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PUBLICATIONS

A Critical Review of Measurement Practices for the

Determination of pH and Acidity of Atmospheric

Precipitation

George Marinenko and William F. Koch

Center for Analytical Chemistry

National Bureau of Standards

Washington, D.C., 20234

This review surveys current literature on the measurement of pH and acidity of atmospheric precipitation.

Current practices for calibrating pH measuring systems

for atmospheric precipitation applications are reviewed and

possible sources of error are discussed. Determinations of

acidity are grouped in accordance with the type of end-point

selected for titration: color indicator, fixed pH, Gran

plot, and closed loop.

Key words: acidity, acid precipitation, acid rain, acidimetric

end-point, pH determination, pH standards, pH measurement errors,

titrimetric errors.

INTRODUCTION

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The purpose of this paper is to review the procedures that are

being used to measure two of the more controversial rainwater parameters ,

pH and acidity. It is generally recognized that even uncontaminated

rainwater is naturally acidic because of the dissolved atmospheric

carbon dioxide. In the absence of any acidic components other than

atmospheric CO^, the pH of rain would be 5.7 [1,2]. However it is

unlikely that this laboratory ideal is ever encountered in natur> ,

•

in remote areas, due to the natural sulfur and nitrogen cycles.

The intense concern in the last decade over acid precipitai;

not focused on the innate acidity of the precipitation, but ratia-r

the acid-forming components other than C0 9

. Of particular

the anthropogenic components, which could lead to irreversible s

in the natural balance [1]. Many of these components are derived from

combustion processes and are dependent on the nature and origin of the

fuel used. The most prevalent among the acidic components of combustion

products are oxides of sulfur, oxides of nitrogen, and hydrogen chloride.

The current literature contains many generalizations and conclusions

about the trends in precipitation acidity, its possible sources (and

origin), as well as projections of the environmental effects in the

future [1], Such conclusions, however, must be based on accurate analytical

data. Consequently this literature review was undertaken to evaluate

critically the published procedures for the measurement of pH and acidity

of atmospheric precipitation, and to assess their comparability.

BACKGROUND

Only in very special cases can the two parameters, pH and acidity,

be considered interconvertible. In the majority of cases, the pH of a

solution cannot be readily reduced to hydrogen ion concentration (acidity)

without some knowledge of the other ionic components and their concentrations in the solution, as well as an understanding of systematic errors

inherent in the measurement process. Conversly the knowledge of the

total acid content of the solution is not necessarily sufficient to

calculate its pH. Both of these parameters, however, play an important

part in defining the state of a system.

Formally defined, pH is the negative logarithm to the base 10 of

hydrogen ion activity. It is related by some to "free acidity", or the

concentration of "free" hydrogen ions in solution [3]. Often it is

conceptually defined as the negative logarithm of the effective concentration of hydrogen ion. It is known that in dilute solutions, thi .

2

effective concentration or activity is always less than the total concentration of hydrogen ions. As the concentration approaches infinite

dilution, the two quantities (activity and concentration) converge.

Total acidity, on the other hand, is the measure of the total

concentration of hydrogen ions both "free" and undissociated (provided

they are "acidic" hydrogen atoms) [3]. Generally, total acidity is

determined by acidimetric titration with strong base.

In other words, pH is an intensive factor, representing the free

acidity of the solution, whereas the total acidity is an extensive

factor, representing the acid capacity of the system. pH is generally

the more important factor in chemical equilibria and in kinetics; total

acidity dominates in stoichiometric considerations.

REVIEW OF pH MEASUREMENTS

Measuring pH in precipitation samples may at first appear to be a

straightforward procedure. In reality, differences in the metrological

details of the measurement can very easily produce widely variant

results, often greater than 0.5 in pH. Current measurement practices

used by acid rain researchers are quite diverse. A summary of these

practices is presented in Table 1. Many articles concerned with acid

rain studies make no mention of the pH measurement details but rather

report only the measured pH values. Since no conclusions can be drawn

from such articles regarding metrological validity of the procedures,

those references are not included in this review. Many of the reported

procedures are quite incomplete as can be seen from Table 1.

As evidenced by the data assembled in Table 1 there is no unit orm i tv

in the procedures used for pH measurements in acid precipitation me.i

ments. While in some cases the accuracy of the measurement is at least

3

considered [4], data supporting the accuracy claim are lacking. In some

cases it is difficult to establish exactly what measurements were performed.

The following is an example of this: "Precipitation pH was determined

by calibrated electrode meters" [5]. This is the only statement describing the pH measuring instrumentation. The use of unfamiliar terminology

(what is an electrode meter?) and the lack of experimental detail renders

the above report virtually useless when assessing the reliability of the

data.

There is a point to be made regarding the precision and accuracy of

pH measurements. Several authors report the precision of their measurements with an implication that the accuracy is of similar magnitude.

However, measurement accuracy is determined by two types of error

components or uncertainties: systematic and random. In testing the

repeatability of pH measurements, (i.e. the precision, generally expressed as the standard deviation) , one is testing only one of the two

error contributions, namely the random error component. Included in the

systematic component are measurement biases, often quite subtle in

nature and unlike standard deviations, not as amenable to mathematical

computation.

The systematic error assessment in pH measurement is complex. One

factor contributing to the difficulty of assessing systematic errors is

the presence of the liquid junction between the solution to be measured

and the reference electrode. The presence of this junction generates a

small potential across the boundary. In itself this would not be a

serious problem, since this potential would be incorporated into the

calibration measurements with a standard reference solution ot known pH.

Unfortunately liquid junction potentials are not constant and ire de; en4

dent on the nature of solutions involved. Thus, calibrating the pH

measuring electrode system with a standard buffer solution of known pH

does not guarantee the accuracy of pH measurements on the unknown. Even

under ideal conditions, these biases could be as large as 0.1 in pH,

depending on the buffer used for calibration as well as the nature,

ionic strength and the pH of the unknown solution.

To make matters even more complicated, the magnitude of these

liquid junction related errors is also dependent on the type of liquid

junction involved (e.g. ceramic plug, flowing, etc.), on the condition

of the junction, and on the history of the electrode (types of solutions

tested with the electrode)

.

For these reasons it is advisable for the electrode systems used in

acid rain pH work to be used exclusively for this task and to be monitored with respect to their performance with known reference solutions

intended to simulate the unknown samples. However, for documentation

purposes, calibration with the certified buffer solution is essential.

In addition to liquid junction potential errors, streaming potential

errors can be produced by stirring or agitating the test solution during

the pH measurement. The magnitude of this streaming potential depends

on several factors: (a) type of reference electrode junction; (b) rate

of stirring; and (c) ionic strength of solution. Galloway [6] and Koch

[7] have observed that stirring low ionic strength solutions can produce

apparent changes as large as 0.5 pH. To eliminate such errors, it is

generally recommended that pH measurements be made in quiescent solutions

[6,7,8].

Galloway, Cosby and Likens [9] point out difficulties associated

with accurate electrometric measurements in low ionic strength solu5
-
tions, such as atmospheric precipitation. Their calculations show that

if activity coefficients, liquid junction potentials, streaming potentials and non-Nernstian behavior of the electrode system are ignored,

errors of 50% or more may be encountered in calculating hydrogen ion

concentration. According to Galloway et al., residual liquid junction

potentials, arising mainly from the differences between ionic strengths

of the calibration buffer solutions and those of the samples, can produce

errors as large as 0.04 pH. Koch, Marinenko, and Stolz observed even

larger residual liquid junction potentials in a systematic study of

various junction materials [7],

In 1979, in an attempt to reduce the effect of residual liquid

junction potentials, Galloway et al . [9] recommended using solutions

made from strong acids for calibrating the glass/reference electrode

system. The use of dilute solutions of strong acids along with a double

junction reference electrode has been recommended by Liljestrand and

Morgen [10] for calibrating the glass electrode system. The use of

dilute acid calibrants tends to minimize the junction potential effects.

In such a calibration approach the structure of the liquid junction

between the reference electrode and calibrant solution matches more

closely the junction between the reference electrode and the unknown.

Three years later, in 1982, Galloway et al . [6] suggested a different,

procedure for pH measurement. Instead of using strong acid solutions t

calibrate the pH meter, a two point calibration procedure using pH '.0<

and pH 4.00 buffers is used. The performance of the pH measurement

system was periodically checked using a dilute H^SO^ solution '

concentration. This is in keeping with the recommendati |, ns t •

Marinenko [11] . A similar procedure was developed by Peden ft :

.

6
"
for the U.S. Department of Energy. In their procedure, pH 7 and pH 4

buffers are used to calibrate the pH meter and a 2 min equilibration

period is required before the pH measurement is made.

Some researchers use an inert atmosphere such as nitrogen [13] to

shield the solution from atmospheric CO^. Shielding, however, is not

the only effect of such a practice. Gas/solution exchange processes are

relatively fast and therefore nitrogen shielding will remove some or all

of the dissolved CC^, depending on the time involved and the pH of

solution. Below pH 5 removal of CC^ is rapid while at pH 7 it is rather

slow.

Some procedures invoke the British Standard [14,15] which recommends adding CuSO^, as a fungicide, to rainwater samples at the time of

collection. Adding this salt will invariably affect the acid-base

equilibrium.

In one instance, a colorimetric method was employed [16] for the

determination of pH in single drops of rain. Colorimetric pH indicators

were sprayed on paper and calibrated with solutions of known pH. This

is, however, a special application. In general, the interest in pH of

rain is on the bulk scale rather than on the drop size scale. The

colorimetric determination of pH on the bulk sample may have some merit.

However, a narrow range, accurate, calibration scale of sufficient

sensitivity must be developed. Furthermore, colorimetric procedures ire

generally more involved and time-consuming than electrometric pruredur- .

There appears to be one point of consensus: the majority ot researchers who use electrometric measurements prefer combination elo -

trodes [4,7,17,18,19,20].

7

The diversity of pH measurement practices is quite apparent, which

may be partially or wholly responsible for some of the conflicting

conclusions of researchers in the area of acid precipitation. A more

uniform approach to measuring pH in precipitation is desirable to eliminate

some of the ambiguities. A self-consistent system should be developed,

which will require specific standard solutions for calibration, and will

specify the minimum performance characteristics of the electrodes.

TOTAL ACIDITY

The measurement of total acidity generally involves titrating acids

in the sample with a standard solution of sodium hydroxide. In some

cases, the standard base is generated coulometrically in situ . The

methodologies employed in acid precipitation measurements of total

acidity are summarized in Table 2. The principal distinction between

these different experimental approaches is the manner in which the

titration end-point is determined.

Three different end-point techniques are used:

(1) visual, using color indicators

(2) electrometric to a fixed pH

(3) electrometric using Gran plot [21]

The simplest and most convenient technique for field work requires

a color indicator [15]. Corrections must be made for the difference

between the volume of titrant consumed to cause the color change of the

indicator and the volume required to reach the equivalence point.

It is quite common to use a glass electrode while titrating t >

fixed pH. Examples of this technique are illustrated in Table 2

'

entries 5, 6 and 7. The choice of the end-point ranges from pH 7 t

8

pH 9.5. Depending on the end-point pH value, the amount of titer will

vary. First, there is a certain amount of reagent required to change

the pH of pure water from the neutrality point to some pH value above

it, the end-point. Naturally, the higher this end-point pH value, the

more reagent will be required.

Second, if the end-point pH value is too low, not all acids present

will be titrated. Only the strong acids will be completely titrated at

pH 7; the weak acids will be partially titrated. Carbonic acid is one

of these weak acids. In all these procedures, valid corrections must be

applied to reconcile the differences between the end-point pH and the

equivalence point pH. The existing nonuniformity in establishing the

end-point pH is perhaps the single most serious variable in total acidity

determinations. There is little doubt that the total acidity value

obtained when titrating a sample to pH 7, will be quite different from

that obtained by titrating the same sample to pH 9.4 (-Methods 6 and 7,

Table 2)

.

The Gran plot technique requires that pH measurements be made after

several (well spaced) aliquots of standard NaOH solution are added to

the sample. Then dV/dpH=f(V) can be calculated. This function is then

extrapolated to the abscissa. The use of a Gran plot is one convenient,

way of manipulating titration data and may enable one to determine the

strong acid contribution to the total acidity.

The closed loop technique provides a rapid and convenient way :

obtaining the total acidity [7]. In this technique, an electrolyte

solution is first adjusted to some convenient pH value before a measure i

amount of sample is introduced and titrated coulometr ica 1 1 y to the ;am«-

pH. A convenient system involves the use of 0.01 mol/L . di-m \-

solution as the electrolyte.

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«
CONCLUSIONS

Measuring pH and total acidity in low ionic strength solutions,

such as acid precipitation, is a formidable task involving intrinsic

difficulties associated with the measurement system (such as residual

liquid junction potential, and choosing an equivalence point, respectively) even under a relatively constant set of conditions. If measurement conditions vary or if they are not given, as is indicated by the

reviewed literature, the matter becomes complicated even further, to a

point where no significant correlation can be made between the results

of different experiments.

A uniform approach to the determination of pH and total acidity of

precipitation is necessary to make all measured values comparable. This

approach must include a sampling protocol, sample storage and handling

conditions, traceability to national calibration standards (such as

Standard Reference Materials issued by NBS) , minimum performance requirements of all measurement instrumentation, and a valid statistical evaluation.

ACKNOWLEDGMENT

Portions of this work were funded by the United States Environmental

Protection Agency, Contract Number IAG-AO-13-F-2-535-0 , Subagreement 3.

10

Table 1. pH Measurement Practices.

Method/ Instrumentation Calibration Remarks Ref

1. Electrometric/Sargent- - Accuracy [4]

Welch PBL combination claimed

electrode +0.03 pH

2. Electrometric/Corning

12 or Orion 801

instrumentation lab. combination, Corning glass

and Orion 90-01-00 reference electrodes

atmosphere [17]

3. Electrometric/SargentWelch PBL combination

electrode

[18]

4. Electrometric/glass "dilute acid [10]

electrode with double

junction reference

electrode

solutions"

5. Electrometric/Corning Fisher certi- [19]

12 pH meter with fied pH 4.00

combination electrode and pH 7.00

buffers

6. Colorimetric/ indicators "calibration with [16]

sprayed on paper solutions of known

pH"

11

Table 1 (Continued)

.

Method/ Instrumentation Calibration Remarks Ref

7. Electrometric/Orion

combination electrode

8. Electrometric

9. Electrometric

10. Electrometric/combination electrode

with ceramic junction

standard buffers

pH 4.0 and

pH 9.2

standard buffer

pH 4.00 and

pH 7.00

Dilute solutions

of strong acid

pH 4.004 NBS buffer

pH 6.863 NBS buffer

[ 20 ]

standard devi- [12]

ation of measurement is

0.05 pH

[ 6 ]

[7]

Certain commercial equipment, instruments, or materials are identified

in this report to specify adequately the experimental procedure. Such

identification does not imply recommendation or endorsement by the

National Bureau of Standards, nor does it imply that the materials or equipment identified are necessarily the best available for the purpose

12

Table 2. Total Acidity Measurements.

Titration method End-Point Remarks

1 . Microtitration of 5mL Electrometric/

sample with NaOH Gran Plot

solution

2. Conventional NaOH Electrometric/

titration Gran Plot

3. Conventional NaOH Electrometric/ Resolution of

titration Gran Plot strong and weak

acid contributions

4. Coulometric generation Modified Gran plot [

of OH" or closed loop

5. Conventional NaOH pH = 9 CO

2 has large

titration effect on acidity

6. Conventional NaOH pH = 9. 4-9.

5

titration

7. Titration with pH = 7.00 N

2 purging to re5.10

_3 mol/L NaOH move CO^: NaOH

standardized with

KHP

8. Conventional NaOH color indicators;

titration electrometric/

Gran Plot

9. Coulometric titration closed loop pH 8.3

Ref

.

[ 13 ]

[ 22 ]

[

29 ]

[ 10 ]

]

[

23 ]

[

24] [27

[ 3 ]

[

4 ]

[

18 ]

[ 17 ]

[ 19 ]

13

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[7] Koch, W. F.; Marinenko, G. ; Stolz, J. W. Simulated Precipitation

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[8] Madsen, B. C. Atmos . Environ . , 1981, J_5, 853~862.

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,

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[10] Liljestrand, H. M. ; Morgen, J. J. Environ . Sci . and Tech . 1973, 1 1 ,

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[11] Koch, W. F.; Marinenko, G. Proc . Sympos . on Sampling and Analysis

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.

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[15] BSI British Standard 17^7, Methods for the Measurement of Air Pollution, British Standard Institution, London, 1969.

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[19] Hendry, C. D.; Brezonik, P. L. Environ . Sci . and Tech . 1980, 1

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[20] Scott, W. D. Atmospheric Environment , 1977, 1

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[21] Gran, G. Acta Chem . Scan., 1 950, _4, 559-577.

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[23] Seymour, M. D.; Schubert, S. A.; Clayton, Jr., J. W. ; Fernando, J.

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[25] Frohliger, J. 0.; Kane, R. L. Proc . First Internat . Sympos . on Acid

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[26] Klockow, D.; Denzinger, H. Proc . First Internat . Sympos . on Acid

Precipitation and the Forest Ecosystem , 1975, 157-158.

[27] Krupa, S. V.; Coscio, Jr., M. R. ; Wood, F. A. Proc . First In- •rr.v .

Sympos. on Acid Precipitation and the Forest Ecosystem , 1

,

[28] Cogbill, C. V. Proc . First Internat . Sympos on Acid Precipitation and

the Forest Ecosystem, 1975, 363-370.

15

[29] Brosset, C. Proc . First Internat . Sympos . on Acid Precipitation and the

Forest Ecosystem , 1975, 159-179.

[30] Esmen, N. A.; Fergus, R. B. The Science of the Total Environment , 1976,

6 , 223 - 226 .

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16