Ion-Exchange Calcium Electrodes in Biological Fluids: Some

Applications in Biomedical Research and Clinical Medicine, in lon-Selective Electrodes^

Durst, R. A., ed., p. 215, NBS Spec. Publ . 314 (1969).

[2] Lustgarten, J. A., Wenk, R. E., Byrd, C, and Hall, B. , Clin. Chem. 20, 1217 (1974).

[3] Rechnitz, G. A., Chem. & Eng. News, p. 29 (January 27, 1975).

298
[4] Guilbault, G. 6., and Montalvo, J. G., J. Amer. Chem. Soo. 92, 2533 (1970).

[5] Llenado, R. A., and Rechnitz, G. A., Anal. Chem. 45, 2165 (1973).

[6] Baum, 6., and Lynn, M. , Anal. Chim. Acta 65, 393 (1973).

[7] Eisenman, G. , The Electrochemistry of Cation-Sensitive Glass Electrodes, in Advanees

in Analytical Chemistvy and Instrumentation^ C. N. Reilley, ed., Vol. 4, p. 213

(John Wiley & Sons, Inc., New York, 1965).

[8] Bates, R. G., Pure Appl. Chem. 36, 407 (1973).

[9] Ruzicka, J., Hansen, E. H., and Tjell, J. C, Anal. Chim. Acta, 67, 155 (1973).

[10] Robinson, R. A., and Bates, R. G., Anal. chem. 45, 1684 (1973).

[11] Mohan, M. S., and Bates, R. G. , Clin. Chem. 864 (1975).

[12] Morf, W. E., Ammann, D., Pretsch, E., and Simon, W., Pure Appl. Chem. 36, 421

(1973).

[13] Durst, R. A., and Staples, B. R. , Clin. chem. 206 (1972).

299

National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

ELECTROLYTE ACTIVITIES IN HUMAN BLOOD PLASMA

M. J. D. Brand! gnd W. J. Scott

Technicon Instruments Corporation

Tarrytown, N.Y. 10591, USA

1. Introduction

Assessment of a patient's acid-base and electrolyte status is an essential diagnostic

procedure in many acute and non-acute diseased states. Diseases of the renal and respiratory systems primarily result in acid-base and electrolyte imbalance although other

regulatory mechanisms exist [1]^. Care of the critically ill patient may require frequent

measurement of blood gases and electrolytes during potentially life-threatening episodes,

e.g. 3 in cardiac surgery. The use of serum or plasma electrolyte measurements to evaluate

the condition of acutely ill patients reflects the fundamental importance of extracellular

electrolyte concentrations in basic physiological processes.

From a physico-chemical standpoint blood plasma is an extremely complex liquid. It

is moderately concentrated in inorganic salts and contains appreciable levels of polyelectrolytes (proteins). Uncharged species, e.g.^, CO2, glucose and urea, are present and

have important effects on metabolic processes. Because of the complexity of this chemical

system, the physician bases his clinical judgments on measurements of total electrolyte

concentration. Indeed, he has had little choice as the analytical methods in routine use

in clinical laboratories provide only concentration measurements for Na"*", K"*", CI", and

HCO3. The total concentrations of these species give little information on the true

composition of plasma, which is influenced by specific ion-ion interactions (metal-1 igand

complex formation) and non-specific interactions characteristic of non-ideal solutions.

The physical chemist prefers to define for each species present an "activity" which may

crudely be regarded as a concentration term, corrected for the various specific and non- specific interactions.

The complexity of human plasma electrolyte composition has been recognized most

clearly in the measurement of ionized calcium [2]. The preferred practical method of

making this measurement uses an ion-selective membrane electrode [3]. Such electrodes are unique among analytical sensors in that they have a response determined by ionic activity,

rather than total concentration. Routine serum electrolyte measurements in many clinical

laboratories are now made with ion-selective electrodes used in the Technicon Instrument

Corporation systems SMAC and Stat- Ion. These systems dilute the serum sample so that

ionic interactions are controlled, and the results are obtained as concentrations.

There have been several reported applications of ion-selective electrodes in the direct

measurement of electrolytes other than ionized calcium in whole blood, plasma and serum. These have included the determination of sodium and potassium [4-7] as well as descriptions

of more complete systems for acid-base and electrolyte measurements [8-11].

The use of ion-selective electrodes in whole blood, serum or plasma as opposed to dilute serum presents numerous problems, not least of which is an understanding of what is actually being measured. Rational interpretation of the results of electrode measurements requires a knowledge of the true nature of each ion species measured. Previously,

Walser [12] and Robertson [13] have considered ion interactions in plasma ultrafi Urates.

^Present address: Department of Chemistry, Texas A&M University, College Station,

Texas 77843.

2 Figures in brackets indicate literature references at the end of this paper.

301
Dahms et at. [14] have discussed the activities of sodium, potassium and chloride in human serum as determined by electrode. In this paper a model is developed for electrolyte ion

activities in human plasma. A complete model would include acid-base, metal-ligand

complex, and redox equilibria as well as those involving gas and solid phases. However in

order to contain the complexity of the problem, only a limited range of acid-base and

complex equilibria will be considered.

This model is intended to represent the extracellular fluid in equilibrium with

erythrocytes and other blood cellular components, i.e., the plasma fraction of blood. The

term plasma implies in vivo plasma rather than the in vitro counterpart which contains an anticoagulant. However, because many of the equilibria constants required have only been

measured at 25 °C, numerical results are obtained at this temperature; this does not detract from the accuracy of the model but only from its interpretation.

A. Water content of plasma

The solids, particularly proteins dissolved in plasma occupy a significant fraction

of the total plasma volume, i.e., the partial molar volume of proteins is not negligible.

An alternative statement is that the activity of plasma water is less than one. Activities may be given in two different scales; molar (mol/1 of solution) and molal (mol/kg of

solution). The molar scale is adopted here mainly because it is in agreement with normal

clinical chemistry practice (meq/1, mg/dl , etc.).

McLean and Hastings [15] have given the water content of plasma as a function of the

protein concentration (in g/dl):

Accordingly, the total concentration of all non-protein species is corrected from mol/1 of

plasma to mol/1 of plasma water using the factor given by eq (1).

B. Specific ionic interactions

This model is restricted to consideration of the clinically significant cations H"*", Na"*", K"*", Ca^"*", and Mg^"*" and their interactions with various anions; metal -protein interactions will be considered below. It is necessary to restrict the range of anionic species

considered; table 1 shows the anions which are thought to represent the majority of those

present in plasma. A great many others, e.g., amino acids, ascorbic acid, etc, could have

been included but were omitted to simplify the model. As a further simplification, not more than one acid-base dissociation was considered for each acid as shown in table 1; this is a reasonable restriction, from comparison of the acid pK with physiological pH

values. In general terms, the model considers the formation of a 1:1 complex between each

metal and each anion. No polynuclear complexes are considered, e.g., of the type M2L, nor are mixed complexes, e.g., M1M2L. Not all of the possible 1:1 complexes were included as for some the stability constant is too small to measure, i.e., the complex is completely

dissociated; table 2 lists the complexes included.

For each ligand:

2. Development of the Model

H^O/100 ml plasma = (100-1) - 0.75 P (1)

M. . L^^M.L^

4^4

K (3)

(2)

^ ^4—^4 K,2i " [M,-] [4]

(4)

302
Table 1. Plasma anions and ionization constants.

CT

CO2 +

H2PO4

HSO4

HCitrate"

Lactic acid

Pyruvic acid

HCO3 + H"^

HPO4 + H"^

SO4 +

Citrate"" + H"*"

Lactate" +

Pyruvate" +

6.35

7.21

1.99

6.49

3..86

2.49

Numerical values from [16].

'chloride is assumed to be completely ionized.

Table 2. Complexes and dissociation constants pK.^

HCO3 ^2^°4 HCit^" Cit^" Lactate" Pyruvate"

Na^ -1.11 -0.72 -0.93

K"^ -1.00 -0.96 -1.10

Ca^^ -1.26 -1.08 -2.70 -2.31 -3.09 -4.85 -1.42 -1.08

Mg^^ -1.16 -l'' -2.50 -2.25 -2.46 -4.57 -1.37 -1 (2)

^Numerical values from [13] and [16].

'^Estimated values.

where [ ] denotes activity. The total concentration of the ligand is:

^ot = ' W/^2 h4]/^M.L^ ' ? [^•4]/^M.L2 (5)

where y represents an activity coefficient. From (2) - (5):

n -1 = lliot (6)

303
[4] = mhv^ (7)

[M.4] = K^,[M.][LJ (8)

[M.4] = 4i[M,.][L2] (9)

Equations (6) to (9) can be solved given L^^^, H, M^. and the activity coefficients. The

values of L^^^ and H {i.e. pH) are assumed known, as are the total concentrations of each

metal ion. The mass balance condition for each metal ion is given by:

Solution of equations (6) to (9) is obtained iteratively using an initial estimate for

M^- and calculated activity coefficients to give an improved estimate for M^. from (10).

The method of calculating activity coefficients and details of the computer program used

to obtain the solution are given below.

C. Protein- ion interactions.

This model is restricted to a consideration of the interaction of calcium and magnesium with albumin and globulin. The effect of chloride binding to albumin, although

included in an earlier treatment [14] is not considered here as the magnitude of such

interactions has been shown to be negligible at physiological pH [17,18].

There have been many studies of calcium binding to albumin, the results of which have

not always been in agreement. Recently, Pedersen [19] has shown that under physiological

conditions, the concentration of bound calcium is given by

(CaAlb) "^l^^^ )

'^^^ot " il H- K,(Ca2+) . 10-PHK 1 ni)

I a

where = 10 » Ki = 241 1/mol, and ( ) represents concentration. This equation is used in the model. Pedersen [19] also studied the binding of calcium to y-globulin.

Although far less data is available for globulin binding, it was shown that the binding

was less and of smaller pH dependence. A linear regression to the data of table 2 [19]

gives the equation.

(CaGlob)/Glob. . = 0.564 pH - 1.79 (12)

at a mean ionized calcium concentration of 3.8 mmol/1 and a globulin concentration of 0.13

mmol/1. Pedersen has given insufficient data to estimate a reliable equilibrium constant

for this reaction and, therefore, the approximation is made here that under physiological

conditions the concentration of globulin bound calcium is linearly proportional to the

free calcium ion concentration.

Relatively few studies have been made of magnesium binding to albumin. A recent

treatment by Frye et at. [20] shows that the equilibrium follows a pH dependent mass law

304
relationship. The data of table 2 [20] can be reduced by linear regression to show the pH

dependence of the Mg-albumin complex stability constant:

= 142.65 pH - 952.79. (13)

Assuming six ion binding sites per molecule of albumin,

6 K,(Mg2"')Alb

(MgAlb) =

' . (14)

1 + (Mg^^)

No data has been found on magnesium binding to globulins and, therefore, the assumption is made here that calcium and magnesium bind identically to y-globulin.

The treatment used in this model assumes that calcium and magnesium do not compete

for the same binding sites on albumin or globulin. Data on the competitive binding of

these two metal ions to proteins does not appear tn be available. However, it must be

noted that under normal physiological conditions not all binding sites are occupied and

therefore the assumption of independent bindjng of calcium and magnesium may not be un- reasonable.

D. Activity coefficients

Activity coefficients for all uncharged species were assumed to be 1, as were the

activity coefficients of proteins and metal -protein complexes. The activity coefficients

for all other ionic species were calculated by means of the Davies equation [21]:

- logy. = 0.509 nj — 0.31 (15)

^ M + /r

where n-j is the ion charge and I the ionic strength. This equation takes no account of

the individuality of ionic activities nor is it particularly accurate for trivalent species.

However, only one such ion (citrate) was included, and at low concentrations, so the equation was considered adequate.

Following the convention of Scatchard [18], the ionic strength was considered to be a function only of non-protein ions but was otherwise defined conventionally;

1-0.5 En? - C. (16)

where C^. is the ion concentration.

3. Computer Simulation of the Model

Figure 1 shows a flow diagram of the computer program used for numerical simulation

of the model. The program is written in Tymshare Fortran IV; full details may be obtained

from the authors.

Input data to the program includes: pH; Na, K, Ca, Mg, total concentrations; albumin

and globulin total concentrations; the total concentrations of each of the seven anions

given in table 1. Initial values for the ionic strength (equal to the sodium ion concentration) and free metal ion activities are set. The estimates for the monovalent cations

are obtained from the products of the concentrations and the activity coefficients derived

from the estimated ionic strength. Estimates for divalent ion activities were obtained by

trial and error; for normal plasma these values were set equal to half the product of

concentration and the estimated divalent activity constant.

305
INPUT

DATA

SET INITIAL CONDITIONS

ACTIVITY

COEFFICIENTS

FREE

LIGANDS

METAL -

PROTEIN

COMPLEXES

METALLIGAND

COMPLEXES

FLOW CHART OF

PROGRAM 'PANIC

Figure 1. Flow diagram of computer program used to simulate model.

The program then proceeds to solve equations (6) to (9) to calculate the free ligand

and metal -ligand complex activities. Next, the concentration of protein bound metals are calculated using equations (11) to (14). At this point it is possible to make a better

estimate for the free metal ion activities using equation (10). The program now tests for

convergence by comparing the initial estimates. If these differ by more than 1 percent,

the free metal ion activities are reset to the mean of the initial and new estimates, and

the calculations repeated. In this iterative fashion a solution is usually quickly obtained to the set of equations.

4. Results and Discussion

A real advantage of a model of a system as complex as blood plasma is that it allows

the effects of variables to be investigated when these may be experimentally inaccessible.

Thus, it is possible to examine the properties of plasma when all the components take

extreme high or low values; clinically such a situation is extremely unlikely. Table 3 shows the calculated ionic strengths and activity coefficients in such hypothetical plasma

samples compared to normal values. Monovalent activity coefficients are in the range 0.76

± 0.02 and divalent coefficients in the range 0.33 ± 0.03. The constancy of the activity

coefficients over the extreme physiological range indicates that either an activity or a concentration scale may be used to obtain medically significant results. In the particular case of ion-selective electrode measurement of electrolytes, it is possible to calibrate the electrodes on a thermodynamic activity scale or on a concentration scale by

appropriate choice of standards at a physiological ionic strength [22]. It must be noted

that significant differences exist between concentrations expressed in terms of plasma

306
Table 3. Comparison of normal and abnormal plasma.

Input data (total concentration) Low Normal High

pH 7.0 7 4 8.0

Na meq/1 120 140 170

K meq/1 2 4 8 Ca mg/dl 5 10 15

Mg mg/dl 1 2 5 5 Albumin g/dl 2 4 7 8

Globulin g/dl 1 2 5 5

CI meq/1 70 100 130

CO2 meq/1 10 30 40

P mg/dl 1

- 3 7 10

S mg/dl 0.5 nu

Q 1.5

Citrate mg/dl 2.4 2. 4 2.4

Lactate mg/dl 10 10 10

Pyruvate mg/dl 0.6 0. 6 0.6

Calculated values

Ionic strength 0.108 0. 152 0.204

Monovalent activity coefficient 0.778 0. 759 0.746

Divalent activity coefficient 0.365 0. 332 0.310

volume and plasma water volume. Table 4 shows the effect for the four commonly determined

electrolytes. Calibration of an electrode with aqueous protein-free standards results in

plasma measurements on the plasma water volume scale.

The model predicts the activities of 46 components of plasma; table 5 shows the

calculated distribution of these in normal plasma. It is also possible to investigate the

effect of one variable on others, e.g.^ the pH dependence of ionized calcium and magnesium

(fig. 2). These results indicate possible applications of the model; however, it is not intended to explore these in detail here.

Table 4. Comparison of normal plasma electrolyte values

on different scales.

Na

K

01

CO.

Total concentration

meq/1 plasma

140

4

100

30

Ion concentration

meq/1 plasma HgO

149

4.3

107

30^

Ion activity

meq/1 plasma H2O

113

3.2

81

23^

HCO3

307
1 un 1^ ID CO LT) o O

1o

1 o

1o o

1o1o

1o O

1o

X X X X X X X X X X X

(Ti CO CM o tn

in CO CM 00 o CO un

<£< CO CM CM in

CD

o

CD

CD

+ -=3- + CO O O Q- O O CM Qn: ic 31

C7> CD

+-> + + o s- O >>

00 DC o _J

C7> C7> CD cn

s: SI oCM

co ID in

(/> 1 1 1 1 1 1 1 E o o o o o o o o c

(O 1 1—

r— X X X X X X X X X

Q- ^ vo 1 CM 1

in

CM CM CM o >=1- Ln CO O

CM >5l- 00

X

CM

X

-l->

a;o

Q-EO

o

10E

(/)

o

J ct CJ3 n3 fO fO O O C_)

COO

X

CM

1- CO o 1 + o o +-> u

<J> CM Q. o "3

zr. :jz nz O _J

10 ro 10 fO to o O o O o o

or—X

CT>

U3

IO

X

CMm

s->>

C_)

oo. o o

o

4-)

s-

!-> o

E

oo

cy> or—

Xo

ino

X X

in

J3

ro o I >=1- o

n3 O

30 o o CO CM ID CO in

1O1o

1 o1o

1 oIo

1 o

Io1o

1O1 o1o1o1o

1o

X X X X X X X X X X X X X X X «^ CM to o CTl CO o CTl in in 00 to CM 00 in o CD CO en CM in CM

Ln o CM CM CO CM CM CM E

E Z3 1 ^ II 3 1 CO o II "S- +-> III o s- ja o CM o D- o o II >=d- 4-> <0 o >> 's- o o CM Q. 00 O <_) _l Q. 5 5 o o n: nr oo JZ O zc _J a: a.

308
Figure 2. Calculated pH dependence of ionized

calcium and magnesium in normal plasma.

U.4

n 9

0 7.0 7.2 7.4 7.6 7.8 8.0 pH

It may be argued that the model is incomplete in that it does not require electroneutral ity--a condition which is always satisfied for an actual plasma sample. The ab- sence of this condition is intentional and, in fact, it is not required of the model.

Electroneutrality is obtained at a given pH by charge balance between the proteins and

other ions. Perhaps the most serious criticism of the model is that it relies entirely on published equilibrium constants, the accuracy of which is not always beyond question. If it serves no other purpose the model emphasizes the need for reliable thermodynamic measurements of these constants at various temperatures and at physiological ionic strength. It

is to be hoped that this treatment of ionic equilibria will lead to a clearer understanding

of the species present in blood plasma and their role in metabolic processes.

References

[1] Collins, R. D. , Illustrated Manual of Laboratory Diagnosis, Second Edition, Chapter

4, p. 35 (J. B. Lippincott Company, Philadelphia, 1968).

[2] Moore, E. W. , Ionized calcium in normal serum, ultrafiltrates and whole blood determined

by ion-exchange electrodes, J. Clin, invest. 49^, 318 (1970).

[3] Weissman, N. and Pileggi, V. J., Inorganic ions, in Clinical Chemistry Principles and

Techniques^ 2nd Edition, R. J. Henry, D. C. Cannon, and J. W. Winkelman, eds.,

p. 662 (Harper and Row Publishers, Maryland, 1974).

[4] Moore, E. W. and Wilson, D. W. , The determination of sodium in body fluids by the

glass electrode, J. Clin, invest. 42^, 293 (1963).

[5] Friedman, S. M. , Wong, S., and Walton, J. H., Glass electrode measurements of blood

sodium and potassium in man, J. Applied Physiol. ]%_, 950 (1963).

[6] Portnoy, H. D. and Gurdjian, E. S., Glass electrode determination of blood sodium and

potassium. Am. J. Clin. Path. 45^, 283 (1966).

[7] Miyada, D. S., Inanu, K. and Matsuyama, G., Direct potentiometric determination of

potassium and sodium in blood, plasma and serum, Clin. Chem. ]7_, 27 (1971 ).

[8] Dahms, H., Automated potentiometric determination of inorganic blood constituents

(Na+, K+, H+, CI), Clin. Chem. U, 437 (1967).

309
[9] Neff, G. W., Radke, W. A., Sambucetti, C. J. and Widdowson, G. M., A computer-assisted

electrode system for measuring blood pH, PO2, PCO2, sodium and potassium, Clin. Chem.

1_6, 566 (1970).

[10] Orion Research Incorporated, Automated Potentiometrio Eleotrolyte Analysis System,

Final Report, Contract NAS9-12117 (National Aeronautics and Space Administration,

Houston, Texas, December 1973).

[11] Steele, R. E., Walker, W. E., McWhorter, M. M., Lewis, K. A., Robertson, C. R.,

and Maffly, R. H., Critical care blood gas, pH and electrolyte analyser, Pvooeedings

of the 27t'h Annual Conference on Engineering in Medicine and Biology, Vol. 16, p. 401 (The Alliance for Engineering in Medicine and Biology, Chevy Chase, Maryland, 1974).

[12] Walser, M., Ion association VI. Interactions between calcium, magnesium, inorganic

phosphate, citrate and protein in normal human plasma, J. Clin. Invest. 40, 723

(1961). ~

[13] Robertson, W. G. , Measurement of ionized calcium in biological fluids, Clin. Chim.

Acta, 2A, 149 (1969).

[14] Dahms, H., Rock, R. and Seligson, D. , Ionic activities of sodium, potassium and

chloride in human serum, Clin. Chem. 1_4, 859 (1968).

[15] McLean, F. C. and Hastings, A. B., The state of calcium in the fluids of the body.

I. The conditions affecting the ionization of calcium, J. Biol. Chem., 108, 285

(1935).

[16] Sillen, L. G. and Martell, A. E., Stability Constants of Metal-Ion Complexes, Chemical

Society Special Publication No. 17 (Metcalf and Cooper Ltd., London, 1964).

[17] Carr, C. W., Determination of ionic activity in protein solutions with collodion

membrane electrodes, in Electrochemistry in Biology and Medicine, T. Shedlovsky

ed., p. 266 (Wiley and Sons, Inc., New York, 1955).

[18] Scatchard, G. and Yap, W. T. , Physical chemistry of protein solutions. The effects

of temperature and hydroxide ion on the binding of small anions to human serum albumin,

J. Am. Chem. Soc, 86, 3434 (1964).

[19] Pedersen, K. 0., Binding of calcium to serum albumin II. Effect of pH via competitive

hydrogen and calcium ion binding to the imidazole groups of albumin. Scand. J. Clin.

Lab. Invest. 29, 75 (1972).

[20] Frye, R. M., Lees, H. and Rechnitz, G. A., Magnesium-albumin binding measurements using

ion-selective membrane electrodes, Clin. Biochem. 7_, 258 (1974).

[21] Davies, C. W., Ion Association (Butterworths , London, 1962).

[22] Bates, R. G. and Mohan, M. S., Studies in the calibration of ion-selective electrodes

for use in biological fluids, Clin. Chem. 21_, 864 (1975).

310
National Bureau of Standards Special Publication 450. Proceedings of a Workshop on pH and Blood Gases held at NBS, Galthersburg, Maryland, July 7-8, 1975. Issued June 1977.

THE KING'S COLLEGE HOSPITAL ION-SELECTIVE ELECTRODE SERUM

ELECTROLYTE ANALYZER

A. D. Hirst, P. Gay, P. Richardson, and P. J. N. Howorth

Department of Chemical Pathology

King's College Hospital Medical School

London, U.K., SE58RX

In recent years there has been a great increase in the availability of ion-selective

electrodes, and these are finding wider use in routine clinical chemistry laboratories

(table 1). The obvious difficulty is that such electrodes measure activity which must be

somehow related to concentration by calibration and conversion factors. Early reports,

e.g., with the Na electrode suggested that the conversion factor might vary in disease

states when abnormal concentrations of protein were present [1]^.

Table 1. Some ion-selective electrodes.

Cations Anions

h"" (pH) F" CI"

NH3 (NHt) Br" r

Na" k" CN" NO3

Ca2"^ Cd2"^

Cu2"^ Pb2"^

We felt that it would be an interesting and potentially useful experiment to construct

a multi-channel continuous-flow serum electrolyte analyzer based entirely on ion-selective

electrodes. This would dispense with the need for chemical reactions to first make a colored

solution, pass light through it, measure what emerges and then by complex electronics, convert

it to the concentration of the substance being measured.

Table 2 shows the general principles adopted in the analyser. The Na electrode is probably the most critical part of the system. The Na electrode is based upon an alkaline

Table 2. General outline of KCH analyzer.

Ne" electrode: acid-serum dialyzed into high-osmolar

alkaline buffer

k" electrode : Na" activity backed-off

HCO3 : pH electrode CO2 from acidified serum

diffuses into Na2C03

Glucose : O2 electrode. High concentration glucose

oxidase causes quick fall in PO2

Urea : NH3 electrode. Urease NH3 from urea

Figures in brackets indicate the literature references at the end of this paper.

317
error in a glass pH electrode. If serum is made alkaline with buffer, then protein interference can be a problem. By dialyzing an acidified serum sample into an alkaline solution

of high molarity, a linear response to changes in Na concentration can be obtained. Figure

1 shows the flow- diagram for the analyzer.

waste

t

dialyzer

Na"*" electrode TT

K"*" electrode

waste

1

dialyzer

H"*" felectrode

NH3 electrode 4

s ample line

ai r 0.2 mol HaSOit

1 mol NHijOH/CI

air

1 mol triethanol amine

air HCl pH 6.7

ai r 0.2 mol HaSOit weak Na2C03'

ai r

I ^0.5 mol NaOH

ureas e/EDTA

.air pH 6

O2 electrode qgl ucose oxi dase

I 0.15 mol NaCl

-ai r

Figure 1. Flow-diagram for KCH Ion-selective electrode serum electrolyte analyzer.

The K-selective glass electrode is responsive to an appreciable extent (activity

ratio about 10:1) to Na ions, but fortunately the Na response is linear and additive.

Accordingly, K is measured after Na by means of time-delay coils and the Na contribution

from the Na electrode is electronically backed-off from the total ion activity. The serum

HCO5 measurement* is very simple. About 50 percent of the CO2 liberated from acidified

serum can be dialyzed across a standard membrane into a Na carbonate buffer whose pH

alters with varying CO2 content.

The measurement of glucose by an O2 electrode gave trouble at first owing to the O2

content in the air-segmentation bubbles in the mixing coils. It was not found necessary,

however, to use inert gases for this purpose. By using a high concentration of glucose

oxidase, the rapid fall in PO2 found in the first 30 seconds was found to be proportional

to the glucose concentration. The urea method is potentially as accurate a method as it

is possible to devise and should be superior to the non-specific diacetyl monoxime reaction.

312
The urease reaction takes place at a slightly acid pH after which the sample is made

highly alkaline (pH 12) before being fed through the NH3 electrode.

The electrolyte analyzer has been assembled at King's College Hospital from commer- cially bought electrodes, sampler, pump, dialyzers, mixing coils, heated water bath,

recorder and electronic components. Unfortunately at the time of writing, it is not yet

in routine operation although all the individual channels have been worked up prior to the

final design of the machine. Figures 2 and 3 show general views of the analyzer.

Figure 2. View of KCH analyzer.

Figure 3. Close-up view of analyzer.

313
All the arguments that have raged for 60 years about the measurement of H concentration or activity apply in some measure to ion-selective electrodes. However, since no one

in practical terms really wants to measure or report urea as "PNH3" or K"^ as "pK" [2],

then we will all make the necessary effort to adapt in a reasonable manner the activity

measurements to concentration units. All our routine methods have inherent errors or compromises of one sort or another. We really do not see any good reason at all why we cannot also make an effort with regard to the pH-glass electrode and use a single unified

set of units such as the SI provides for all our common electrolyte and blood gas analyses.

References

[1] Dahms, H. , Rock, R., and Seligson, D. , Ionic activities of sodium, potassium, and

chloride in human serum, Clin. Chem. ^4, 859 (1968).

[2] Howorth, P. J. N. and Hirst, A. D. , SI units, pH and cH"^, J. Clin. Path. 28, 423 (1975).

I

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WORKSHOP DISCUSSION

The transcription of selected segments of the discussion sessions is included in this

volume because it supplements much of the material in the preceding chapters. In this chapter will be found additional comments on practical problems, instrumentation, methodologies,

and concepts on the subject of blood pH, gases and electrolytes.

The workshop was arranged to provide time for discussion at the conclusion of each

series of lectures on a particular subject. The following discussions were transcribed with

a minimum of editing and represent the spontaneous exchange of remarks of the various participants as they were recorded at the workshop. The speakers have had an opportunity to correct the transcription of their remarks for technical accuracy.

Bates :

Before this session began, several of us were reminiscing about the meeting in 1964 at the New York Academy of Sciences and I recall in 1964 there was a great deal of discussion

about two things, one of them being the validity of so-called "unlogging" the pH value, and

the other one was concentration versus activity, in other words, concentration versus excess

free energy. It seems rather remarkable that in eleven years the utility of one of these

hasn't proved itself. We still have the same problem of which to use, and it occurs to me

that this might be due to the fact that the experimental methods for determining pH, for

example, do depend more closely on an activity than on concentration, and I wonder if this

is the reason that this dilemma still faces us.

Cohen :

My view of this is that this is not a dilemma but a sematics difference. When hydrogen

ion concentration is spoken of, at least in the parlance of clinicians and physiologists, I think what is meant is "the negative antilogarithm of the pH" which is, as you point out, an activity measurement. This is not to say that one is ignoring or brushing aside the difference between concentration and activity, but that one is simply using a shorthand notation

to designate measured quantity in terms that are more readily conceptualized. And I am wondering whether you would view this as an acceptable "derivation" from the measurement of

pH or whether we ought employ a different term for it.

Bates :

Well, it seems to me that we are really begging the question of whether it's an activity

or a concentration that is crucial in biochemical and biological processes, aren't we?

Cohen :

Well, that is what I mean to imply. That is, I don't think anybody really quarrels with

the notion that it is the activity that is fundamentally at issue, that biological systems

and chemical systems must be responding to the chemical activity. That being the case then,

I think it's a matter of arbitrary definition as to what one choses to call the unit of measure as long as it serves the need for unambiguous communication.

Austin :

I think that perhaps it may be the fact that pH is essentially a "dimensionless" factor

that is important. pH serves as an indicator in a given situation, and it avoids the issue

of applying other than a unit dimension. Is it a nano-equivalent? How does that fit with

a milli-equivalent? I think that is one important factor, and it speaks for the continued

use of pH. We are dealing with rather gigantic differences when we're talking about milliequivalents versus nanomoles. One can not argue that we should then speak of the log of the

sodium or potassium concentration, since we administer these in their actual concentrations,

even though their biological activity is what is important. Though we give bicarbonate in

mEq to correct hydrogen ion concentrations/activities, no one corrects acidosis, for example,

by hanging a bottle of H+.

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We is berg :

I have two thoughts on pH versus the hydrogen ion concentration. Physicians have degraded themselves by saying they can't understand pH, whereas chemists as a group use pH,

We're not going to be able to change the rest of the chemical field by saying we're going to use hydrogen ion concentration. Secondly, I find that those who speak for the hydrogen ion

concentration are really deluding themselves because a change of 10 nanomoles in a hydrogen

ion concentration has a completely different interpretation when it's on the highly alkalotic

side versus the highly acidotic side. So you're going to give a greater importance to one nanomole of hydrogen ion where the effect on the body enzyme activity, and so on, is com- pletely different. So my own feeling is that I give both pH and hydrogen ion concentration

in my report on the diagram; you can use whichever you will. But if you realize, as most

of us do, that a pH unit change is a tenth or ten times going the other way, we can get that

concentration effect the same way. We're dealing with numbers, but we're giving undue importance to a number of the hydrogen ion concentration.

Cohen :

Well, I think one could find as many examples of biologic relationships that would be

linear with respect to hydrogen ion concentration as you could with respect to pH. I don't

really think the argument that enzyme systems or other things are responding linearly to pH

rather than to hydrogen ion concentration or activity is a very compelling one. I think the

only compelling argument, at least in my mind, for retaining, or perhaps substituting, the

notation, "hydrogen ion concentration", (with the caveat that I mentioned earlier) is that

it simplifies the mathematics that describe the bicarbonate-carbonic acid buffer system.

That is the only justification that I can see, but it seems to me that is one that can't be

minimized. It has some very useful pedagogical and practical advantages.

Engel :

I understand Dr. Bates' surprise because it has taken a long time since we started these

discussions at the meetings at New York Academy of Sciences. I think also that you should

realize that the people working in this field of acid-base physiology are a very inhomogeneous

group of people. We have pure chemists, we have physiologists, we have medical doctors in

the ward, etc, and the interests of the different groups of people are very different and,

therefore, the emphasis is sometimes put on one side of the aspect. The other thing I would

like to comment on is the old questions around the "great trans-Atlantic acid-base debate"

which was brought up again. I don't think it has ever stopped even though it sort of dampened

for a few years. I think, however, that we should agree to divide our problems in three different groups. First, we have some basic definitions which apply to clinical chemistry as well as to pure and applied chemistry. I think it's very feasible to use these very well

systematized and well defined concepts. Second, as clinical physiologists and as acid-base

physiologists, we might be interested in some derivatives of these well-defined and measurable

variables and, of course, we would like to define some variables or derived variables which

make physiological sense. And thirdly, we have the problems related to the clinicians with

pathophysiological interpretation of these either derived or basically measured values. I think that we have been very good in the past of mixing all three of these groups together in one melting pot, and that lots of the questions which were discussed during the "great transAtlantic acid-base debate" came about because of this maxture of what our different goals

and definitions were.

Howorth :

I'd like to go back to something that people were talking earlier on about pH as opposed

to so-called concentration units. I think many physicians in practical terms need some sort of concentration unit. I mean if a surgeon is aspirating gastric juice, he likes to have some measure of how much acid is being removed from that patient. You also need a quantitative measure with renal excretion of H+ and renal function tests. In everyday practical medicine we need to know how much acid a patient is putting out in his urine on a quantitative

basis. This is where the mole as the amount of substance is superior to the dimensionless

pH concept.

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Cohen:

Is it not true that in other systems where the measurement is done electrochemically,

for example, calcium, we report the substance in concentration units when we recognize that we are measuring it electrochemically? I don't think that poses a problem.

Noonan :

Well, I think a fundamental concept in analytical chemistry is the idea of recoverability. I can say that I am measuring glucose quantitatively if I can recover a certain amount

of pure glucose which has been added to it. The glucose technique is directly measuring the

concentration. However, I cannot take a certain quantity of hydrogen ion and add it to a solution and measure it with an electrode and say that it is concentration because concentration is not being measured. In terms of ion-selective electrodes for sodium and potassium,

and indeed pH, you can get some constant correlation between the activity and the concentration that will work in many, many cases but will not work in all cases. If you really want

to be precise, there should be a difference between a sodium measurement that's made with an electrode and a sodium measurement that's made with a flame photometer.

Austin :

I agree, sodium concentration is important when you are replacing it, but with pH, when

considering dissociation equilibria (Davis,' R. P., Amer. J. Med. 42^, 159-162, 1967), the quantity of hydrogen ion that you're measuring is so much smaller that the quantity of sodium.

They are of an entirely different magnitude. This may create some difficulties with measuring

diminutive quantities of hydrogen ion where we're measuring an entirely different number of

sodium and potassium ions. How does this help you? Apart from concentrations, pH denotes

a more meaningful description of the acid-base equilibrium reaction. If we are talking

about acid-base balance that's one thing; ionic concentrations or activity are another matter.

Recall in the Henderson-Hassel balch equation that we do take the log of the bicarbonate con- centration, even though we give bicarbonate in its sodium form for practical purposes.

Gambino :

I think I'd like to say that I agree basically with Siggaard that it is preferable to go

to some type of measurement or some type of indication of the amount of energy and the quantity of substance, but I'd like to have some measure of the energy. I think it's important

to point out that we cannot utilize the suggestion that the way we make the measurement

should determine the unit. For example, in temperature we estimate the average kinetic

energy, we measure the movement of a mercury column and we measure current in a system. But we don't attempt to say that because we measure current that we should then define temperature

as an estimate of kinetic energy in terms of current. We don't do that. We go back to some fundamental quantity such as the joule. I think we should try to get a quantity into these

fundamental units.

Durst :

There's something that I picked up from our discussions that I might just mention now. That is that we're using the word "derived" in two different ways. One is in reference to the SI system where derived means units expressed as products or ratios of the base units

without numerical factors as opposed to what we are calling derived units such as bicarbonate,

that is to say, a parameter calculated from a measured parameter or parameters. Again, this

in one more example of a semantic problem which could be significant.

Weisberq :

May I add one more semantic problem. You're using the word nomogram erroneously. I have collected 105 diagrams, 17 nomograms, and 12 slide rules. The first diagram goes back

to 1914, and I think we are really referring to diagrams rather than nomograms. Dr. Malenfant

had a true nomogram. A Cartesian nomogram and a diagram are two completely different things.

A diagram does have benefits for a practitioner to see, especially with continuing sets of

data. One set of data means nothing; a point in time of compensation or uncompensation. But

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if you can tell where the patient has been or where the patient is going and thus follow that

on a flow chart. Dr. Siggaard-Andersen had, of course, the time chart going along with that.

So I think we ought to be clear that we are speaking of diagrams most of the time for physicians to look at. The nomogram is a laboratory device to save ourselves some mathematics.

Howorth :

I think a nomogram is something that one uses to derive a figure from primarily observed

data. A parameter isn't something which is directly measured, it's something that's indirectly measured.

Vi sser :

Well, there will be no confusion in this case between "derived" and "derived", because

base excess is a derived quantity, but maybe another derived quantity can be expressed in a basic unit. So we must distinguish between the quantity and the unit.

Cohen :

I'd just like to comment on Dr. Engel's presentation. I had two commentaries to make; one trivial, really, and the other perhaps more substantive. The trivial one being that the use of the term "net acid" in his scheme, if I understood it, is somewhat at variance with

modern definitions of acids and bases; it doesn't seem to conform to a proton donor. The

second is the issue of simplicity. I think there is nothing internally inconsistent about

that system for describing, in phenomonological or empiric terms, what happens to living

organisms when they are challenged. The only quarrel I have with it is that it seems to me

to be unnecessarily compl icated. One of the issues that surfaces in my mind, as a dominant

theme, relates to the fact the system with which we are dealing is inherently complicated,

simply by virtue of its physiology being so multifactorial. We are, therefore, better served

by the conceptual framework for discussing that system's behavior which is as simple as we can make it. I don't see any compelling justification or reason for a system, such as described, in contradistinction to one which deals just with the carbonic acid/bicarbonate

equilibrium and describes what happens to the organism in those very simple terms.

Engel :

If I can take the last comment first, then if you did total body studies of acid-base

metabolism, you would very quickly run into the same problems which we have been running

into. That is, you can't describe them with the set of variables we have been using so far.

When you define the body as a system, I think you should make it just as complicated as is

needed in order to completely describe the system. This is exactly why we wanted divide titratable acid up in three components: Carbonic acid which is regulated by the lung-brain

system, metabol izable acid which is regulated by the metabolism, and the net acid for which

the only route of escape is the kidney, and this makes physiological sense. I don't know

whether the scheme for the breakdown was too complicated or not, but anyway we feel that this

is the only way where you can go in and put a finger on the physiological mechanisms and causes which are involved in an acid-base disturbance. One thing more is that you may measure, as I said, a normal acid-base status in the blood and yet have a very pronounced acidbase disturbance in terms of balance studies.

Cohen :

I don't think there is any question about this latter point. Clearly, we can have a very markedly disturbed external balance without having the deviation.

Engel :

Right, but you won't find out unless you carry out the balance measurements.

Cohen :

Of course, but let me just take those two issues in turn. One is that I think it is possible to designate external balance of acids and bases without utilizing that breakdown

318
factor of acidity in the classic context of the net acid balance. In other words, there is

a certain rate of net acid excretion, by the kidney, which we all recognize as ammonia plus

TA minus the bicarbonate that is excreted. In the steady state, this balances the rate of

endogenous acid production. Now, there is no satisfactory way, nor do I think your system

provides me, for quantitating directly the rate of endogenous acid production. Given a diet, one can't use that scheme, for example, to specify what its net contribution of hydrogen ion or base will be to the body fluids, any more than one can take a diet and analyze

it for this purpose in any other way. So one is left with a conundrum that's been ^n issue

in physiology from the beginning of time; that is, unless one constructs a purely art.-^icial

diet which can be demonstrated to be complete metabolized, a la Relman, external hydrogen ion

balance must be estimated indirectly since we have no handle on the rate of endogenous acid

production. For this purpose, we use the rate of net acid excretion by the kidney, measured

in terms of hydrogen ions which are gaining access to the urine, as a fix on that parameter.

Engel :

Well, I think that you can assess the endogenous net acid production. But there is a discrepancy between what we and you call net acid. Net acid in our terms is defined as the

titratable value (TA) to an endpoint at pH 7.40 and temp. = 37 when carbonic acid (CA)

and metabol izable acid (MA) concentrations are zero. This is the operational definition

and it gives you values which are quite different from the net acid values you are talking

about. NA can be measured in food, you can measure them in feces, and you can measure them

in the urine, etc.

Cohen ;

I certainly don't quarrel with the fact that it can be measured, but the issue is to what end, what use can be made of the information once available. Let me, by way of illustrating that, ask you to analyze the in vivo acid-base disturbance that I illustrated in

my talk in terms of the components of your "titratable acidity" and to reconstruct the se- quence of events that led to that situation.

Gambino :

That's just the question that I was going to ask. My analysis of what you said, based

on what Dr. Engel said, gives a different interpretation of it. You said your paper described

the effects of hypocapnia. But I saw it as a secondary response to using your model to different types of acid load. You had an oral acid load, increase in hydrogen ion load absorbancy, and you had 9 percent oxygen, which I would suspect to be a primary anoxic acid load

with lactic acid.

Cohen :

No, Dr. Gambino, there was no lactic acid production in that experiment. That degree

of hypoxia produces no increase in lactate or unmeasured anion concentration at all.

Gambino :

How do you explain the bicarbonate?

Cohen :

That was the consequence of the renal response to the hypocapnia; the reduction in

bicarbonate reabsorption at the renal level resulted in a steady-state reduction in bicarbonate concentration, unaccompanied by any change in unmeasured anion concentration, no measured

lactate accumulation. In other words, the decrement in bicarbonate was offset entirely by an increase in chloride concentration and the change in bicarbonate being accounted for by a reduction, by my terminology, of net acid excretion. There was an accumulated change in acid

excretion by the kidney, that led to retention of endogenous acids that reduced bicarbonate

to that extent. So what one is seeing in that experiment is the physiologic response of the

kidney to two independent variables. One, the oral hydrochloric acid load, and two, the

effect of the hypocapnia on renal bicarbonate reabsorption.

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Weisberg :

If what you say is true, and I'm not doubting that, if you have that with a normal dog

without having the previous hydrochloric acidosis and have the same delta that you had with

the decrease in the PCO2 with the hypocapnia, why shouldn't you have the same effect in the

kidney where the loss of the bicarbonate is leading to that acidosis?

Cohen :

We find the same retention of hydrogen ion in the two situations, whether one goes from

normal to the chronic hypocapnic state or from the acid-fed to the chronic hypocapnic state. The difference is in the initial level of plasma bicarbonate. In one, it is normal, and

in the other it is low due to the acid feeding, but we are superimposing the same delta bicarbonate on these two different baseline levels; as a result, we see the divergent responses

of hydrogen ion. There is nothing magic about this. Given the change in PCO2 and the

physiologic response of the organism thereto, the change in acidity is predictable from the

relationships described by the Henderson equation.

Siggaard-Andersen :

It should be kept in mind that the same end result is obtained if (1) the initial disturbance is an acute respiratory alkalosis which is subsequently modified by metabolic com- pensation ending in a chronic respiratory alkalosis, and if (2) a metabolic acidosis is

initially induced before the respiratory alkalosis develops.

Cohen :

I certainly agree with that. I don't think the sequence in which these two maneuvers

are employed is crucial. One could reverse them and end up in the same situation, I would

predict. But we haven't done that experiment. But I still think that the issue is how does one unravel such complex situations. What is the conceptual scheme one utilizes to dissect

out the two components which are in that acid-base disturbance; it would be instructive to see how Dr. Engel would do that.

Engel :

Yes, I should first like to draw a parallel to that. We were at one point measuring

acid-base balance on patients with pyloric stenosis and, as you know, the interpretation of

the concurrent alkalosis has always been that, they developed alkalosis, because they are losing hydrochloric acid in the vomit from the stomach. Now making, for five or six days,

balance studies on patients with pyloric stenosis showed that the alkalosis only partly was due to a change in the net acid balance of the patients, and that most of the deviation in

Base Excess could be related to the fact that the patient at the same time lost large volumes

of water. Thus, it was rather a "contraction" alkalosis. I think that if we want to bring

about new ways of assessing acid-base metabolism in practice, and if you want to learn new

things, you cannot avoid measurement of external acid-base balances.

Siggaard-Andersen :

One of the problems is that the lactate ion behaves chemically as a neutral substance

when we try to titrate it to pH = 7.4. Nevertheless, we prefer to consider it a base from

a cl inical -physiological point of view. In Engel 's terminology, lactate is called a metabolic organic acid (MOA) with a negative sign. I prefer to call it a metabolic organic base,

in other words to distinguish between metabolic acids and metabolic bases, in order to avoid

talking about a component with a negative sign. The negative sign should not refer to the

component but to the measured quantity: the change in the amount of substance of the com- ponent in the organism or the body fluid.

I would like to use this opportunity to make a plea for the SI system. It is important

that the SI system, which has been advocated quite generally for science and technology, is also adopted by the clinical sciences. It is important that those of us who are involved in the measurement of physical and chemical quantities know the terminology of physics and

physical chemistry which constitutes the basis of clinical chemistry and clinical physiology.

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However, we cannot expect all physicians to remember all their physics and chemistry, and

therefore it is our responsibility to translate the information contained in the clinical

chemical data into common language.

The acid-base data including electrolytes represent a complicated set of data, perhaps

even including redundant data {e.g.^ total CO2, bicarbonate, base excess, buffer base). Much

of this information may be contained in "diagnoses" such as: partly compensated metabolic

acidosis, hypokalmeic alkalosis, etc. We may have to make slightly longer statements but

anyway this set of terms (acidosis, alkalosis, etc.) is very useful as a supplement to the mere list of measured quantities.

Engel :

I agree completely with Siggaard, and I think that there is one pitfall we should avoid,

and this is to let our instruments mix interpretation into the measured quantities or into

the derived quantities. I think the manufacturers should stay far away from that. Interpretation is the business of the clinical physiologist or the clinician dealing with the patient

Bates :

I'd like to ask Dr. Austin a question. Was your pK' a concentration constant or was it

the same as Dr. Maas' constants?

Austin :

It was the pK arrived at by determining the pH of whole blood, fixing the PCO2 by

known gas concentrations and measuring the CO2 content by the Van Slyke manometric technique.

Bates :

In other words, the hydrogen ion is supposedly an activity but the others are partial

pressures.

Austin :

Right.

Bates :

They are remarkably constant.

Cohen :

Harkening back to Dr. Bates' first question from this morning, I would like to ask

Dr. Maas if he would have any objection to simply taking the negative antilogarithm of the

equation which you put up, I've forgotten what the numbers were, I think 16 or 17 or something of that series. In any event, would he be willing to express the hydrogen ion con- centration, in quotation marks, as being equivalent to the ratio of PCO2 and bicarbonate

ion concentration multiplied by the negative antilogarithm of your pK value. In this way,

simply converting into linear units of activity or concentration, the logarithmic units of

pH, in order to specify the level of acidity.

Maas :

You can consider an equilibrium between gaseous CO2 and bicarbonate directly without

looking at the H2CO3 in between. So the answer to your question is, yes, the ratio PCO2 and

bicarbonate.

Noonan :

I wonder if we can get some discussion on this whole idea of talking about the plasma

pH, plasma PCO2 versus the similar quantities in whole blood. Does anyone have any feelings

about which it should be, or should we have one, or should we always talk about plasma pH in

measuring blood pH.

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Austin:

The difference is about 0.01.

Weisberg :

The difference is that we now have convenience. The machines usually use whole blood;

remember, we used to use plasma. Plasma is wonderful because you never clog the capillary

electrodes. When you have whole blood, you have all the problems. But we're lazy. Machines

are made that way, so it's easier to put in whole blood. And, of course, it's easier mechanically, but technically, as Dr. Austin pointed out, we're measuring the same thing. So it

makes no difference technically. Clinically, that 0.01 difference between plasma pH and

whole blood pH is not going to make any difference. And there's no major difference between

the PCO2 in plasma and whole blood as far as clinical purposes are concerned. I don't see you point, but it is interesting.

Gambino :

We don't have thermostated separation systems. So you have serious temperature errors

in pH but even more so in oxygen and PCO2. The other factor will come out tomorrow. These

data are utilized by anesthesiologists and ventilation therapists, etc., for the immediate

treatment of patients within minutes of having the results. Whole blood is essential. It's not just a matter of being lazy.

We were discussing pH and base and this is a workshop on blood gases and most of my

blood measurements are performed because there is a need to know CO2 gas pressure as a measure of ventilation and the partial pressure of oxygen as a measure of diffusion. These mea- surements are used to change ventilation therapy. And that's the most common cause for the

high volume testing, and the acid-base part of it is important but not as dominant in the

minute by minute therapy in the intensive treating. So I think we will have to also consider

the fact that there is a need for gas measurements for ventilation therapy.

Laver :

I prefer not to address myself to that problem. Dr. Rispens, to be sure I understood

you correctly, did you say that the ratio of the CO2 content of plasma to that for whole

blood was constant over a specific temperature range? Did you measure the CO2 content of

whole blood?

Ri spens :

In anaerobically stored blood the ratio of total CG2 in plasma to total CO2 in blood is

constant. I measured total CO2 both in plasma and in blood. I emphasized that it was not influenced by the temperature at which centrifugation of the blood took place.

Laver :

Did you say whole blood? In other words, the CO2 content of plasma plus and red cells?

Rispens :

Yes.

Laver :

So, in other words, the ratio of the CO2 content in plasma to the CO2 content of whole

blood was constant over a large temperature range.

Rispens :

I emphasized that it was not influenced by the temperature at which centrifugation of

the blood took place.

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Laver:

Let me start over again because I think it important we do not misunderstand. The

patient is at a low body temperature. CO2 content of whole blood rises more rapidly than

the CO2 content of plasma. If blood is equilibrated at a low temperature, let us say 25 °C,

and we now measure the CO2 content of whole blood and plasma, what happens to the ratio between the two as compared with equilibration at 37 °C?

Rispens :

Well, we measure the pH at 37 °C and put that in the equation [i.e., equation (2) in the

paper "Quantitative relationships ..."] to calculate the ratio between total CO2 in plasma

and total CO2 in blood.

Laver :

Let us assume we measure it by your technique, that of Van Slyke, or with the pH and

PCO2 electrode. I still keep coming back to the plasma versus whole l?lood CO2 content.

Austin :

CO2 content is constant regardless of the temperature. Dr. Severinghaus and Dr. Nunn

have shown that.

Laver :

I am sorry to be persistent, but you have not understood my question. If you equilibrate whole blood at a certain PCO2, for example 40 mm Hg at low temperature (25 °C), either

in a patient or in a tonometer, measure the CO2 content of whole blood as well as the CO2

content of plasma and compare that to the content for each when blood is equilibrated at

37 °C, what happens to the ratio between the two?

Rispens :

I said that in anaerobically stored blood the ratio between CO2 in plasma and CO2 in

blood is constant. Thus if you manage to get in the blood the same pH at 37 °C by equilibrating at 37 °C with a higher PCO2 than you used at 25 °C, the total CO2 concentration will

be the same and so the ratio.

Laver :

Yes, you keep saying total CO2. Yet CO2 content in the red cell and CO2 content in

plasma refers to two different things.

Gambino :

You've come back to something reported many years ago, namely, that whole blood with 15

grams of hemoglobin and with a changed temperature, you don't get any change in pH with

changing temperature. Let me draw it on the board. This is invariant with the temperature

when it is whole blood, but its not invariant with the temperature when it's plasma. I think

that's what you're driving at, right? If the temperature changes in a whole blood system,

the solubility of CO2 gas causes changes in one direction and the ionization of the proteins,

including hemoglobin, changes in the opposite direction so that the net pH remains constant.

Rispens :

That is certainly true, but that is not the point I wanted to make. I took a blood

sample, centrifuged it at 20 °C and 37 °C and found that the ratio of CO2 in plasma and

CO2 in blood was not affected.

Howorth:

It wasn't entirely clear about the method for measuring total CO2. Do you always use plasma or do you sometimes use anaerobic blood?

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Rispens :

I use blood.

Howorth:

One of the common techniques used in England is to separate the serum or plasma and then measure total CO2 by acidification extraction using the Technicon AutoAnalyzer. You can get

serious losses of CO2 from the analyzer cups prior to analysis.

Weisberg :

When Skeggs first described the AutoAnalyzer in 1957, that was a fine thing. Then in

1960, he designed the technique for the CO2 combining power, and, you may recall, he put the

entire tray into a box which was equilibrated with CO2 gas. Two years after that, about

1962, Masters had a different device. Some of you may remember the AutoAnalyzer when

they used to do hemoglobins and they added a little stirrer just before it was aspirated, to bring up the hemoglobin. He had a crook for the CO2 gas going in three times for each cup

and so he didn't have to put the entire tray into a box, but they did a CO2 combining power,

and therefore, representative of bicarbonate. Dr. Cohen and I were just talking here quietly

and he mentioned Dr. Gambino's technique with the al kal inization. Dr. Gambino and I had many

discussions on this. There is no way as far as I'm concerned, by definition, that the

AutoAnalyzer can give you a CO2 content, because it's not anaerobic and you don't have the

entire blood as a true plasma specimen. Whether you do the al kal inization or not, it's not

going to be, by definition, and that's it. Dr. Whitehead in England showed that very definitely.

Austin :

I think what we're saying is that there is a difference between the CO2 content of whole

blood and plasma. Van Slyke showed this many years ago.

Ladenson :

I think the point is that when most of us measure the CO2 content by the AutoAnalyzer,

we are off by perhaps as much as 3 or 4 milli-equivalents per liter due to CO2 losses to the

atmosphere in these little sample cups.

Weisberg :

You're really getting the bicarbonate.

Gambino :

That's something that Harry Weisberg and I argued about for a long time. I think you're

absolutely wrong. In the al kal inization technique, we drive the PCO2 to about that of the

atmosphere and when the CO2 of the sample is at atmosphere, in a 15 minute period, you don't

get any significant exchange. You don't measure bicarbonate. When you lose CO2, you have

a loss of carbonic acid and bicarbonate and you come down to a PCO2 of about 4 or 5 mmHg.

We've done a lot of work in this area and we do 400 or 500 6/60 's a day and several hundred

blood gases. Many of the patients have the same measurements and physicians are looking at the calculated bicarbonate and the measured bicarbonate, and if they disagree by more than

2 milli-equivalents, we hear about it right away.

Weisberg :

You just said bicarbonate, not total CO2.

Gambino :

No. We give them a calculated total CO2, which you can get from the pH and PCO2.

Austin :

I agree with Dr. Gambino. There is very little difference between the calculated CO2

content of plasma and the AutoAnalyzer CO2 content.

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Gambino :

When the sample is handled properly.

Cohen :

May I add my signature to that document. I agree. We have done thousands by the

Gambino technique and they are extremely close to the Van Slyke...

Wei s berg :

You put them into the equation that Dr. Rispens and Dr. Maas were talking about for

research purposes, or were you talking clinical?

Gambino :

This in cl inical

We is berg :

OK. Fine. There's the difference. By definition it is not a CO2 because it is not completely anaerobic. If we do say that it's a true CO2 content, we have to change the

original definition.

Gambino :

That's a different story.

We is berg :

That's what I'm talking about.

Gambino :

I'm not recommending it as a primary standard method. If you don't alkalinize, you get

disasterous results.

Weisberg :

What do you do with a patient who has a high PCO2 of 100 and you alkalinize? You're

changing it down to room conditions, and you're having to change what was existing in that

patient at that time at his temperature. Under most normal circumstances I would agree with

you. But when you have extreme changes in PCO2, high or low in a patient, you're bringing

him up or you're bringing him down to those room conditions and, therefore, you're not having

the same conditions for your patient.

Rispens :

I hope there is no misunderstanding on our method of measuring total CO2. There is no escape of CO2 possible before the sample is added to the acid reagent.

Engel :

I should like to return to Dr. Rispens paper in which he mentioned five reasons not to use base excess or standard bicarbonate. They are not required he says. But what is re- quired? You can use the same five reasons against actual bicarbonate concentration. If you

have enough information to define the system you can of course calculate or derive the remain

ing variables.

Ri spens :

After measuring pH and bicarbonate you can, of course, calculate base excess with any

nomogram, but what I stress is that you do not need it for assessing disturbances in patients

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Engel

But we have a very convenient method to determine base excess. Why shouldn't we use it

then.

Ri spens :

Because you do not need it.

Engel :

But you don't need bicarbonate. I don't need bicarbonate if I measure base excess, of

course. Why should I need bicarbonate?

Ri spens :

Actually you need it for calculating base excess. Calculating base excess is in fact

always based on quantities related to bicarbonate through the Henderson-Hassel bal ch equation.

Deriving a new quantity from the same measurement results can not produce any information

which is not already contained in measured values.

Engel :

If you go to the second reason I would like to say, and so what? This is again the

business about the interpretation of results and not measurements of data. An interpretation

always has to be done with regard to the patient and the ongoing disease and the history of

the patient, etc.

Ri spens :

Now that's what I have been saying continuously. My objectives against base excess and

standard bicarbonate is that they are presented as a quantitative measure for the metabolic

component of an acid-base disturbance. Base excess does not always mean excess of base. In the education of medical students we teach them in the second year how to assess acid-base

disturbances from pH, PCO2 and bicarbonate. We feel that they then clearly understand what

is going on. In the fifth year they go to the clinic and, coming back in the afternoons to the laboratory, they appear quite confused by all those derived quantities such as standard

bicarbonate and base excess.

Engel :

I can tell you that in our hospital it's exactly the opposite, because we taught them

from the very beginning about base excess and they think they do understand base excess.

Weisberg :

Can I alleviate this heated discussion between Ray and myself in the sense of what we use, and I'll have a different track when I talk a little bit later. Mark Twain said, "To do good

is noble; to tell others to do good is also noble, but much easier."

Ladenson :

I was very pleased to hear the separation of diagram and nomogram made because I think

that the two have completely separate uses. If I understand correctly, you use the diagram

as a clinical aid and therefore it has to be evaluated on that basis. But I would submit

that the nonogram is very close to a cop-out; because all you're doing with a nomogram is

trying to calculate a parameter that you either cannot measure or choose not to measure. In many cases, I think this can be dangerous. In tommorrow's session perhaps we can comment on McLain-Hastings calcium nomogram, but even today Dr. Mai enfant commented on a nomogram for

oxygen saturation which cannot be used at high levels of carboxyhemoglobin or methemoglobin.

In our institution, these are the situations when we often need an oxygen saturation estimate, when we are analyzing the blood of firemen or burn patients. Also we have seen patients

following an industrial chemical exp.losion with severe methemoglobinemia in whom oxygen

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saturation measurements would have been useful. This limitation may not be common but I think that what is really needed is a better method which would be accurate under any and all circumstances rather than a nomogram. I think it really is misleading and very dangerous

to depend on a nomogram.

Meisberg :

Not quite. Jack. Nomograms are fine because they're based upon data derived in the

laboratory. They are not theoretical; they may be based upon a formula based upon facts.

But what you're talking about is that they are not taking into account the methemoglobin and

carboxyhemoglobin. They don't have nomograms that way. And that's where the mistake is,

that you don't realize that you don't have it all as oxyhemoglobin.

Ladenson :

But the problem in that, Harry, is that new compounds have to be continually added to correct the nomogram. I think if you look at some of the current ones you'll see that you

have to measure six parameters to get at one. Perhaps we should have been spending the time

getting a better method to measure that one.

Weisberg :

Agreed.

Runck :

This question is directed to Dr. Weisberg, Dr. Visser, and Dr. Siggaard-Andersen. Is it possible for the three of you to agree on an equation for base excess? I've heard three

different versions of the base excess equation being given.

Weisberg :

I'll agree with Dr. Visser, that you need a big computer to take all these things in.

It is impossible to do it otherwise.

Vi sser :

You see that when you change the pH of blood, there is a water shift from the plasma to the cells. You could call that buffering. Because of the displacement of water from the

plasma to the cells when you increase the PCO2, you decrease the water content, the bicarbonate

concentration is increased and so the resulting difference in plasma pH is smaller than you

expected. So now you have to define the base excess including the water shift. That makes

it extremely complicated. You introduce the Donnan equilibrium, you have to introduce

osmotic pressure, you have to explain why the erythrocytes don't burst, and then we need a big computer to do a small job. You may have a practical definition of base excess which is

accurate within one unit. That's another thing you can aim at, but if you want to have an exact definition of base excess, an exact calculation, you're trying to shoot a fly with a cannon.

Weisberg :

The other way you're trying to shoot an elephant with a popgun.

Siggaard-Andersen :

Which is the best equation or algorithm for base excess? All equations or nomograms are based on empirical data for titration curves of blood and plasma, or buffer values for hemoglobin and plasma proteins. All approaches involve certain approximations and it is not possible to say generally that one algorithm is better than the other. The reference value will

always be the value measured directly by titration.

Cohen :

I would just like to ask Dr. Engel and Dr. Siggaard-Andersen whether or not they would

agree with the assertion that I made, that if one is beginning from the vantage point of

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wanting to develop a system that adequately described the acid-base status of a living

organism, whether or not they believe that there is any less information required in developing such an analysis utilizing the concept of base excess or standard bicarbonate than

is necessary utilizing the bicarbonate and PCO2 itself?

Siggaard-Andeirsen :

For the evaluation of the patient we actually do not need any of the derived parameters

(bicarbonate, standard bicarbonate, base excess, buffer base). All we need is the pH and the

PCO2 and a chart (like the one Dr. Weisberg showed). By plotting pH and PCO2 in the chart we can evaluate the degree of metabolic acidosis or alkalosis as the distance from the

reference line or the reference area for acute respiratory disturbances.

Those who are critical of calculating the excess concentration of base in the extracellular fluid as a measure of a metabolic acid-base disorder, but prefer to estimate the deviation from the reference line in the chart more loosely, are actually deceiving themselves

because estimating the distance from the reference line actually involves estimating the excess concentration of base in the extracellular fluid.

Cohen :

I'm sure you understand that I didn't' mean to imply that I thought there was anything

less informative about the base excess or standard bicarbonate. What I'm trying to understand is whether or not your position is that there is more information contained in these

parameters, i.e., do you require less physiologic understanding of the behavior of the organism in order to unravel a complex situation when using these parameters as opposed to the

pH and PCO2 and/or bicarbonate.

Siggaard-Andersen :

It requires the same amount of physiologic understanding to interpret the base excess

value as the interpretation of the bicarbonate. However, there is more relevant information

contained in the base excess value as exemplified by the direct proportionality between the

change in the extracellular base excess and the rise in the lactate concentration during

muscular exercise.

Engel :

Well, I agree with Siggaard. I think that in most cases, maybe 80 percent or more of

the cases, it doesn't matter whether you are using one or the other measure, but for certain

specific cases, for example the exercise case, I also think that you get more information out

of using the base excess than you do using the carbonic acid-bicarbonate system. But we want to go one step further, and think that if you in the exercise experiment had divided the

base excess, and preferably the base excess measured at a PCO2 of zero, up into the net acid part and the metabol izable acid part, then you would have had one further piece of

information; namely, that what you were dealing with was an organic acidosis as opposed to a net acid acidosis.

Cohen :

The difficulty, of course, is that when we are in real life, using various systems for

analyzing acid-base disturbances, we are presented with an unknown. We don't know that

exercise and/or anoxia and/or salicylate intoxication or whatever is at issue or what

combination of those factors is at issue. So it seems to me that the burden of the system

is whether or not it does, in fact, facilitate our understanding of these clinical disturbances. And in that context, I still remain to be convinced that I know anything more about

that individual when I know his "standard bicarbonate" or "base excess" than I already have

in my head when I know his bicarbonate and PCO2. I don't see that my understanding has been

improved by those additional parameters. This is an old argument. I know that we've discussed

it many times. But it seems to me that if, in fact, there is nothing more that I've learned

from these additional parameters, then I have to ask myself: Have I given up anything; is there a price or a risk in using these parameters? I'm not saying that it's not as good.

Certainly, one can develop an internally consistent system that describes nature perfectly

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well using base excess and standard bicarbonate. There is no quarrel with that. But, the

question is, if I'm not gaining anything, am I losing something? And if there is anything

lost, then it seems to me that the risk-benefit ratio would argue strongly against its

adoption. And I think there is a potent loss or risk. To the extent that the term "base

excess" or "standard bicarbonate" implies to the unsophisticated some information vis-a-vis

diagnosis or therapy, there is a potential for misuse. This is obviously not, as I said in

my presentation, an indictment of the system but is something that has to be considered as a liability in introducing such terms. I would submit that this is not an issue when one uses bicarbonate and PCO2 directly for diagnostic purposes.

Weisberg :

Just one comment in reference to Dr. Cohen. Unfortunately, many people, especially in

the United States, think that the M.D. after our name stands for Million Dollars. That is not so, I'll guarantee that on a Bible. But also, on the other hand, I don't think the M.D. should stand for Mentally Deficient. And, unfortunately, we have a problem that we can give

all the laboratory data, those of us working in the lab, and if you're on the other side, the

practicing clinician, we can give all the laboratory data to the clinician and, if he's not going to use it properly, he's going to have a big problem. Let me give you an example. If we have a patient who has 25 mill i-equivalents per liter of keto-acids, calculated, what

would your guess be for the diagnosis on an acid-base evaluation on that patient?

Cohen :

I don't have enough information.

Mei sberg :

That's exactly the point. Because you find more keto-acids with metabolic alkalosis due

to perniciou^; vomiting then you do with keto-acidosis due to diabetic acidosis. You gave me

the answer and we agreed; you need more information which you get from the bedside. You can utilize the laboratory data very well whether you use the base excess concept or the PH-PCO2

concept. The average clinician is not like that, unfortunately, and therefore that's the

difficulty that we, who are servicing that type individual with the laboratory data, have.

Co hen :

I'm not so sure from what you've said, what side of the question you've come down on.

Weisberg :

I'm on the side of the question saying that they agree; all you need is a pH and a PCO2 presented as numbers to distinguish seven diagnoses of acid-base imbalance. You need

a bicarbonate (or total CO2 content) for two additional ones--to distinguish between a mixed and a respiratory imbalance. The only time I use a base excess (or delta content) is to distinguish between the acute and chronic respiratory conditions. Even with that, I still

would have to see the patient to make a definitive diagnosis. I'm not hedging and straddling,

but that's a statement of the fact of the case that you cannot make a diagnosis without

knowing the condition of the patient.

Engel :

Well, I think I agree very well with Dr. Weisberg, and maybe one of the major problems

in understanding acid-base is that most people don't know enough physiology.

Layer :

Well, you remind me of a comment made by Joseph Barcroft 50 years ago, namely that the

art of successful advertising consists of saying something which is true but totally irrelevant. At the bedside, nothing categorizes the critically ill patient better than the un- steady state. To make much out of derived values in the unsteady state is a potentially

dangerous proposition. Despite my experience with critically ill patients, acid-base balance,

and gas exchange, I would not know how to use these numbers, and I would be loathe to react

upon them. Although we do not report derived data in our institution, they are available

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upon request. I like to think that it has not made much difference in the care of the critically ill patient. What is the value of such interpretation when the derived data, based

on hemoglobin concentration, are received at the bedside after the patient has had an infusion

of two units of packed red cells and the hemoglobin has increased from 10 to 15 g/100 ml? In exercise, a trained athlete may have a blood lactate concentration of 40 mM at a particular

pH and PCO2, and yet a patient with exactly the same values for lactate, pH, and PCO2 may

be moribund. How do we interpret that? Derivation is important when applied to the steady

state. Its relevance to the critically ill, characterized by the unsteady state, is highly

questionable.

Gambino :

I have just a few comments. One, regarding plastic syringes versus glass and the data

that you showed. We've done some studies in our own lab and find that if the blood gas comes to the laboratory 50 minutes after it's drawn, in my opinion, it's practically useless,

and what we do are two things. You'll hear tomorrow that we have peripheral laboratories so that we decrease tremendously this transport problem that blood gases should be measured as close as possible to if not in the site where they are to be used, to eliminate that problem.

Secondly, we use ice not so much to prevent change but the abundant ice melts and spills and

the person who is given the specimen to carry, sees the melting ice and knows he has to hurry. The third thing we do is that the person who brings the specimen returns the report,

because, as mentioned, the actual measurement can be done very rapidly. The specimen comes

to a central area and the person who brings the specimen returns the report. That way, the

person who sent the specimen to us knows what's happening because he expects that report to come back with the messenger and is less likely to be lost because he knows he must come

back.

Regarding the syringes, the dominant reason for measuring PO2 frequently, is to monitor

oxygen therapy so you don't get above 100 mm and below 60 mm, approximately. And most of

the specimens that are critical are in that 60 to 100 range and are coming within 15 minutes.

In my experience, there is no practical problem with plastic syringes versus glass, if the

blood gases are done rapidly. There is a problem with dilution. I am concerned about the

amount of heparin in the syringe and too much heparin will introduce PO2 errors in a sample.

We've had problems with glass syringes when the barrels are loose. There is a need for an improved collection system that would be impervious to gas and would not have any gas stored

in it and would have a sealed barrel that moves very, very easily. Al kal inization that we perform for the CO2 content, done routinely, is performed when the Vacutainer arrives, with

the stopper on. It is centrifuged sealed, so there is no change in total CO2 content. As soon as the stopper is removed, that's the al kal inization point right at that moment before

any transfer takes place, because as soon as you transfer you lose CO2, much as you do in pouring a Coke or champagne. You're going to lose CO2 when you move the specimen. Finally,

another unpublished study, done at the Presbyterian Hospital in San Francisco by Dr. Burns

and the cardiovascular group there, has shown that the a venous-arterial pH and PCO2 is very

small as long as the cardiac index is above three liters. As soon as you start getting

decreased cardiac output, you increase that spread and, of course, this is the basic problem

with venous samples and capillary samples. In very, very sick patients in shock, the patient

with a serious problem that requires frequent blood gas measurements, those are the patients

that may give you the largest errors. So I think that for the critically ill patient, we must get that arterial specimen.

Ladenson :

The use of mercury in the syringe is very dangerous as there have been two documented

cases of embolism due to its use.

Unidentified :

I think that this was intended to be after the specimen was taken.

Gambino :

Even then it's not needed. We used it a long time ago; mercury in the syringe, but it

is dangerous to the laboratory personnel as well. But I think it's been shown that you get

adequate mixing by shaking and I think that is the most important. It is essential that the

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sample presented to the instrument be mixed whole blood and not have any separation. We do see samples coming into the lab that have been iced and you get extraordinary errors in PO2

if you don't mix that specimen uniformly. So I think that's a critical aspect of presenting

the sample to a blood gas machine.

Ladenson :

What is your median delivery time using your peripheral laboratories?

Gambino :

The peripheral labs are in the intensive care units and there are the four intensive

care units in the medical center. There are four peripheral labs, and they service those

patients in the intensive care unit. So it is as fast as you get the blood and walk to the

machine. I would say the mean time is three, four, or five minutes. Now the mean time to the central laboratory will vary from ten to twenty minutes, because of the elevator and

walking time, but never longer than 30 minutes.

Weisberg :

Do you have a special runner going up to get the blood?

Gambino :

No, the bloods are drawn at the site, not by the laboratory.

Weisberg :

You mean the house staff. .

Gambino :

Or whoever; the inhalation therapist, the anesthesiologist, the nurse, the nurse practitioner, a variety of people draw the specimen. Because radial artery puncture with a 23

gauge needle and no anesthetic is in many ways easier than drawing venous blood. They do

multiple punctures every day. And in the 6 years I've been at Columbia, there has been

absolutely no serious morbidity, none, from the radial artery. Not true of the femoral or

brachial

S^rensen :

I should like to comment on those plastic containers because we have started looking

at these in the last few months. I feel the problem is not so much the plastic as it is the

plunger that causes the problem. Because in the plastic syringe, the plunger friction is

much larger than in a glass syringe. If you can avoid this by using radial arterial puncture

with small needles and having air bleed in the plunger you get a much better arterial blood

sampl e.

Austin :

I've been an advocate of venous blood gases, but I have seen recently a number of patients

where the A-V PCO2 difference has been quite wide. We followed the instructions religiously,

the instrumentation I think has been OK, and I just cannot explain it when a patient has not

been in shock, but it just seems to be cropping up in about 5 or 10 percent of the patients

where the spread is greater than the 10 mm Hg. I just have no answer for it. It disturbs me and I've searched continuously and I cannot find it.

Burnett :

How do you do tha- determination? Do you draw venous and arterial specimens simultaneously using the same collection techiques?

Austin :

Yes, same collection technique using glass syringes.

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Gambi no

These are people at rest?

Austin :

People at rest. Patients in hospitals. Warm patients. I'm on one side, someone's on the other side and we're doing arterial and venous puncture. It doesn't usually change the

diagnosis or the clinical evaluation, but the spread of PCO2 does seem to be wider, and I just have no explanation for it. I wonder if anybody has seen that sort of thing.

Gambi no :

I have no recent data since we're doing almost exclusive arterial blood work now.

Runck :

Two comments. First of all. Jack Ladenson commented about the O2 decay. If we're

referring to PO2 decay in a blood gas analyzer, this, in fact, is probably due to a leak

in the system. If the blood has been presented as a properly iced sample, whenever PO2 decay

is observed, it can be corrected by really looking at the fit of the electrodes around the

system. Hanson and Neville in Syracuse actually determined this when they put their blood

gas analyzer in a great big plastic bag and filled the bag with 100 percent oxygen and, lo

and behold, the PO2 decay suddenly disappeared. The second comment concerns use of plastic

versus glass syringes. I would agree with Dr. Gambino, you should not discount the use of

plastic syringes. There still are people around using glass syringes who allow the bottom

of the syringe to fall out when the syringe is full. I think we have to recognize that not everybody is as proficient as we'd like in terms of handling samples; use of plastic syringes

make sample handling easier and more reliable.

Howorth :

I have a comment about Dr. Ladenson's paper. This is in connection with the plastic cover used on the AutoAnalyzer. I understood that the reason one has to have an orange

plastic cover is because the bilirubin is extremely labile to fluorescent lighting.

Ladenson :

There are two types of plastic covers, one of which is amber and the other clear,

Howorth :

Well, OK. It's essential to stop bilirubin from disappearing due to photolysis prior

to analysis. Then the second comment is about the rate of fall in vitro in blood PO2 in leukemia. I understood that primitive white blood cells utilize anaerobic metabolism to a greater extent than normal white cells.

Ladenson :

I have no evidence of my own. The one study that I have found concerning PO2 decay

and white blood cell count did not distinguish as to the cause of the elevated white cell count.

Gambino :

I think it's important that the plate cover did no good in the AutoAnalyzer. One very

important reason, when you use the little disk that was described for the cup, what's happening there, of course, is that the disk is effectively reducing the surface area of the

sample exposed. The flow rate or volume of CO2 escaping per unit of time from a particular

surface area is decreased. But when you have the cup, like this, you have a big plastic

thing, there are all kinds of cracks in it. It doesn't mean a thing, and so the surface

area exposed to the atmosphere is the same as it is right out in the open room. This does

absolutely no good at all. That's why we never reported on it. It's absolutely useless.

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Now as far as the bilirubin is concerned, it is equally useless. If you come to my lab you'll

see that we have a stainless steel cover that's opaque to light because the bilirubin will

deteriorate through that semi-colored cover. So if you want to get the light effect out, we

use a metal cover over the AutoAnalyzer tray, and it is an important factor if you have bright

1 ights or sunshine.

Weisberg :

That kind of cover must be used on the AutoAnalyzer. However, if you allow the cups to rotate for an hour, you will find that if you have it uncovered, the dehydration of the

proteins will be increased by about 8 percent. So that's an absolute technique for getting

to the one with the acid-base blood gases.

Gambino :

That's not as effective as it should be either, because you still have the leaks.

Weisberg :

Yes, but it's going to reduce it. Without it you're going to have an 8 percent change.

Siggaard-Andersen :

The papers by Dr. Gambino and Dr. Laver were very interesting, but it occurred to me that neither of them stated directly that another worthwhile effort might be to try to reduce

the number of requests. I think Dr. Laver hinted yesterday that we produce all these numbers

but what do the doctors do with them? There were a number of redundancies among the tests

you mentioned, e.g., serum osmolality and serum sodium.

Gambino :

I think your point is well made, I'm very aware of it, and I believe it can be approached

in several ways. First, the peripheral ization of the laboratory as I conceive of it and as

Dr. Laver was suggesting, when the physician himself or the people taking care of the patients

actually do the measurements, there is direct control. You see, the laboratory is simply

an extension of physical diagnosis. It's a more refined physical diagnosis. Now, the physician, himself, when he performs a physical exam, knows he has only 24 hours in a day and

then he has so much time to give to the patient, so he rations his time accordingly. When

you have a large central laboratory, and all you have to do is write an order on a chart,

that is conceived of as an infinite resource, much as air and light and water. There's

plenty of air and plenty of water, but we know it's a finite resource, and a laboratory is

a finite resource. So, one of the ways of avoiding abuse of this finite resource is to put

the resource directly into the hands of the person who is using it so that he, himself,

experiences that. And that's what happens with the peripheral unit. In other words, if they

start doing more blood gases, they'll just buy the blood gas machine, set it by the patient,

so that providing laboratory services at the bedside is one way and having the person who is

utilizing the lab tests become more active as part of that, either by actually doing it or by

actually bringing it to the laboratory. In the case of stats in our hospital, a telephone

call is required and a stat voucher is made. We do not pay any attention to written stats. There must be a phone call by the physician, it must be an action so that there is involvement. I mentioned yesterday about the specimen being brought to the laboratory. When it comes to the central lab, the person who brings the specimen must wait to bring the report

back. That closes this chain. And so we are very, very concerned about it. The second

thing we do to eliminate tests, and we just last month no longer do single SGOT's as isolated tests. We don't do HBDH. When we introduce a new test, we always look to see if we can eliminate another test. So just as the physician only has an hour or a half an hour--or whatever the time--to do his physical exam, I only have a certain amount of time

and resource and total energy in the laboratory to expend on these results, so I must monitor

what's happening and how useful it is. But I agree that it is easy for this to take off

and be abused. But when the person is doing it himself, it is much less likely. So that the

second thing is that I don't think there is going to be any slackening in the lab work because

the lab is such an essential component of physical diagnosis. You don't want to wait until

the patient is blue to diagnose anoxia. And we don't want to wait until the patient has

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cardiac arrest to diagnose digoxin toxicity. Therefore, the physician requires these techniques and, as I told Dr. Laver last night, if you deny a physician laboratory tests, you're

effectively denying him access to his patient. It's as if you locked the door on the room

and you denied access to the patient. I think this is extremely important for the clinical

chemist to appreciate and the people at NBS and the regulatory agencies to realize, that

laboratory testing is such a direct continuance from physical diagnosis that you are interacting very directly with physicians' activities. And the physician is not as dumb as sometimes has been alluded to--that they don't know what they're doing. They, indeed,

have a pretty good idea of what's happening, and I think that the key thing is for laboratories to provide those tests that are physiologically and biochemically correct. When we understand the physiology and we understand the biochemistry of the disease process, then

the testing becomes important. You will find that as our knowledge of any disease process

increases, testing volume will increase because the test becomes critical. Now that we understand pulmonary physiology better, the tests are extremely critical. When we don't

understand a disease, such as diabetes, testing is very vague and diffuse and not very

useful

.

Austin :

I think maybe this has been referred to, but it's my contention that pieces of equipment

need a mother, so to speak, and we have found that when errors develop and when deviations occur it's because too many people have access to a piece of equipment. They really need to be, in my opinion, supervised by a very few people and I think when you get too many people

using it, too many handling it, that's where it breaks down. That's my theory.

Laver :

In regard to Dr. Siggaard-Andersen ' s comments--yes , you are quite right about osmolarity. We introduced it a long time ago when we did not have immediate access to a blood

sugar measurement and, of course, the blood urea nitrogen. Now, we receive few requests

for serum osmolarity. However, matters are quite different when one is dealing with blood

gases. These patients are usually in the early phase of acute respiratory failure. It was our original intent to provide continuous measurements by repeated analyses and a short

turnaround time. Depending on severity of disease and the number of patients in need of

ventilator support, we perform an average of 400 blood gas determinations daily in some 80

patients. Fortunately, the number of patients in need of our services has not increased, and

the average number of samples per patient has not changed. The overall demand is not likely

to be altered until we have reliable, continuous blood gas measurement devices.

Richards :

I direct this to Dr. Gambino. One of the figures I jotted down was a relative standard

deviation of 11.3 prcent for PO2. If you took that at the 2 standard deviation level, it

would mean about 22 percent.

Gambino

Yes. It's not tolerable.

Richa rds :

It seems to me, if this kind of quality figure is being generated, it might be better

for a laboratory to cut back on the work in order to provide a better quality statement.

For those of us who participate in proficiency testing, 22 percent is really very high.

This is the sort of thing you see for some of the enzymes.

Gambino :

I didn't want to give the impression that I found that acceptable. I'm telling you that

this is what was seen when a particular study was performed in different laboratories. In other words, across laboratories, and of course you have to realize that indicates that there are problems with instruments in the field.

334
Kreuzer:

I would like to return to Dr. Laver's remark. It is also my contention that the future

is continuous recording, and I would like to draw the attention again to the group in Brussels which is led by Dr. De Meester. This group has embarked on a rather large program to apply these continuous recordings. Of course, I am well aware of all the problems, but I think it's worth the effort. My second remark also refers to Dr. Laver's paper. I wonder

whether anybody has had experience with the microporous membranes as applied to measuring

problems e.g., in polarography, because microporous membranes might be just the material we are looking for. This is due to the fact that we are concerned here with a suspended gas

phase in essense; a large part of the membrane is essentially gas phase. You have high

diffusivity and practically no solubility. Has anybody had experience with these membranes?

Brand :

I tried using such a membrane. You refer I take it, to something like the microporous

Teflon membrane that Millipore makes, that kind of material. In our experience, it clogs

with red cells very quickly. It works very well for about half an hour and then clogs up irreversibly.

Engel :

I would like to return to the problem of the decentralized laboratories and I agree with

Siggaard that when you get up to a size as we have in our laboratory, with more than 200 em- ployees, it is comparable to a power station where the requests come in just like when you

turn on a switch. But I also agree with Dr. Gambino that establishment of peripheral laboratories probably will reduce the number of tests which are not really necessary in order to handle the patient. We have a very good example, when we made a peripheral laboratory in our

pediatric department, the number of tests from that department went down by more than 50 percent.

Vi sser :

I think that the need for blood gas analysis can be enormously decreased by using con- tinuous CO2 monitoring in the expired air. In Europe, this has been in use for about 20

years. Twenty years ago I had to build my own infrared analyzer because I had no money to buy one. It is now everywhere in use, hundreds of analyzers, and especially for respirator

treatment of patients with normal lungs, you don't need the blood gas analysis because, in

that case, it is especially the CO2 and the pH that is of importance. And the O2 doesn't

change so rapidly. So this is also decentralized, and you can do it as continuous monitoring

or intermittently, if you don't want to buy so many instruments. Of course, you can't do

that for oxygen. There is a big difference between Europe and the United States, in that

respect.

Laver :

I think that comment requires a vigorous rebuttal, particularly from someone whose

experience has indicated that end tidal PCO2 measurements are worthless in the critically

ill. First, because the difference between end tidal and arterial PCO2 can be substantial.

Second, the blood-gas exchange problem in patients with acute respiratory failure (ARF) is

oxygenation, not CO2 removal. The latter is apparent as a terminal event in ARF and common with patients in chronic respiratory failure. To my knowledge, there appears to be little

difference of opinion on the subject. Monitoring of end-tidal CO2 may be of value to follow

performance of the ventilator but not for evaluation of blood-gas exchange in the patient

with acute respiratory failure and previously normal lungs. I find no justification to con- done this return to a questionable practice.

Visser:

That is not what I am proposing. I am proposing a decrease in the number of blood gas

analyses by replacing a number of blood gas analyses by measurement of alveolar or end tidal

CO2, taking into account the difference between end tidal CO2 and arterial CO2. If you have

a patient's arterial and alveolar PCO2, you can follow changes of PCO2 by measuring expired

gas instead of blood. So the number of blood gas analyses is decreased but not replaced by .,

335
Laver:

Unfortunately, things do not work out that way in practice. The number of blood gases

is not reduced because one must still follow the level of arterial oxygenation. Occasionally,

samples must be drawn every 10 to 20 minutes because repeated evaluation of arterial PO2 is necessary. However, such determinations include the PCO2 and pH. Addition of acid-base

evaluation does not complicate or lengthen the procedure by more than a few seconds.

Kreuzer :

I'm inclined to agree with Dr. Laver, and I might add that the electrodes which we have

described are equally applicable to the gas phase, to the continuous recording of PO2 in the

respiratory gas. We have done hundreds, even thousands of experiments over many hours, with

recordings in all kinds of situations. We have even developed a telemetric method. I might

add here, since I couldn't go into this for lack of time, that we succeeded to telemeter, over distances, of one to two kilometers, the oxygen uptake of a subject by means of, among

other parameters, the continuous recording of PO2 in the respiratory air.

Si ggaard-Andersen :

Has anybody had experience with the use of the isotonic bridge described by Maas some years ago? The isotonic bridge eliminates the effect of varying hemoglobin concentration on the liquid-junction potential but it does not eliminate the effect of varying ionic strength

in the sample. The latter effect can possibly be calculated by means of the Henderson equation. In dehydrated or over-hydrated patients the ionic strength of the plasma may vary

considerably which might cause a considerable variation in liquid-junction potential. This

would be of special importance with the calcium electrode where the electrode sensitivity is

only half of that for a monovalent ion.

Austin :

I'm not sure I'm answering your question, but I think saturated liquid-junctions give more stable readings in my experience. I'm not sure that's what you're asking, but that's

my experience.

Bates :

I'd just like to point out, in connection with that figure of Dr. Maas, that strictly

speaking, these standards, which have an ionic strength of 0.16, are not on the NBS scale

exactly because the NBS convention was limited to 0.1 ionic strength. However, I think the

excellent agreement that you find experimently might be justification for relaxing this re- striction.

Weisberg :

May I change this subject before Dr. Laver leaves and I'd like to have this on the

record. Dr. Kreuzer could answer this too, because back in 1971, Severinghaus and his group

reported an error on the oxygen electrode due to the reduction of the halogenated hydrocarbons,

especially halothane. This has been in the literature and I've asked several anesthesiologists and it seems to be a puzzle. I've discussed this with Dr. Laver but I'd like him to put it in his own words for the record and perhaps Dr. Kreuzer could answer too. Does halothane have an effect on the polarographic electrode for oxygen?

Laver :

The answer is no. We have run polarograms with whole blood in equilibrium with up to 4 percent halothane, a variety of electrodes, a variety of membranes and buffers (pH 7 and pH

10). We have found no discrepancy attributable to halothane. This holds true for methoxyflurane, fluoxene and other halogenated anesthetics, irrespective of oxygen concentration. The

effect of halogenated anesthetics on the readings of the PO2 electrode need not be of concern

to clinical laboratories.

336
Kreuzer:

First of an, I would like to confirm this, but maybe one word about CO2 should be added.

You keep reading statements in the literature about CO2 sensitivity of the electrode. We have investigated this very thoroughly, too, and in our experience, if an electrode is markedly

CO2 sensitive, then it is apt to be a bad electrode. If the electrode is constructed properly,

{i.e.:, if there is sufficiently fast exchange of ions between the electrolyte covering the

cathode and the bulk of the solution) and if the bias voltage is chosen correctly it is practically not sensitive to CO2.

Austin :

Just getting back for a moment to the liquid-junction subject, I would be interested

in Dr. Bates' feelings about the molarity of the solution for a liquid-junction.

Bates ;

Well, I really don't have any special information beyond that presented by Dr. Maas. I think everyone realizes that the early work indicated that, in systems which are homogeneous,

the stronger the concentration of KCl the more reduction in the liquid-junction potential

that ensues. But I think that with a system such as blood where you have the danger of cre- nation of blood cells and precipitation of proteins, etc., you may have a special problem.

Perhaps that outweighs the matter of the actual magnitude of the liquid-junction potential

that you would expect in completely homogeneous systems.

Runck :

I have a comment dealing with Dr. Gambino's statement on the performance of the Versatol

Controls and Dr. Noonan's statement on the Coke machine concept. I think Dr. Gambino's data

on the Versatol Controls presented two aspects of performance. The first aspect is the

precision, or repeatability, of replicate measurements. The second aspect wasn't mentioned

and that is the accuracy, i.e., the bias, seen by the laboratories reporting the values.

Our experience has been that the precision problem is actually an operator problem. And that

can be corrected very quickly by going in and retraining people, showing them how to do it

properly. The bias problem is more of a systematic problem, and that has to be looked at in

instrumentation. Now the Coke machine type instruments, the automated instruments, tend to deal with the precision problem very well. We find very good precision performance with

automated type instruments.

Ladenson :

My question is actually in the form of a request to Dr. Durst and the Bureau of Standards. I hope that future certified pH standards would be checked with systems that have

liquid-junctions. For example, when we analyzed the certified Tris buffer against the certified phosphate buffers we got a value that was 0.01 to 0.02 lower than the certified value

{Clin. Chem. y 20^, 1337, 1974). This means that the two certified buffers will give different

results on the type of equipment that is used in virtually all laboratories.

Durst:

Preliminary studies made at NBS several years ago indicated that this a real effect,

but the magnitude of the problem is still under study. Unfortunately, we do not have clinical pH instruments available for our testing and must rely on classical liquid-junction

measurements.

Bates:

Didn't you designate that as a secondary standard?

Durst :

It was supposed to have been issued as a provisional standard until such time as we demonstrated consistency with the pH scale for cells with liquid-junctions. But at the

337
point where we certified it, the Office of Standard Reference Materials decided that they

weren't going to issue provisional standards anymore, and it went out as a final certificate,

without any indication that it was to be considered a secondary standard. It was an internal

communications problem in that respect. It went out essentially as a primary standard, when

it really was provisional as far as we were concerned. The certificate has since been revised

to reflect these considerations and also carries the secondary-standard designation.

Kreuzer

I would like to make two remarks with particular reference to the PO2 measurement; one, the gas-liquid factor and two, the tonometer. As far as the first point goes, what I often

miss in these graphs relating gas PO2 to fluid or blood PO2 is the exact description of the

apparatus. In other words, the blood liquid-gas factor is nothing else than a stirring factor or flow factor, or a flow dependency as we call it in our continuous system, and it has

only meaning if we know exactly what the electrode is, that is to say, if we know the diameter of the cathode as well as the thickness and the material of the membrane. And this is

not always self-evident. It may be self-evident for the user, but not for the listener.

This factor may be quite different depending on the type of electrode. Radiometer Company,

e.g., has used different types of electrodes over the years with different diameters of the

cathode. Concerning the second point, the tonometer, it has been described already decades

ago that in tonometers there is apt to be a separation between plasma and erythrocytes. If

you draw a sample from the tonometer you may easily have blood that has a higher hematocrit

than native blood and this, of course, as Dr. Runck showed in one of his graphs, will affect

the gas-liquid factor. That is to say, you have to know when you are tonometering what

the situation is in terms of the hematocrit.

Gambino :

I wanted to ask Dr. Maas about his laboratory. You presented data on various instruments

in your laboratory in regard to your aqueous quality control calibration system. Which junction do you use in routine practice? Do you use your sodium chloride junction or, when you

use the instruments, do you use the junction material recommended by the manufacturer?

Maas :

We have used the junction recommended by the manufacturer.

Gambino :

So that you are not in actual practice, in your hospital, carrying out on patient samples,

the isotonic junction?

Maas :

No, we don't use that type.

Gambino :

Do you think you will? In other words, we woul<l like to know, to help us to decide which

way to go.

Maas :

I feel that in practice by changes of the concentration of the isotonic solutions, for

example, by diffusion of phosphate, poorer results are obtained than using saturated KCl

.

Gambino :

So, at the present time, you would not favor recommending moving in that direction.

That clarifies it then.

338
Sigqaard-Andersen :

Concerning reference methods, I think it might be possible to develop a reference method

for blood pH on the basis of the operational definition given by lUPAC, using a capillary

electrode, standardized saturated KCl liquid-junction, etc.

Concerning PCO2 and PO2, the reference method would probably have to be based on gasometric analysis of a small gas bubble in equilibrium with the blood, i.e., the Schola ^er technique. This techinque was too complicated as a routine method but with new techniques """t may

prove useful as a reference method.

Ladenson :

There have been some statements made that this reference material would have to be blood.

If it were blood, how could we standardize its use? I cannot see any way of doing this

because everyone will tonometer it differently. Are we going to have a standard tonometer?

From a practical standpoint, the problem is that all instruments do not give the same results.

To those of us in the field, the problem is which one is correct, and how can we ascertain it. To the manufacturer, I am sure the problem is similar when it comes to instrument design. The

aqueous material (General Diagnostics) described previously, apparently parallels the differences one sees with blood. This material which was apparently shipped all over the

country under all kinds of conditions has very similar results in various laboratories. It appears like we now have a good start on a reference material and we need a central unit

to provide values for it. I think we are now sitting in the place that can do this better than

anyone and can do it independent of any particular manufacturer's electrodes. I believe the

National Bureau of Standards should play a major role in ascertaining whether there is a material that is stable and whether values for pH and blood gases can be assigned to it based

on rational measurements independent of a particular commercial instrument.

Durst :

I would like to respond to that briefly insofar as we are being put on the spot. This

was one of the purposes, as far as I was concerned, of the workshop; to try to get ideas of

what could be used as a good standard material for calibrating pH and blood gases. We were

thinking along the lines of one of our buffers in isotonic saline and coming up independently

with a gas mixture of certified CO2 and O2 with the balance in nitrogen. Then one would have

to develop a reference tonometer to put them together and come up with a reference solution.

But it seems from the discussion of the materials that Dr. Gambino described that there

may be a material already available that we could use in this way.

Gambino :

I want to clarify that. I think it's essential that the National Bureau of Standards

provide a reference material for gases, in other words, a certified CO2-O2-N2 mixture as a reference sample. So I think that's step number one. That can be done. Your mass spec can help you there. You've got the technology to do that. That will fix a reference point. Now

then I think the material such as I showed, and the material that Dr. Maas has developed in

Holland, are secondary materials and will be dependent upon the primary gas standards. If

you don't go to blood, you can define the PO2 of those systems, and you can also measure,

you can recover, you can do gas chromotography , you can analyze that solution which is a simple solution, a single aqueous sterile solution, and quantitate the amount of gas. And

from the solubility coefficient, etc., you'll know what PO2 to expect. But I would see it as kind of a secondary reference material, not something that the NBS could successfully do

immediately, or should do. I think you should definitely do the gas as soon as possible.

Austin :

The same thing should hold true for buffers.

Gambino :

Well, they've done that. Along those lines, I think the NBS should consider doing more

work at 37 °C on other buffers. I think that would be essential for more physiologic research.

339
Me is berg :

What about bracketing the normal pH with something above 7.4?

Gambino :

Exactly, something higher, a buffer certified at 37 degrees that would go above 7.4.

Ri spens :

I think as far as PCO2 is concerned, it is much easier. You do not need gases. You can make a mixture of a bicarbonate solution and a phosphate buffer which has a definite PCO2

after mixing.

Gambino :

Do you supply gases for any other type of standard?

Durst :

Yes, for example, we have certified air-pollution gases which are mixtures of CO2 and

nitrogen.

Runck :

Blood gas measurements shouldn't be any different than any other clinical chemistry measurement. We need both a standard reference material, as suggested by Dr. Noonan, which is

traceable to NBS, but in addition, we must have a standard reference method for the determination of blood gases, just as we have atomic absorption for calcium and hexokinase for

gl ucose.

Ladenson :

The only problem I see in stopping there, is that some manufacturers have designed instruments that would not be compatible in routine use with a primary gas standard, e.g., the

Radiometer ABLl system. For this reason, I believe that we need something that will be

applicable to all instruments. I wonder, and I address this to the manufacturers, would

such a non-blood material be acceptable to you as a starting point to try to get these values

together.

Runck :

Absolutely.

Ladenson :

Then I think that is what has to be done.

Austin :

Everything is relevant to that.

Ladenson :

Well, we can all agree that with such an NBS certified material we can begin to standardize the results obtained with different instruments.

Gambino :

It would be possible with an NBS gas in a defined tonometer system to reproduce ...

340
Ladenson:

What I'm saying is that the material has to be made or purchased in large single lots

in the same way as all the other Standard Reference Materials. NBS writes the specifications,

analyzes the material, and provides a certified value for a material which you can ship

around the world.

Unidentified :

I don't know how other people feel about this, but I found using a wide variety of different types of tonometers, made in different ways, that it's not that difficult to tonoeter a solution accurately.

Runck :

That depends, in our experience, on the level to which you tonometer your blood samples.

Most of the data we have seen are at normal values for PCO2 and PO2. Those don't stress the

instrument. It would probably be advisable to have, in addition to normal ranges for PCO2

and PO2, a control material or a reference material at the extreme values as well.

Gambino :

I would disagree with you, at least clinically, maybe in physiologic research you might

want it, but clinically we don't want patients to go above 110 millimeters at all.

We is berg :

What if you have a patient on therapy?

Gambino :

That is why we do so many oxygen tension measurements and nobody is supposed to go over

100 millimeters, period. That's why we do it. You keep lowering the inspired oxygen and

that's why, in my particular laboratory, I couldn't care about accuracy at very high P02's.

Weisberg :

You get levels of 300 and 400 in cardiac bypass.

Gambino :

So what? It doesn't make any difference.

Weisberg :

The point being that they want those with accuracy, and they're trying to make a decision on it.

Gambino :

No they don't. They don't want it at all.

Runck :

I think there's controversy on that. There are two points on the table right now. Number one, can you tonometer accurately? Number two, if you can, what are the correct

levels to use? Each has to be treated separately.

Gambino :

There may be some special physiologic situations in a hyperbaric situation where some- body's trying to maintain a high PO2. In a hypothermic situation, they want a high amount

of oxygen physically dissolved. Those are special cases. I think when you see the problems

341
that are occurring, however, the critical decision levels are at 100 and at 60. Those are

the dominant decision levels. It's essential that we be accurate and precise at those i levels. I think you get into tremendous sample handling problems with a high PO2. At high

' PO2, you have to go right from your tonometer or your patient directly into the instrument. 1 If you have any type of sampling handling, you're not going to get the right value. !

Siggaard-Andersen :

| Just a comment on the technical problems of tonometry. If you have very small bubbles,

the pressure inside the bubbles is higher than atmospheric pressure. If you have a very

tiny outlet from the tonometer, you may also have a higher pressure. It is not always easy

to obtain exactly the same temperature in the tonometer as in the electrode, so if we want ii accuracies of 0.1 mmHg it is not going to be easy.

'

Runck :

I

My comment relates to Dr. Gambino's comment on the levels of PCO2 and PO2 at which the

operator should test the instrument. I think we must keep in mind that there are two purposes

in testing an instrument. First, is to assure yourself that patient treatment is being

carried out on good numbers. Second, is to give the operator early warning of impending

instrument problems. In the case of CO2 measurements, if we choose 35 millimeters as our ! checkpoint for CO2, the carbon dioxide electrode could actually turn to stone and you would

still get very nice readings at 35 millimeters because the electrode is always kept in the area of 35 millimeters. It is very important that a CO2 level other than 35 millimeters be

tested. And the same thing with PO2. The early warning of a PO2 problem witli an instrument

is going to be very apparent at a high PO2. If you have excessive atmospheric contamination

or a leak in the sample chamber, giving relatively small problems at 86 or 90 millimeters, you

will see this in exaggerated fashion at 200 or 300 millimeters. This is why manufacturers

should state performance characteristics at these levels as well.

Gambino :

I agree with that. You also see it if it is very low.

Cali :

I Would it be interesting to the group to find out how we would handle the question of

reference materials here within the Bureau, in order to arrive at a reference material?

Group : I

i Yes. I

Cali :

j I've heard you talking about accuracy and precision. Where it is possible, we believe | that all standardization should be based upon accuracy. I won't go into the complete logic \ of that proposition but if you are trying to achieve compatabil ity in measurement and by ; that I mean simply that everyone measuring the same sample gets the same result, then that 1 goal can be achieved directly if all concerned are making measurements that are accurate. So, | we start off by asking what is the basis for achieving accuracy? Well, accuracy is related j to the true value, but let's not get into a philosophical fight on what we mean by "true , value." But you can say in an operational sense, that somehow you've got to have access to

| the base and derived units of measurement, which is now the SI system. We have access to the SI by what we call here at NBS definitive methods. Now, definitive methods do not

exist, in which case as an alternate, you then use two or more reliable methods. What | constitutes a reliable method is a matter of expert subjective and, hopefully, objective

j opinion. The preferred way, however, is via the definitive method. A definitive method

[ must fulfill several criteria. You must have a well-established theory and usually it 'j should be expressable in mathematical terms, usually via appropriate equations. When you [ make such a definitive measurement, using the equation, the derived end result is expressed

j in terms of one or more measurement parameters. Call them A, B, C, and so on. You must | then make your measurements in such a way that each of these parameters. A, let's say is I

I

342
mass, has to be directly traceable, in a well defined series of steps, to the base system in

units. And similarly, if B is temperature, you have to know what systematic errors are involved in the making of that measurement, and so on. When you are finally done, you can then

say that your measurement is within some error bounds in terms of accuracy, or inaccuracy,

to be more correct. Using such a proven definitive method, you are now in a position to take a material and to build into that material the accurate values of whatever property or

properties are under consideration. From this we arrive at what we call here at NBS a Standard Reference Material. Now the rest of the world does not like this word "standard",

in the phrase "Standard Reference Material". It's a word very difficult to translate into

other languages. The accepted term now would be reference materials, in general, and what

we make at NBS would be called certified vefevenae materials. So we've got certified

reference materials (CRM's) and, more generally, reference materials. At this point, because

CRM's are directly related to the SI by a one-step process, we call these primary reference

materials. Now, this methodology is usually very difficult to apply in the typical clinical

laboratory situation. For example, we may use such techniques as isotope-dilution mass spectrometry which involves instrumentation running well over $150,000 or $200,000 plus very

highly trained technicians and scientists and so on. Therefore, we want to transfer this

accuracy in some way that is of a more practical nature. So now, we go from definitive

methods via the CRM to reference methods, which is what you've been discussing. Given then

a CRM and a reference method we now have the possibility of the manufacturing community

producing accurately known secondary materials, in a matrix that is really of interest.

Such secondary materials can be used for two purposes: (1) quality control on a daily basis

through these secondary materials, and (2) through the use the reference method and the

primary reference material to check in turn the accuracy of what I call the field or routine

methods used on a daily basis. This whole process has, of course, variations, depending

upon the particular scientific or technological field. In clinical chemistry, I have just

described the logical way to proceed toward achieving measurement compatibility. This

system is, however, just in its initial stages. Many more definitive methods, reference

methods, and CRM's need to be developed before major advances can come about.

Have I heard correctly that, although you may not all be in exact agreement, what you

would like from NBS would be a series of gases at the appropriate partial pressures of

oxygen and carbon dioxide in nitrogen as a first CRM. As a second CRM, perhaps a buffer

that's equilibrated with appropriate partial pressures of those two gases, and then preferably, blood or blood-like matrices prepared the same way.

Engel :

Right.

Cal i :

Have I got it about right?

Several voices :

Yes.

Cali:

You'll have it tomorrow!

Ladenson :

I'd like to ask Dr. Bates a question. You said you used Tris buffer to control pH. Did

you do any studies with Simon's calcium exchanger on the effect of Tris buffer? Both the

Orion and the Ruzicka exchangers will show apparent calcium binding to Tris buffer.

Bates :

We used Simon calcium exchanger, and our studies suggested that there is no significant

association between calcium and Tris buffer.

343
Ladenson

relationship.

Bates: ^i.^ the linpar relationship between the concentration of calc,

;tn'poss?b,r"tSrr4.f:n^iraSor?h,t L responsible for the slope being ,

than theoretical

No 0 nan:

1li

J- '"'-i-nreJ 'E.f ;%Sinr^s^:rorrbe\oi:?rt°b:t^:^^e z.,^

anything since then, but I th nk some or zu^ V extensive studies i, ^mmmm

correlate them to some normal ratio between activity and concentration is a question

s?r?:srdiusi "nnLS^:d=^r"bui levels. i^i--','-:-i-,,s^oi:s%"hrd

levels bv a total chloride measurement. Looking at them with these ^ ' ^^^^^^f

'deviatfonsVrom normality. And likewise variations in tota pr^^^^^^^^

to make electrolytes look differently, since they would directly affect the average

activity coefficient.

Cohen :

How do you interpret that?

Gambino:

Yes what could a physician do with that data? I think it's important to consi

activity and the electrode measurements, but I don't think enough work has been don^Je

fyoi

In (

ies

Jxpe

'ita

Noonan :

When you look at the total, the correlation is excellent. So that you could s we're qoinq to give you 140 on electrode and 140 on your flame photometer, but once i it; while that electrode is going to give you 130, or its going to give you 150. The d the studies have just never been done to demonstrate the usefulness of this approach. ! -

does seem to me that the physiological activity of these ions relating to the thernj

of ion movement across cells, etc., are basically a function of activity, and not q tion.

Weisberg :

Seligson did that about 10 or 12 years ago with Dahms and Rock, and they had ihi;, publication in Cliniaal Chemistry but it's just hanging in limbo because you can'tli'jto

your number to concentration. As Ray pointed out, the clinician is thinking in coi for electrolytes, and therefore, it may be a very good method, but you have now go factor to correlate it to the concentration.

i|5 re, Cohen : i "itei

. )

Did The issue is whether or not these are really telling us something that is goti/hyi

the patient.

344

trei
is old data suggesting that chloride is measureably bound to albumin.

hat is true. One of the things that was missing from my model, of course, was actions with albumin. I specifically left out chloride because I don't believe

ignificant binding at physiological pH. If you look back, Dahms used chloride

explain some anomalous chloride results that he got. If you look back to the

rk of Scatchard, and it's been confirmed on a couple of occasions, then it shows

lloride binding of albumin is markedly pH dependent, and at physiological pH, it's

igible. So I think that there is little justification for including chlorideding. In terms of CO2, of course, there's plasma carbamates which ought to have

d in the model that I presented, but I just didn't because the data wasn't readily

me.

Ik about chloride binding in the physiologic range, what magnitude are you

s far as I can remember, the thing that's reported is number of ions of chloride

|nd per molecule of albumin, and the number in the physiological range is less

here is less than one chloride ion bound per molecule of albumin. When you get

of around three (or perhaps five, I don't remember exactly) it goes up to about

ions per albumin molecule.

ersen:

nsi

the report by Dr. Brand very interesting but I wonder how you have estimated the

the protein polyelectrolytes on the ionic strength? Just taking the net protein

tration and assuming the influence equivalent to monovalent ions is questionable,

lated the ionic strength of normal plasma on the basis of the pK value for carbonic

fna as compared to pure aqueous solutions and found a value of about 0.17 mol/kg.

lyou have any comments to that.

e ionic strengths that I show are excluding protein. They are electrolyte ionic

In other words, I am following the convention of Scatchard, that the activities

tes in protein solutions are unaffected by the proteins. Now I don't think that's

experimentally by anyone other than Scatchard, so I can't comment, but that's a sr^ntative source. The model also, I should perhaps have commented, does not force

fility. It does not require that the system is in electroneutrality and it simply

the excess charge, whichever way, is made up by the protein. I think a model

?1oped forcing electroneutrality. It would be an interesting thing to do but I

lerni it, and it wasn't included in the model presented.

)t c

n'tcof

I go'

I

goi

tuggested that, as a final item, we have summary made of what's happened over

'days, and I didn't want to do this because I am very subjective about the coni[ disqualified myself. Since I wanted to get a disinterested person, I asked

g to provide us with a summary.

e really disinterested, I wouldn't have come. But there's an old Army saying,

nited States and I am sure in the army of every country, "never volunteer." I Dick a little earlier asking if a summary would be given and you've heard his

why he isn't giving it. I had suggested some names and some of those named

t reverted back to me again, you see; so here I am.

I 345
I apologize because if I had known I was going to do this, I would have listened mt better to you people. If I leave anyone's name out, please accept my apologies, and I cl' ask you please to contribute until the time of your leaving for the airport, if I have n|; any errors or if you object to what I say. I will honestly try to be objective and lea\/'

out my personal bias. I have to tell you what Dr. Laver said yesterday when I was kiddij'

him, having known him for some years, why he didn't contribute as much as I did by inter

ting, and he said he learned his lesson. He was at a conference several months ago and i spoke so much that he now has 15 pages to proof on his added comments; so I am stuck aga; on that. >

I think that we can have some subheadings for Dr. Rubin and the idea for the Expert

Panel. Number one would be the groups of the Definition of Quantities and Concepts to iK combined with the Recommendation of Nomenclature and Physiological Terminology and SymboipJ

As mentioned earlier, there is a phoenix that has arisen from 1964, and obviously we stij

have the same problems we had at that time. We have the problem of overlapping laborato

and clinical terminology. We have the problem of which parameters to use in reference tif the diagnosis of acid-base imbalance. Here, of course, I'm mentioning names. Dr. Cohen

introduced some additional terms regarding the duration and compensation of acid-base ' disturbances, and Dr. Engel introduced additional parameters, especially regarding the I'^^y^ titratable acidity.

I think the entire problem is more complicated in comparison to 1964, especially by! the introduction of the SI units. The United States is lagging far behind the other I countries, even though the other countries are still not quite uniform; for instance in i England, they are going to use hemoglobin as grams per deciliter because of the confusio'

that would result with the use of millimoles. We have been lagging for the following

reasons: number one, we're not in the metric system and, in reality, in order to get a metric system and SI units, we should be doing this with our school children rather than'

;are

idiis

lefa

:liet

lion, off

ifis ai

fi talking ex cathedra from this acy level to go down. We are really going to have our problem'

in the United States. In addition to this, we have a problem that we can accept it in Vf^^

laboratory; I calculated that there are nine practicing clinical chemists in the audienc

We can do it, but the fact is that the physician is not going to do it

In addition to the problem of SI units, we have the problem of in vivo and in vitro

diagnosis. The question of extra-cellular fluid regarding pH, especially, and blood gas

Several years ago when there wasn't much distinction, they had a controversy. I think m^"^'' and more today we're going over to the in vivo type of consideration and the clinicians ft'' going to have to learn to adjust.

we

i'tie r

ato »urs

ved

ed Our third area is the question of temperature correction. To quote two good areasDr. Laver is not here now and we can say that his is a good area, Massachusetts General,

the Mayo Clinic. They are routinely doing temperature correction for pH and blood gases^'^ys

all routine determinations as well as for the cardiac by-pass type of surgery. I think F'ist gave my stand yesterday on the plea that it should be done. The question of us today isr^e

what advantages would there be? We do not question as to what the normal would be at thi"'

patient's body temperature. We get the situation, such as Dr. Gambino has written aboutj

his monthly newsletter, of the patient coming in who has a PO2 of over 150 breathing roof

air; the physician says it's impossible, and when you check back, of course, that patienf

a low temperature. And finally, another area of controversy is the question of pH versujj

hydrogen ion concentration; and I won't make any comments about that.

Possible areas to be considered in greater depth by the subcommittees in the future

perhaps would be P50, or as our European friends call it T50, oxygen content, and 2,3-DPi

But there has been some agreement. Number one, we agree whether pH or hydrogen ion wouli

be the area of determination for acidosis/alkalosis or acidemia/al kalemia. The PCO2, or

if you wish, the carbonic acid concentration or total CO2 content by calculation would b^Jssi reported for respiratory factors. We still have our century of the Tower of Babel in the

' we still lack agreement on which metabolic term and which metabolic test to be used. 1 Whether the values should be reported per se, to use calculated values or those based on'M

the determined ones, should be looked into to see what can be done. As a subgroup of th

subcommittee A, is the question of the Evaluation of Nomograms and Algorithms. Here it was agreed, at least I think it was agreed, that there is a distinction between a diagrai

which would be a graphic representation for education and possible diagnostic help to the'|cti

physician and the nomogram which is a means of calculating additional parameters for thef

'

01

tan

346
a y — •'^ • J I • •• — — . -w • — i^w >, vvy V.I ^1 Ijr ^ I V liagram or a nomogram in the better centers, but I think it does have some area" of use

ivijhe boondocks, if you would, where they don't have

til

ratory personnel in order to report to the physician. Now it is not necessary to use 'I think it does have some area of us( ve quite the expertise as this group

today.

A subcommittee B would be entitled, with our subsections today, Blood Sampling,

jjling, and Storage. I think this is important enough to be a subsection of the proposed

rt Panel and it may need a restatement of the state-of-the-art rather than any further

^tigation. But you heard the comments and the- difficulties; as Dr. Ladenson showed,

b are indefinite or insufficient data but there have to be more definitive areas of

1 we're doing. I think here again, with my own personal bias, that for an evaluation

I2, you must have arterial blood, for the routine acid-base evaluation where you don't

a PO2, no matter how small that percentage, that venous blood would be perfectly

factory. This has been substantiated time and time again. Dr. Gambino, of course,

tdhe of the leaders in that area.

'DCfl

IS i)'

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or

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tl

it

gramli

tlie

A subcommittee C would combine, I think, the Establishment of Reference Values,

ity Control and Standards, and the Development of Reference Methods. And, as is Dus from this afternoon, this was the area of greatest interest, contributed no doubt

le fact that Mr. Cali gave his beautiful exposition of what the NBS does and so on. I here again, the problems came up with tonometry. Should we use blood, an aqueous

ion, or a glycerol solution as advocated by Dr. Maas? I think this requires a great

of expert input, and we would have to have that subcommittee do that. The question

Vpe of tonometer, and I apologize to Dan ahead of time, the simple Noonan tonometer,

simple Noonan, but the simple Noonan-tonometer, versus the more complex ones. I think

s agreed that the NBS should work in the areas of certifying gas mixtures for standardin for tonometry and for standardization of the instruments with a probable ultimate

euifacy within one percent. It was also agreed, I believe, that the NBS should work in t|area of certifying buffers or buffer compounds for buffers of pH. I think the suggesthat was made is extremely important that these buffers be certified at 37 °C and we bracket the 7.40. It is crazy to have a pH in the low sixes and under 6.8, and so

ve really should be bracketing and my own preference would be somewhere about 7, 7.4

7.6 or a 7.7 in that area. It was agreed again to use an operational definition for

jjslAs far as PCO2 and PO2 is concerned, there was some discussion for mass spectrometry,

iner this is a definitive method or is this really a secondary method for accuracy. A

,inetric Scholander-Riley technique would be the one to be utilized in the reference

atory. The problem came up of the question of normal levels for PCO2 and PO2 and,

)urse, that evolved itself with a give and take, but we won't mention the people

ved. I believe, that an abnormal level, as it seemed to be the consensus, should

;ed primarily to check out instruments rather than for the beneficial knowledge to )hysician for diagnostic levels. Dr. Gambino reported on some data using General

lostic's new aqueous acid-base control which contains PO2 values, and this would

lately replace the one that they have now which does not have a PO2 and is a lyophilized

'ial, which I think has a great many problems attacked to it. If the liquid material

;! out well, as I think it will, it should be substantiated and supported. My own

iistion to Dr. Durst is that I think there have been enough people working with it,

iambi no, Mr. Komjathy, and myself, that this should be included in the final mono1, because you don't expect a company to publish somebody's in-house, out-of-the-house

of work but with the monograph I think it will give some good data to show on that

of support.

The next area of section D is the Instrument Specification and here, of course, we

'the best available and then beyond that. I think the consensus that everybody would

5 to is that we need better service manuals for maintenance and repair at all hours,

ksible. But what we're really saying is we need a VIP instrument; VIP instrument

hs for Very Idiot Proof. I think the manufacturers should be listening, as they were especially to that Section C with the Reference Methods and Standards, because even

ph they are in the business to make money, we also have to have some of the accuracies

on. The future problem, as emphasized by Prof. Kreuzer, is the question of the in

monitoring. Here, of course, with a PO2 electrode that he described, micro and so an be extremely convenient. It was extremely interesting to see this new PCO2

ictivity electrode because this is now opening some new parameters whereby we may make

advances.

347
And finally, we have the Electrolytes. And this I find a little bit good and bad

because it's sort of an afterthought. We had 2 days on the other groups and had only 2 hours for electrolytes. This is like the New Orleans type of talk of lagniappe, something

extra for nothing. It came out very well because this is leading into my conclusion. And

by taking it out of context, Dr. Brand gave a very nice talk and his comment that reporting

the values as per liter of plasma water is extremely important. I think this has been

suggested many times in the past. I was very pleased to hear Dr. Bates talk about sea water. Number one, because he is the other bald-headed, white-haired man in the audience,

and number two, the sea water study that he is working with now takes me back to 1916 when

the diagram (number 3 in my list) by McClendon also worked on sea water. So really we've

gone full cycle now for another 60 years.

I wish to thank Dr. Durst for the honor of summing up the deliberations and, on behalf of all the participants, to congratulate him on being the honorary obstetrician for

the birth and/or delivery of this conference. I hope that the monograph will be as successful as the one on Ion-Selective Electrodes.

348
I!BS-1 14A (REV. 7-73)

U.S. DEPT. OF COMM. 1 . PU B L IC ATION OR R E PORT N 0 . 2. Gov't Accession

1 BIBLIOGRAPHIC DATA No.

1 SHEET NBS-SP 45 0

3. Recipient's Accession No.

i. TITLE AND SUBTITLE

j Blood pH, Gases and Electrolytes

j Proceedings of the Workshop on pH and Blood Gases held

I at the National Bureau of Standards 7/7-8/75

5. Publication Date June 1977

6. Performing Organization Code

. AUTHOR(S)

1 Richard A. Durst, Editor

8. Performing Organ. Report No.

PERFORMING ORGANIZATION NAME AND ADDRESS

NATIONAL BUREAU OF STANDARDS

DEPARTMENT OF COMMERCE

WASHINGTON, D.C. 20234

10. Project/Task/Worlc Unit No.

11. Contract/Grant No.

2. Sponsoring Organization Naipe and ComDlete Address (Street^ City, State, ZIP) American Association for Clinical ChemistryAmerican Society of Clinical Pathologists

International Federation of Clinical Chemistry

National Committee for Clinical Laboratoiry Standards

TT,S. Department of Commerce, National Bureau of Standards

13. Type of Report & Period Covered Final

14. Sponsoring Agency Code

5. bOFPLEMENTARY NOTES

Library of Congress Catalog Card Number: 76-608179

6. ABSTRACT (A 200-word or less [actual summary o( most significant information. If document includes a significant

bibliography or literature survey, mention it here.) On July 7-8, 1975, a workshop was held at the National Bureau of Standards to discuss the status and needs of this very important area of clinical measurement,

' A major goal of this workshop was the initiation of cooperative efforts on an international level toward the standardization of pH and blood gas measurements and the

various quantities and terms used in this field.

To this end, the first technical session was concerned with the acid-base status

of blood and included the topics: Definitions of Quantities and Concepts; Recommendations of Nomenclature, Physiological Terminology and Symbols; Reference Values;

and the Evaluation of Nomograms and Algorithms. The second session addressed

itself to the more practical aspects of this subject and included the topics: Blood

Sampling, Handling, and Storage; Instrument Specifications; Quality Control and

Standards; and the Development of Reference Methods. Finally, a brief session on

the newer topic of the electrometric measurement of blood electrolytes.

This volume contains all of the papers invited for presentation at the workshop by

some of the leading clinical and medical authorities on this subject and also includes

a transcription of the extensive discussion sessions.

17. KEY WORDS (six to twelve entries; alphabetical order; capitalize only the first letter of the first key word unless a proper name; separated by semicolons

)

Acid-base status; blood electrolytes; blood gases; blood pH; calcium; carbon dioxide;

hydrogen ion concentration; nomograms; oxygen; ^ qq > pH; Pq ; potassium; sodium.

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