WRD LIBRARY COPY 

U. S. GEOLOGICAL SURVEY

Methods for Collection 

and Analysis of Water 

Samples

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1454

LIBRARY
Methods for Collection 

and Analysis of Water 

Samples

By F. H. RAINWATER and L. L. THATCHER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1454

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1960
UNITED STATES DEPARTMENT OF THE INTERIOR 

STEWART L. UDALL, Secretary

GEOLOGICAL SURVEY 

Thomas B. Nolan, Director

The U.S. Geological Survey Library has cataloged this publication as follows:

Rainwater, Frank Hays, 1921 

Methods for collection and analysis of water samples, by 

F. H. Rainwater and L. L. Thatcher. Washington, U.S. 

Govt. Print. Off., 1960.

ix, 301 p. illus., diagrs., tables. 24 cm. (U.S. Geological Survey. 

Water-supply paper 1454) 

Includes bibliographies.

1. Water Analysis. I. Thatcher, Leland Lincoln, 1923 

joint author. II. Title.

REPRINTED 

1962

1965

For sale by the Superintendent of Documents, U.S. Government Printing Office 

Washington, D.G., 20402 - Price $1.50 (paper cover)
PEEFACE

The Geological Survey has the responsibility for measuring and 

evaluating water moving through that portion of the hydrologic 

cycle between the time the water from the atmosphere reaches the 

earth's surface and the time it is returned to the atmosphere or enters 

the ocean. During this part of the cycle, water may appear in many 

environments and under different conditions, and a full understanding of the problems of hydrology requires the application of various 

specialized scientific techniques. The Quality of Water Branch of 

the Water Kesources Division determines and appraises the chemical 

and physical characteristics of the Nation's water resources.

These activities include (1) systematic collection, compilation, 

and evaluation of basic data relating to the chemical and physical 

quality of surface and ground waters as required for the. development and utilization of industrial, municipal, and agricultural water 

supplies; (2) research and development studies to improve investigational techniques; (3) fundamental studies of the occurrence and 

significance of dissolved and suspended substances; and (4) preparation of results of water-quality research and investigations for 

publication. This manual covers only the data-collection segment 

of the work between and including the selection of the sampling site 

and the completion of the laboratory determination of water characteristics attributable to the presence of dissolved material. Other 

phases of chemical-quality investigations are discussed by J. D. Hem, 

1959, in "Study and Interpretation of the Chemical Characteristics 

of Natural Water," U.S. Geological Survey Water-Supply Paper 

1473.

This manual was prepared by the U.S. Geological Survey, Water 

Resources Division, under the immediate supervision of W. F. White, 

chief, Chemical Quality Section of the Quality of Water Branch. 

Many associates contributed helpful suggestions and reviewed the 

analytical procedures.

CONTENTS

Page

Preface_ _ _--__--_--__-_____________-___--___-_-__-_----__-____-_ in

Abstract_ .______________________________________________________ 1

Introduction._____________________________________________________ 1

Section A. Collection of water samples._______________-_-____--___--_ 3

Sampling-site selection (A:l)__-______--____-__--___-_---_--_-_ 3

Selection of surface-water sampling site (A: la)________________ 3

Selection of ground-water sampling site (A:lb)________________ 6

Sampling frequency (A:2)______________________________________ 7

Sampling equipment (A:3)_____________________________________ 8

Sample containers (A:3a)___________-_________-____-_-_____- 8

Apparatus for sampling water that contains nonvolatile constituents unsusceptible to aeration (A:3b)__________________ 10

Apparatus for sampling water that contains dissolved gases and

constituents susceptible to aeration (A: 3c) __________________ 13

Sampler suspension apparatus (A:3d)________________________ 15

Sampling instructions (A:4)_____________._______--__-_-_-_--___ 15

Safety precautions (A:4a)__________________________________ 15

Collection of representative samples (A:4b)______-___---___-__ 20

Tabulation of sample data (A:4c)___________________________ 20

Special treatments, preservatives, and fixation (A:4d)________ 27

Section B. Handling of water samples before analysis.________________ 31

Time between sample collection and analysis (B:l)_______-______ 31

Sample-bottle cases (B:2)--_____--__--_-_----------_--_-------- 32

Sample storage (B:3)--__---____--__---_---_-------.---__------- 32

Preparation of composite samples (B:4)__________________________ 32

Time-weighted composite samples (B:4a)_____________________ 36

Discharge-weighted composite samples (B:4b)________________ 37

Removal of turbidity (B:5)______-_-__--__----____--_-----_----_ 39

Section C. Analysis of water samples_________________________________ 43

Analytical techniques in water analysis (C:l)_____________________ 43

Spectrophotometry (C:la)________________________________ 43

Terminology and procedures (C:la-1) ___________________ 44

Applications of Spectrophotometry to water analysis

(C:la-2)--_____-___-___---_____---__---_----------_ 47

Flame photometry (C:lb)__________________________________ 55

Instruments (C:2)____--________________---___-----_----_---_ 58

Conductivity meters (C:2a) ___-__--___--___-__-__---__----_ 58

Balances (C:2b)______-____________________________--_-_-_- 60

pH meters (C:2c)_____.__________.______________.____-_ 60

Spectrophotometers (C:2d)_________________________________ 61

Flame photometers (C:2e)____________-____-_---_--_------ 68

Water-color comparators (C:2f)_____________________________ 69

Turbidimeters (C:2g)__...__________.___-__---..---.-_----- 70

Photometric titration assemblies (C:2h)____________-_____---- 71

Glassware and other containers used in the laboratory (C: 3) -------- 72

Choice of analytical methods (C:4)_____________-_-_____--_-----_ 75

Explanation of terminology (C:5)_-________---_.______-_-_------ 75

v
VI CONTENTS

Section C. Analysis of water samples Continued Page

Scheme of analysis (C: 6) _._____ ______________--___---__-_____ 78

Significant figures for reporting results (C:7)____________________ 79

Evaluation of the accuracy of analytical results (C:8)____________ 81

Chemical balance (C:8a)___________________________________ 81

Relation of residue on evaporation to calculated dissolved solids

(C:8b)_________________________________________________ 83

Specific-conductance relation (C:8c)___________--_---__--____ 84

Interrelation of constituents of water from the same source

(C:8d)_________________________________________________ 85

Method for checking partial analyses (C:8e)__________________ 86

Section D. Analytical procedures___________________________________ 87

Acidity (D:!)____________________________________________ 87

Total acidity (D:la)___________________________________ 88

Volumetric method (D:la-l)__-----------------------_- 88

Immediate acidity (D:lb)____________________________ ______ 90

Volumetric method (D: lb-l)________________---________ 90

Potential free acidity (D:lc)__-___--___---__----_--------___ 91

Volumetric method (D:lc-l)____________-_-_-_-_---____ 91

Alkalinity (D:2)______________________________________________ 93

Potentiometric method (D:2a-l)____________________________ 94

Aluminum (D:3)____-___-_-_____-_-_-______-__---_---_---_-__- 97

Ferron-orthophenanthroline method (D:3a-l)^________________ 97

Arsenic (D:4)_______________________________________________ 101

Gutzeit-volumetric method (D:4a-l)________________________ 101

Barium (D:5)______________________________________ 107

Gravimetric method (D:5a-l)____________-_---_-_---___-___ 107

Complexometric method (D:5a-2)_-___-_--__-__--------_--__ 109

Boron (D:6)~ __ _._ ________ _ _____ _ ___ _ ____ 113

Dianthrimide method (D:6a-l)_____________________________ 113

Potentiometric method (D:6a-2)____________________________ 115

Carmine method (D:6a-3)____________-_-_-____--___-_-___ 118

Bromide (D:7)________________________________________________ 121

Oxidation method (D:7a-l)_-___----------_---_---_-------- 121

Bromine displacement method (D:7a-2)____________________ 124

Calcium (D:8)__-___________________________._ __.____ _ ___ 127

Complexometric method (D:8a-l)_-_-------_-------_-_------ 127

Permanganimetric method (D:8a-2)_______________-___---_-- 129

Gravimetric method (D:8a-3)____---------_---_-----_-_---- 131

Turbidimetric method (D:8a-4)___________________________ 134

Carbon dioxide (D:9)_-_______-__-___-_-_---_-_---__--_-_------ 137

Calculation method (D:9a-l)_________________-___-_-_-_____ 137

Chloride (D:10)__________________________________ 141

Volumetric method (D: 10a-l)_______________________________ 141

Gravimetric method (D:10ar-2)___---_---------------_------ 143

Chlorine residual (D: 11). ____________-__________--_----------_- 147

Orthotolidine-arsenite method (D: 1 la-1)_______-_-_----__-___ 147

Chromium (D:12)____________________________________ 151

Hexavalent chromium (D: 12a)__-_-_---_---------_-_-------- 151

Diphenylcarbazide method (D: 12a-l).___________________ 151

Chromium (D:12b)_____________________________ 153

Permanganate-azide method (D: 12b-l)__________________ 153
CONTENTS VII

Section D. Analytical procedures Continued Page

Color (D:13)_________________________________________________ 155

Comparison method (D: 13a-l)_______--_--_-__-___--_______ 155

Copper CD:14)________________________________________________ 157

Diethyldithiocarbamate spectrophotometric method (D:14a-l)___ 157

Density (D:15)_______________________________________________ 161

Gravimetric method (D: 15a-l)____.___._..__.__.__.__._____ 161

Fluoride (D:16)_ _._._.___.___.__-__._._.._...........___.. 163

Zirconium-Eriochrome Cyanine R method (D: 16a-l)__________ 163

Zirconium-alizarin method (D: 16a-2)._._..__-__.___..__._.__ 166

Isolation of fluoride by distillation (D: 16a-3).-_-----_-----.__ 169

Hardness (D:VT).............................................. 173

Hardness (D:17a).._..____._...___._._.._.._.__.._..__.._. 174

Complexometric method (D: 17a-l)____-_..___.__.__..___ 174

Calculation method (D:17a-2)________________ 176

Carbonate and noncarbonate hardness (D: 17b) _______________ 177

Calculation method (D:17b-l)___.____.________ 177

Iodide (D: 18) -------------_-___-_--____-_-_---_-----_---___ 179

Oxidation method (D; 18a-l) _ _ _____________________________ 179

Starch method (D:18a-2)___.______________________________ 181

Iron (D:19)__________________________________________________ 183

Iron (D:19a)_____.._____._..__..____._._.._...._._....__. 183

Bipyridine method (D:19a-l)_.______._.-__.__-_-_-. 184

Ferrous and ferric iron (D: 19b)___________---____-_-___-_--_ 186

Bipyridine method (D: 19b-l)-_.-___--________.___---- 186

Lead (D:20)_. ................................................ 189

Dithizone method (D:20a-l)_________._..____________ 189

Lithium (D:2V)............................................... 193

Flame-photometric method (D:21a-l)_._.____.___.__._.____- 193

Loss on ignition (D:22)_.___._._._..._..__.___..__.__-_--._---_ 195

Gravimetric method (D:22a-l)...._..___..-._______-.-._.-- 195

Magnesium (D:23)_______________________________________._. 197

Calculation method (D:23a-l)--__--___--_---_-_--__-------- 197

Pyrophosphate method (D:23a-2)_-.____-----_._-_---------- 198

Eriochrome Black T method (D:23a-3) ______________________ 201

Manganese (D:24)_____..._______.-.____-__.-_____________ 205

Permanganate method (D:24a-l)..____.-.__-___-____-_-__-- 205

Tetrabase method (D:24a-2)._____________.____.___... 207

Nitrogen (D:25)_.___._.__-__.___..___..______......___.__.... 211

Ammonia nitrogen (D:25a)____.___.__.________.__-__-_----_ 211

Nesslerization method (D:25a-l)_.._____._.___--__-__- 211

Titration method (D:25a-2)_____.__..._.__....___.. 214

Test-paper method (D:25a-3)___...______________ 215

Nitrate nitrogen (D:25b)__.____.___________________ 216

Phenoldisulfonic acid method (D:25b-l) _____.____-____-- 216

Reduction method (D:25b-2)________________... 219

Nitrite nitrogen (D:25c)______________________.____ 221

Diazotization method (D:25c-l)____________._--__-___-- 221

*J Organic nitrogen (D:25d)_________________________.. 223

Kjeldahl method (D:25d-l)....__...... .__.-.... .- 223

Oils and waxes (D:26)__________________________________.______ 227

Wet-extraction method (D:26a-l)____.___________---_-_- __-- 227
VIII CONTENTS

Section D. Analytical procedures Continued Page

Oxygen consumed (D:27)___________.__________________________ 229

Permanganimetric method (D:27a-l)______-_-_--_-_______-__ 229

Oxygen dissolved (D:28)_--.-_-__-_-.__--_--__-_---_-_--_-.--__ 233

Alsterberg (azide) method (D:28a-l)________________________ 233

pH value (D:29)._ -..____ _. . 237

Instrument method (D:29a-l)________-_.___.-_-___.____--_- 237

Field instrument method (D:29a-2)____--..__.._.____.__-_._ 238

Paper method (D:29a-3)_.____._.________-________.--__ 238

Phenolic material (D:30)._____._____________--_-----___--___-__ 239

Gibbs reaction (D:30a-l)_____.____________________________ 239

Phosphorus (D:31)____________________________________________ 245

Orthophosphates (D:31a)_____-_-_---------------_--------- 245

Phosphomolybdate method (D:31a-l)___.__--____.______ 245

Phosphorus (D:31b)_____.____.________________. 247

Phosphomolybdate method (D:31b-l)______._._._.____._ 247

Potassium (D:32)_____________________________________________ 251

Flame-photometric method (D:32a-l)___.______-__.__-__---_ 251

Selenium (D:33)_--_-_____.___________________--_-_-______.--_ 253

Norris-Fay titration method (D:33a-l)______________________ 253

Silica (D:34)___..___..._________._..._______.... 259

Molybdate blue method (D:34a-l)___---------______________ 259

Gravimetric method (D:34a-2)_____________________________ 261

Sodium (D:35)__._..__..__...._..__._______._._._._.__.... 265

Flame-photometric method (D:35a-l)_______________________ 265

Zinc uranyl acetate method (D:35a-2)_______________________ 265

Solids (D:36)______________...______._'___...____._ 269

Dissolved solids (D:36a)___._._.________--_-._-_-__._.-.- 269

Residue-on-evaporation method (D:36a-l)________.__-_._ 270

Calculation method (D:36a-2)__----_--_---_.__.__.___ 271

Suspended solids (D:36b)______________________________ 273

Filtration method (D:36b-l)___________________ 273

Specific conductance (D:37)____________________________________ 275

Wheatstone bridge method (D:37a-l)__.________-______-.__ 275

Sulfate (D:38)________________________________________________ 279

Visual thorin method (D:38a-l)____--__-_____-_-_-____-_-_ 279

Spectrophotometric thorin method (D:38a-2)_________________ 281

Gravimetric method (D:38a-3)________-_-_-___,________---_ 283

Turbidimetric method (D:38a-4)___._____---___-_-_-____-___ 284

Sulfides (D:39)_____________________________________________ 287

lodometric method (D:39a-l)_____---__---_--___-______---_ 287

Turbidity (D:40)________________________________ 289

Hellige turbidimeter (D:40a-l)_____-_______-_______________ 289

Viscosity (D:41)___________________________________ 291

Brookfield viscometer (D:41a-l)___________-__-__-_-____-___ 291

Zinc (D:42)________________________________________ 293

Dithizone extraction method (D:42a-l)_-_-_-_______-_-__-___ 293

Direct color-comparison method (D:42a-2)___________________ 295

Index____________________________________________________________ 299
ILLUSTRATIONS

Page 

FIGURE 1. Depth-integrating samplers__-______-___-_----__________-_ 11

2. Point samplers._.____________-_______-----____-____--___ 12

3. Apparatus for sampling water that contains dissolved gases and

constituents susceptible to aeration______________________ 14

4. Sampler suspension apparatus_____________________________ 16

5. Tag for surface-water sample----____------------------_--- 22

6. Tag for ground-water sample____________________________ 26

7. Sample-bottle shipping cases.______________________________ 33

8. Sample-filtration assembly_____________-____-_-_-_________ 40

9. Relative-error curve for spectrophotometers without sensitivity adjustments_-_____-_____---____----__-------_-_- 47

10. Absorbancy curves for colored and turbid samples ___________ 51

11. Comparative energy transmission of filters._________________ 62

12. Forty-mm cell carriage for Beckman Model B spectrophotometer_______---________--______-_---_----__-___--_-_ 65

13. Schematic representation of photocells.____________________ 66

14. Photometric-titration assemblies.__________________________ 72

15. Arsine generator for the isolation of arsenic.________________ 102

16. Variation of the ratio of actual to apparent fluoride content

with aluminum content._______________________________ 167

17. Fluoride-distillation apparatus.___________________________ 170

TABLES

Page 

TABLE 1. Tolerances for volumetric glassware.--__-____--____-_--__--- 74

2. Factors for converting parts per million to equivalents per million. 80

3. Orthophosphate mixtures as related to pH_____________---_-_ 82

IX
METHODS FOR 

COLLECTION AND ANALYSIS OF WATER SAMPLES

By F. H. RAINWATER and L. L. THATCHER

ABSTRACT

This manual contains methods used by the U.S. Geological Survey to collect, 

preserve, and analyze water samples. Throughout, the emphasis is on obtaining 

analytical results that accurately describe the chemical composition of the 

water in situ. Among the topics discussed are selection of sampling sites, 

frequency of sampling, field equipment, preservatives and fixatives, analytical 

techniques of water analysis, and instruments. Seventy-seven laboratory and 

field procedures are given for determining fifty-three water properties.

INTRODUCTION

In January 1887, a committee on methods for stating water 

analyses reported to the Chemical Society of Washington: "Indeed, 

the researches of Clark, Frankland, Armstrong, Miller, Wanklyn, 

and others, within recent years, have brought the subject of wateranalysis into such a satisfactory form that little seems to be desired 

in addition, except some uniformity as to the statement of results." 

Nevertheless, the science of water-quality investigation has progressed during the intervening years. Spectrophotometry, flame 

photometry, the Schwarzenbach reactions, and other advances have 

revolutionized water analysis in recent decades. Further advances 

are on the horizon. Better methods for collecting, fixing, and preserving samples can be expected.

The purpose of this manual is to set forth methods of collection 

and analysis of water samples used by the Geological Survey in 

making water-quality investigations. An earlier manual, containing 

only analytical methods, was released in 1950 in preliminary form 

and was intended for use only in laboratories of the Survey. This 

manual has been expanded appreciably both in respect to scope and 

number of analytical methods. Although excellent and authoritative manuals on water analysis are available, most of them emphasize 

primarily either domestic, industrial, or agricultural water utilization. No single reference or combination meets all requirements 

as a guide to the broader phases of water-quality studies conducted 

by the Survey.
2 METHODS FOR COLLECTION AND ANALYSIS OF WATER SAMPLES

It is not practical to set forth procedures which can be followed 

with all the varied types of water and conditions throughout the 

United States. However, the procedures generally will be applicable 

to the study of most natural waters. In addition to the detailed 

instructions compiled here, based largely on experience of Survey 

chemists in the field of water chemistry, additional information can 

be found in the references given. This manual can serve both as a 

general reference for chemists and others engaged in water-resources 

investigations and as a field and laboratory guide for newly 

appointed personnel.

The development of new and improved methods is a continuing 

process. The selection and publication of the best methods in a 

manual of this kind are done with the knowledge that some of them 

may be superseded by better methods by the time the manual is 

printed. Nevertheless, it has been found essential to the work of 

the Geological Survey to provide a comprehensive manual of methods of analysis and pertinent related information for the guidance 

of technical personnel. It is expected that the manual will have to 

be revised in the next several years.

Several manuals describing the analysis of water samples have 

been prepared by technical associations, and a need has been recognized for the development of uniform methods by which these 

analyses can be done on a practical basis. For example, essentially 

the same method for the determination of chloride may be included 

in manuals of three different associations. Details of this method, 

however, such as strength of solutions and volumes taken for titration, may vary. It is believed that uniformity can be achieved 

without affecting the accuracy or application of the method. To 

this end the Joint Committee on Uniform Methods of Water Examination, composed of agencies in the United States that publish 

methods for water analysis, has been established to explore ways 

and means of effecting uniform procedures. The Geological Survey 

is identified with this effort.

Trade names used in the manual in connection with equipment, 

supplies, and reagents do not constitute an endorsement of the 

instrument or product. Knowledge of the operating principles of 

the various instruments is essential to performance of the duties of 

the analytical chemist. Trade names are used when necessary to 

describe the exact conditions under which statements pertinent to 

performance of the various operations in water sampling and 

analysis apply.
SECTION A. COLLECTION OF WATER SAMPLES

Water samples are collected and analyzed to ascertain characteristics of a body or mass of water. The sample is usually only an 

infinitesimal part of the total volume and is. therefore representative 

of the total mass only to the degree that uniformity of chemical 

composition exists within the total mass. In their natural state, 

surface and ground water are subjected to forces that promote mixing 

and homogeneity. The fact that such tendencies exist, however, is 

not sufficient cause for assuming that a body of water is so well 

mixed that no attention to sampling technique is required. Often, 

because of local conditions, the body of water may not have uniform 

composition.

The composition of water is subject also to change with the passage 

of time. The chemical quality of surface or ground water is the 

resultant of the geologic, hydrologic, biologic, and cultural environment of the water and varies from time to time as well as from 

place to place. Generally, changes in the quality of surface water 

are more pronounced and rapid than in ground water. However, 

marked changes in ground-water quality can, and often do, accompany such shifts in hydrologic equilibrium as variations in recharge 

or discharge rate, salt-water encroachment, or induced infiltration 

of surface water.

A:l SAMPLING-SITE SELECTION

The type of investigation, purpose of the study, and anticipated 

variation in chemical quality determine to a large degree the location 

of the surface- or ground-water sampling site and the frequency of 

sample collection.

A:la SELECTION OF SURFACE-WATER SAMPLING SITE

Stations for sampling dissolved material in surface water are 

usually operated by the Geological Survey to determine the discharge of dissolved constituents past a point, to describe the changes 

in water quality with respect to time, to collect data that will aid 

in predicting water quality in the future or in estimating the nature 

and magnitude of past events, to study the effect of geologic, hydrologic and cultural changes on water quality, or a combination of 

these purposes.

Adequate sampling of a flowing stream must take into account 

initially the need for defining the water quality in the stream cross

3
4 COLLECTION OF WATER SAMPLES Alia

section at the sampling site at the time of sampling, and secondly 

the need for a sufficient number of samples distributed in time to 

define the changes in quality of the water passing the sampling 

station. One sample can adequately define the water quality at a 

given time if the chemical quality is homogeneous throughout the 

cross section. Theoretically, if the water in the cross section is not 

uniform, a sample representing the average composition might 

be obtained by compositing depth-integrated samples of equal volume taken at several points simultaneously at centroids of equal 

flow, but, generally, on small and medium-size streams, and on some 

large streams, a sampling site can be found where the composition 

of the water is uniform. The problem of obtaining adequate samples 

is thus simplified, and the cross section can be represented by one 

sample taken at or near the center of flow. However, more than 

one sample may be required in the cross section of large streams.

Heterogeneous quality throughout the cross section can result from 

incomplete upstream mixing of ground-water inflow; tributary inflow; industrial, domestic, and agrarian effluents; or from tides and 

other backwater. Unmixed tributary inflow may be discernible for 

many miles down broad shallow streams; tidal stratification may be 

more perceptible in deep channels. The extent to which mixing 

occurs laterally and vertically in the cross section is governed principally by turbulence, which characterizes the movement of water 

in the stream. Differences in water velocity, both horizontally and 

vertically, aided by the configuration of the channel, also promote 

mixing. The effects of temperature and wind are of minor importance except in sluggish streams (especially tidal streams), lakes, 

and reservoirs.

Although differences in mineral content may be detected in the 

field by specific-conductance measurements or simple chemical tests, 

frequent laboratory analysis of samples for ascertaining uniformity 

is advisable.

The study of the chemical quality of tidal rivers is often beset 

with complexities not present in investigations of nonticlal streams. 

Keighton (1954) points out some of the natural forces and physical 

factors affecting the planning of water-quality investigations in 

tidal rivers and suggests sampling procedures that will determine 

adequately the variable quality of the water.

Easy access to the station, reliability of sample collectors, and 

availability of streamflow records are other items to be considered 

in selecting surface-water sampling sites. Locating a chemicalquality station at a stream-gaging station has decided advantages 

because adequate water-discharge data are essential for computing
A: la SAMPLING-SITE SELECTION 5

the dissolved load carried by the stream. However, the criteria for 

a good stream-gaging station and a water-quality station are not 

identical, and sometimes the most suitable location for one may be 

undesirable for the other. Rather than consider the presence of a 

stream-gaging station as a prerequisite for a chemical-quality station, 

the chemist may explore the possibility of obtaining water-discharge 

data at the chemical-quality station by direct or indirect means.

Each sampling station should be selected, so far as possible, to fit 

into a comprehensive network of sampling stations to cover a State 

or drainage basin. Even though an individual station is established 

to meet a specific need for information, the possibility of placing 

and operating it to supply data for other studies of chemical quality 

should not be overlooked. If two or more sites will give about 

equally usable information for the immediate needs, preference 

should be given to the one that most nearly fits the following 

criteria:

1. Best fits an overall plan for evaluating chemical quality on a broad basis.

2. Gives information that can be correlated with periodic information at other 

sites or that will supplement such periodic information.

3. Gives the total of dissolved material discharge from an area.

4. Is at a transition from the surface outcrop of one geologic formation to 

another.

5. Is at a place where future storage or diversion may be developed.

The selection of sites for sampling lakes and reservoirs depends 

on the degree of accuracy required in determining the composition 

and on the degree of mixing. Water circulation and movement 

caused by wind, temperature changes, and currents between points 

of inflow and outflow cause mixing. Other factors such as irregularity in shape of the body of water, difference in composition of 

inflow and stored water, evaporation, solution or evolution of gases 

at the air-water interface, and activity of aquatic plant and animal 

life all tend to bring about heterogeneous quality. A thorough 

study of water composition can be made by sampling along a three 

dimensional grid pattern; samples can be collected at different 

depths at each grid intersection. A more economical and conservative approach might be to sample at different depths along selected 

cross sections in the body of water. The adequacy of such a sampling program depends on the judicious selection of the cross section 

and sampling points. When only one sample is collected to define 

the average character of the lake or reservoir it should be collected 

near the center of the water mass. Conversely, in order to evaluate 

quality of reservoir water for the user downstream, the sampling 

site or station should be located at or near the point of discharge.
6 COLLECTION OF WATER SAMPLES A lib 

A:lb SELECTION OF GROUNI>-WATER SAMPLING SITE

Ground water is analogous to a surface-water reservoir in that 

most usable ground water is in motion, although the rate of movement may be very slow and the areal extent very wide. A well can 

be considered as a sampling point in a large body of slowly moving 

water, which differs in chemical composition vertically as well as 

areally. Most of the forces which cause mixing in bodies of surface 

water are absent or much weaker in ground-water reservoirs. Turbulence is virtually nonexistent. The major forces that tend to mix 

ground water are probably the differences in velocities as the water 

moves through material of heterogeneous permeability, pressure 

differentials and, to a lesser extent, ionic diffusion. The degree of 

movement induced by pumpage and natural discharge affects the 

quality. The diversified nature and solubility of the rocks with 

which the water comes in contact and variations in rate and chemical 

composition of recharge from precipitation and from the surrounding area tend to make the water heterogeneous.

Sampling programs are planned to determine the mineral content 

of ground water through the aquifer, although a completely comprehensive answer is not always practical. Efficient collection of 

water-quality data and intelligent selection of the ground-water 

sampling site generally require more judicious consideration than the 

selection of a surface-water sampling site because the elements influencing water quality are not as easily observed.

Because of the diversified purposes of ground-water investigations, 

it is impractical to attempt specific direction for the selection of 

sampling sites. Nevertheless, some general suggestions can be given. 

If changes in ground-water quality are not considered in the investigation, there are perhaps two equally satisfactory approaches to the 

problem of adequate and economical coverage of ground-water quality of an area; both employ comprehensive and partial analyses. 

One approach utilizes the determination of key constituents (one or 

several) in a large number of samples collected over the entire area. 

By this means an areal water-quality pattern is developed that is 

then of value in selecting the sites for collection of samples for 

comprehensive analysis. In some investigations the identity of the 

key constituents may be unknown at the beginning of the investigation. Then, the reverse approach may be required, and comprehensive analyses may be made early in the study, and these data 

augmented by partial and additional comprehensive analyses of 

samples collected at other sites to complete the water-quality picture. 

Either method requires the greatest of care in the selection of 

sampling points from available sites.
A:2 SAMPLING FREQUENCY 7

In selecting a sampling site to detect water-quality changes or 

stability, control of the variables that affect the change or stability 

should be considered. The hydrologic regimen during the investigation should not be modified. For example, if irrigation is begun 

in the vicinity of the well selected as a sampling site to study 

variations in quality resulting from natural-recharge patterns, the 

data collected may be essentially valueless.

Although some ground-water studies may be concerned only with 

surveying the chemical characteristics of the water, the data are 

often used in conjunction with other geologic and hydrologic information. Consequently, the value of the water analysis is usually 

directly proportional to the facts known about the source of the 

sample. One general observation is pertinent: The most useful 

samples are collected from wells for which good well schedules and 

other data are available.

A:2 SAMPLING FREQUENCY

Samples of surface and ground water should be collected at 

intervals such that no important change in quality could pass 

unnoticed between sampling times. This requisite immediately gives 

rise to two additional questions: What magnitude of change is 

important, and what are the physical and economic factors that 

must be considered in obtaining the record? By necessity the 

sampling schedule adopted is usually a compromise between the 

accuracy and detail desired in the water-quality record and the funds 

and personnel available for the work.

Streams are constantly, subjected to forces and changing environment which bring about variations in the chemical character of the 

waters. These variations may be relatively small in areas subjected 

to intensive weathering for long periods of time, such as the Piedmont province of the Atlantic coastal region. Younger, incompletely 

leached rocks and scattered, infrequent and often intense rainfall in 

the arid and semiarid Western States constitute an environment 

conducive to erratic and often extreme changes in water quality. 

Since the beginning of water-resources investigations in this country, 

the chemical-quality hydrologist has been confronted with the problem of the frequency of sampling required to detect adequately the 

changes in the chemical composition of a stream. On the basis of 

water-quality studies of a large number of streams and the practical 

economics of sampling-station operation, the U.S. Geological Survey has arbitrarily concluded that the collection of samples once 

daily is the minimum frequency necessary to define adequately the 

water quality for the majority of water uses, and daily sampling is
8 COLLECTION OF WATER SAMPLES Al3

considered as standard for the usual comprehensive chemical-quality 

investigation. However, less frequent sampling schedules are often 

used, and the value of analysis of weekly or monthly samples in 

reconnaissance or unit basin studies should not be minimized.

The chemical quality of ground water at the sampling point may 

vary in response to changes in rate of water movement, to pumpage, 

or to differences in rate and chemical composition of recharge from 

precipitation and from the surrounding area. Although concentrations of dissolved constituents in ground water from any one well 

may vary widely, sometimes severalfold, in general the changes are 

much slower than those commonly associated with surface water. 

Until disproved, it is safer to assume that the quality of water from 

a well fluctuates rather than that it is uniform for long periods of 

time. Changes in ground-water quality usually can be described 

satisfactorily by monthly, seasonal, or annual sampling schedules.

A:3 SAMPLING EQUIPMENT

The water chemist has a wide variety of sampling equipment 

from which to choose. Only those items that have been used satisfactorily by the Geological Survey are discussed here. This is not 

to imply that omitted items are unsatisfactory.

A:3a SAMPLE CONTAINERS

Factors that are pertinent in selecting containers used to collect 

and store water samples are resistance to solution and breakage, 

efficiency of closure, size, shape, availability, and cost. Preferences 

for one type of container over another are varied, and selection is 

guided largely by experience, supposition of the possible effect of 

the container on the water sample, and use of containers in the 

individual laboratories. No adequate study of all factors has been 

made.

Hard rubber, polyethylene and perhaps other plastics, and some 

types of borosilicate glass are believed to be suitable on the basis of 

experience within the Survey and the reports of others in water 

chemistry. A limited investigation, conducted by the Survey, of the 

relative merits of four common types of bottles showed that storage 

in Pyrex and polyethylene did not significantly alter the silica, 

sodium, total alkalinity, chloride, and boron content, or the specific 

conductance, pH, or hardness of the water during a storage period 

of about 5 months, although some investigators have avoided Pyrex 

because of suspected contamination from the boron in the glass. 

Glass bottles marketed under the name "No-Sol-Vit" added boron 

to the samples but were otherwise satisfactory. The increase in 

silica content of water stored in soft-glass bottles (citrate of mag-
A:3a SAMPLING EQUIPMENT 9

nesia type) was significant after 2 or 3 weeks. Prolonged storage 

in these bottles also increased decidedly the sodium content and 

hardness of the water. The chemical type and concentration of the 

stored water apparently has a definite influence on the rate and 

nature of materials dissolved. Much additional research is required 

before the findings will be conclusive. The American Society for 

Testing Materials (1954, p. 134) does not recommend soda-lime 

glass bottles as sample containers unless they are coated internally 

with paraffin.

Before use, all new bottles must be thoroughly cleansed, filled 

with water, and allowed to soak several days. The soaking removes 

much of the water-soluble material from the container surface.

The Geological Survey uses 4-ounce to 1-gallon bottles for sample 

collection. A sample of 1 or 2 liters is usually sufficient for the 

normal routine chemical analysis. Excluding containers for samples 

in which dissolved gases or constituents particularly affected by 

atmospheric gases are to be determined, the bottle design is immaterial. Nevertheless, uniformity within any one organization has 

definite operational advantages. The standard BOD (biochemicaloxygen demand) bottle with pointed stopper, which prevents entrapment of air bubbles in the bottle, should be used to collect 

samples that are to be analyzed for dissolved oxygen, carbon dioxide, 

cyanide, sulfides, and, probably, oxygen consumed. If the sample 

is supersaturated with carbon dioxide, a pressure-seal bottle is preferable.

Because of their design and construction material, some bottles 

are more satisfactory for transporting vnter than others. The 

bottle must be resistant to impact and to internal pressure, which is 

increased by expansion of the water or by release of dissolved gases 

at elevated temperatures. Well-sealed fragile bottles of liquid when 

shipped by air may not withstand the large differential pressures or 

freezing temperatures in the rarefied atmosphere. In respect to fragility alone, polyethylene bottles are more satisfactory than glass. 

However, samples in polyethylene bottles must be protected from 

compression, else the liquids may be squeezed out around the threads 

of the cap.

Samples subjected to freezing temperatures are very likely to be 

lost through breakage of glass bottles but are retained by polyethylene. However, this advantage may sometimes be more apparent 

than real. The chemical analysis of a previously frozen sample is 

always suspect because the original chemical character may not be 

completely reconstructed after the sample thaws. Although the 

analysis of the previously frozen sample is usually of some value,
10 COLLECTION OF WATER SAMPLES A:3b

the use of plastic bottles does not lessen the necessity for protecting 

samples from freezing.

The closures for sample containers should be glass, well-washed 

rubber, tin-wrapped cork, or inert plastics. Many of the common 

screw caps have waxed-paper or cork liners that are unsatisfactory. 

Corrodible metal should not be used. Pliable sheet plastic is often 

used to protect samples from contamination by direct contact with 

rubber stoppers, but this precaution is rarely necessary for the usual 

chemical constituents determined. No practical, completely gastight closure has been developed, but the efficiency of the seal is very 

important in the selection of caps and stoppers. The bottle seal 

must remain tight during transport to prevent spillage and gain or 

loss of gas. Some caps differ in their sealing capacity; shallow 

hard-plastic caps are notoriously susceptible to loosening during 

transport, but the deeper caps are somewhat better in that respect. 

Various types of stoppers have been used satisfactorily. The success 

with stoppers is largely dependent on the shape of the bottle neck 

and stopper and on the composition of the stopper. The flexibility 

of polyethylene permits a tight seal that withstands vibration. A 

practical, tight closure that the Survey has used is the patented 

stopper commonly associated with "citrate of magnesia" bottles. This 

stopper is permanently attached to the bottle neck and clamps firmly 

on a rubber washer, an advantage that may outweigh the disadvantage of solubility of the bottle when the time of contact between 

sample and bottle can be limited.

A:3b APPARATUS FOR SAMPLING WATER THAT CONTAINS NONVOLATILE CONSTITUENTS UNSUSCEPTIBLE TO AERATION

Depth-integrating samplers and point samplers are used in the 

collection of water for determination of nonvolatile constituents and 

those unsusceptible to aeration. The depth-integrating samplers 

consist only of a mechanism for holding and submerging the bottle. 

When the bottle is lowered at a uniform rate, water is admitted 

throughout the vertical profile. Hence the sampler must be sufficiently heavy to submerge the bottle in swift water. The method 

of securing the bottle to the sampler should require a minimum of 

manipulation. Belt-and-buckle devices are generally avoided because 

of the difficulty in handling such items in cold weather. Some of 

the depth-integrating samplers used by the Survey are shown in 

figure 1.

Point samplers are used to collect water at a specific depth below 

the water surface and may be very simple or complex in construction. The principal types are shown in figure 2. Sampler A is 

used only for collecting water samples from the bed of a stream or 

lake. The cork is tied to the suspension line in such a way that a
A:3b SAMPLING EQUIPMENT 11
12 COLLECTION OF WATER SAMPLES A:3b
A:3c SAMPLING EQUIPMENT 13

sharp pull on the line will remove the cork and allow the bottle to 

fill. In operation the bottle is corked and then thrown into the 

water with an excess of line; when the bottle comes to rest on the 

bed, the suspension line is jerked, removing the cork. The bottle is 

allowed to fill in place and then stoppered when it is withdrawn 

from the water. Sampler B is a ball-valve type and was designed 

for the collection of ground-water samples at any desired depth. 

When the desired depth is reached, the sampler is filled by raising 

and lowering several times in rapid succession, which opens the ball 

valves at either end and traps the water. The 1-gallon sampler is 

3 inches in diameter and the 1-quart size about iy2 inches. Sampler 

C, the Foerst sampler, is used to collect surface-water samples from 

a given depth and has been used with some success to sample water 

from wells. The sampler is lowered to the desired depth in the 

open position; then the messenger weight, which trips the closing 

mechanism, is run down the suspension line. The messenger weight 

must meet the triggering device on the top of the sampler squarely. 

Hence a set of messenger weights drilled to accommodate suspension 

lines of different diameters permits the use of the sampler with a 

wide variety of ropes and cables. The Foerst sampler is available 

in 4 sizes: 400-ml capacity, 2 inches in diameter; 1,200-ml capacity, 

'21/2 inches in diameter; 2,000-ml capacity, 3*4 inches in diameter; 

and 3,000-ml capacity, 3% inches in diameter. The samplers are 

constructed of brass and thus resist corrosion. Sampler D is used 

exclusively in sampling surface waters. It is lowered with stopper 

in place; the messenger weight then pulls the stopper and the bottle 

is raised to the surface open.

Detachable weights are useful with practically all samplers. Appreciable weight is necessary in excess of that required to submerge 

the bottle. In moving wrater, inadequately weighted samplers drift 

with the current and tend to ride on the water surface at the end 

of the suspension line. Added weight also decreases the angle 

between the suspension line and the vertical, thereby increasing the 

accuracy of the depth measurement.

A:3C APPARATUS FOR SAMPLING WATER THAT CONTAINS DISSOLVED GASES AND CONSTITUENTS SUSCEPTIBLE TO AERATION

The collection and handling of samples for the determination of 

dissolved-gas content and constituents susceptible to aeration require 

special equipment and careful technique. The sampler assembly for 

determining dissolved-oxygen content and biochemical-oxygen demand described by the Alnerican Public Health Association and 

others (1955, p. 250) or a modification thereof, is generally accepted 

as the standard apparatus for sampling open water. This sampler
14 COLLECTION OF WATER SAMPLES A:3c

provides for a threefold displacement of water in the sample bottle 

without aeration. Although this apparatus is efficient, satisfactory 

samples can be collected by other methods. A nonaerated sample 

collected in a Foerst sampler can be transferred by means of a tube 

connected to the outlet valve projecting into the bottom of the sample 

bottle. Twofold or threefold displacement of water in the bottle 

is recommended (see fig. 3).

FIGURE 3. Apparatus for sampling water that contains dissolved gases and constituents

susceptible to aeration.

Rather elaborate instructions are given by the American Society 

for Testing Materials (1954) for collecting samples from closed 

systems, and those procedures or modifications thereof can be used. 

Often a sample can be collected satisfactorily by running the water 

through a tube to the bottom of the bottle and using twofold or 

threefold displacement. The water should be run in slowly to minimize agitation and the resultant excessive aeration or loss of gas.

The samples are generally collected in narrow-mouthed biochemical-oxygen-demand bottles that have pointed glass stoppers to
A:4a SAMPLING INSTRUCTIONS 15

avoid entrapment of air in the sample. Unfortunately this type of 

stopper does not provide a seal that permits much transporting or 

handling of the sample unless the container in which the sample is 

transported exerts pressure on the stopper. "Citrate-of-magnesia" 

bottles (see sec. A: 3a) can be pressure sealed without entrapment 

of air bubbles and, consequently, are used frequently for transporting samples for the determination of dissolved gases and constituents 

susceptible to aeration.

A:3d SAMPLER SUSPENSION APPARATUS

Sashcord with or without a wire core is the most common means of 

suspending depth-integrating samplers (fig. 4). When long lengths 

of line are required, commercially available clothesline reels (fig. 4L4) 

are useful. These reels also have a locking device, which facilitates 

use with point samplers. Figure 4JS shows a versatile assembly 

with which the exact depth can be read from a depth indicator on 

the reel; different weights can be interchanged as required by the 

current. In use the crane is tipped forward against the bridge rail. 

Modified booms having similar elements can be used from boats.

A:4 SAMPLING INSTRUCTIONS 

A:4a SAFETY PRECAUTIONS

Water samples are collected under a wide range of conditions, 

and the work is not without some hazard. A knowledge of the 

hazards involved and means by which they can be minimized should 

be helpful in preventing accidents and in providing greater safety 

for sample collectors. Surface-water samples are collected from 

bridges, docks, cableways, and boats, and by wading, or through 

ice. Inability to swim and difficulty in freeing oneself of burdensome equipment if suddenly plunged into deeper water are perhaps 

the most serious hazards in taking samples from streams, lakes, 

and reservoirs. The collection of ground-water samples has its 

own perils, primarily in connection with pumping equipment, snakebite, and unfriendly dogs. Dangers include loose or rotten well 

coverings and entanglement with belts and drive shafts. Many wells 

are surrounded by heavy brush which makes an attractive habitat 

for snakes, particularly during dry, hot months. High boots will 

reduce the likelihood of snakebites considerably, as most bites are 

below the knee. Beating the grass and proceeding cautiously will 

give the snakes an opportunity to glide away instead of striking.

Wading. Wading is one of the easiest methods to collect samples 

from many streams and also affords the collector the best opportunity to "size up" the flow and decide where to collect the samples. 

Rubber boots or breast-high waders are standard equipment. A
16 COLLECTION OF WATER SAMPLES A:4a

FIGURE 4. Sampler suspension apparatus. A., Hand assembly for shallow water ; B, crane

assembly for deep water.

wading rod or similar probing instrument is essential to safe wading. 

By probing ahead, the collector can estimate the current and locate 

holes, benches, and even quicksand. In shallow7 water the most 

satisfactory method of extricating oneself from quicksand is to 

fall flat on the bed and craw7! to firmer ground; in deep water 

one's effective weight in the sand can be lessened by swimming
A:4a SAMPLING INSTRUCTIONS 17

motions. In a swift sandy-bottom stream or where the bottom is 

covered with ooze, the wader may find that his feet slowly sink into 

the bottom if he stands in one position too long. Lake or pond 

bottoms may be more treacherous than the beds of flowing streams.

A general rule of thumb is that wading should not be attempted 

if the depth of the water in feet multiplied by the water velocity 

in feet per second equals 10 or more, but this criterion must be 

modified by many factors peculiar to the site and to the season of 

the year. The depth as related to the velocity of water that can 

be waded safely is closely associated with the body weight. If 

the wader has any uncertainty about his ability to wade a stream, 

he should attach a rope or tag line securely to the bank. An 

extra change of clothing is advisable if much wading is to be done, 

particularly during the colder months.

Bridges. Traffic is the most serious hazard when working from 

bridges. Sometimes the bridges have walkways for pedestrian 

traffic or catwalks suspended at the side or beneath the bridge, 

but more often than not the collector must work in the traffic lanes. 

If necessary to interfere with traffic, suitable arrangements with 

local authorities should be made in advance. Warning signals and 

signs of various kinds are some protection when sampling from 

a bridge, and approval for their use is generally granted by local 

traffic officers. Commonly used signals include flags, signs, red 

lanterns, and blinking red lights. For some reason motorists respect certain signals more than others. For night work a blinking 

red light is usually more effective than other warning signals. 

Elevating the light several feet above the top of the car is effective 

in congested traffic; this affords some warning to drivers whose 

vision might be obstructed by preceding vehicles.

On highway bridges that have particularly heavy traffic, and 

when cranes are used to handle sampling equipment, 1 lane of 

traffic may be blocked off by placing a barricade and suitable 

lights or "1-lane-traffic" signs on both bridge approaches. For 

additional protection, the car or truck may be parked in one lane 

and between the collector and oncoming traffic. After all the 

prescribed traffic warnings are installed the collector still must 

keep an eye on approaching vehicles there are always drivers 

unwilling to cooperate.

When working on railroad bridges a knowledge of train schedules 

is essential, and at no time should the collector use equipment that 

cannot be removed quickly.

A hazardous condition in which the sampler is likely to inflict 

serious injury on others exists at bridges over navigable streams.
18 COLLECTION OF WATER SAMPLES A:4a

Boat operators cannot see the small suspension line until they are 

almost ready to collide with it, and the sample collector frequently 

cannot see a small boat approaching. Red pennants attached to the 

suspension line or to an independent line in the same location are 

recommended.

High-tension powerlines strung close to the bridge are also dangerous. If it is necessary to sample repeatedly from a bridge that 

has hazardous powerlines, a warning sign should be painted on 

some part of the bridge structure directly above the hazard.

Cableways. The principal hazards of cableways are not the 

possibility of their structural failure, but rather those that attend 

the operation of the cable car. Proper instruction of inexperienced 

personnel in cable-car operation may prevent many accidents. Before releasing a car from its mooring, check the car pullers; in 

some districts the puller is left in the car and in others it may 

be carried as a part of field equipment. Also check the brake lining 

attached to the car puller; this lining is used to scotch the sheaves 

and hold the car stationary while the sample is collected. Once 

the car is rolling free on the cable, no attempt should be made 

to retard its speed by use of hands or car puller until it has slowed 

almost to a standstill. Some of the later model cable cars are 

equipped with brakes. If there is any danger of the car continuing across the stream and crashing into the cable mooring on 

the opposite bank, the sample collector should ease the car out 

from the landing platform with the oar puller or brake and never 

allow the car to run free.

Boats. Water samples are often collected from boats. The degree of hazard involved is generally dependent on the selection 

and operation of the boat. Lifejackets are essential equipment.

Much of the hazard in sampling large streams, lakes, and reservoirs is related to the seaworthiness of the craft and to changing 

weather. A sampling trip in a small boat begun on mirror-calm 

water can suddenly develop into a perilous expedition if the wind 

comes up and large waves form. Only an experienced man can 

judge the roughness in the center of the reservoir from the breeze 

on the shore. In many parts of the country the wind often 

follows a rather definite pattern during any 24-hour period, depending on the season of the year. For example: During the summer and early fall in the Great Plains area, the water is generally 

calm in the early morning, the wind rises about midmorning, reaches 

its peak in the afternoon, and dies shortly before sunset. Other 

regions have their own wind patterns. It is always advisable to
A:4a SAMPLING INSTRUCTIONS 19

check the weather forecast before planning an extended samplecollection trip on a large lake or reservoir. All work should be 

done at a safe distance from spillways, "glory holes," and other 

areas of water discharge.

Boat operations on streams during floods is particularly hazardous, and, where measurements by boat are required, the services and 

equipment of skilled local personnel should be utilized where possible. Not only must the boatman cope with the fast current, but 

he must be on the alert for shallows and for floating and submerged 

drift. Floating ice is very dangerous; even when no danger to 

the boat is involved, changes in wind, current, or tide may shift 

the position of ice floes in such a way as to cut the boat off from 

the shore. On navigable waters in fog it may be necessary to stop 

the motor at intervals to listen for other vessels.

The good boatman checks such items as oars, life preservers, 

buoyancy tanks, lights, motor performance, gasoline supply, and 

spare parts (particularly shear pins) and inspects the boat for 

leaks before he leaves the dock. He loads the equipment so as to 

provide the most efficient operation of the boat and installs his 

booms where the danger of capsizing or swamping the boat will be 

at a minimum.

Ice. Samples are sometimes collected through ice at considerable 

risk. A hole of sufficient diameter to collect a sample from a 

bridge may be broken in thin ice by dropping a weight attached 

to a handline. But if the ice is thick, or no bridge is present, 

it may be necessary to venture out on the ice. The ice in a stream 

is likely to be of variable thickness, and the strength of the ice 

cannot be estimated from the apparent thickness near the edges. 

In advancing across an ice-covered stream it is advisable to test 

the ice with a sharp blow of an ice chisel every few steps. A 

few inches of new snow on top of ice can conceal dangerous places 

that would be visible if there was no snow. Work from ice during 

breakups is particularly hazardous. Moving blocks of ice are as 

destructive to sampling equipment as other debris. Vehicles should 

not be driven over ice-covered bodies of water.

First aid. By definition, first aid is the immediate and temporary 

care of an injured or suddenly ill person until a doctor can be 

obtained; it is not a substitute for competent medical attention. 

Each field party should be supplied with a standard first-aid kit 

and first-aid manual. "First Aid Guide," published in 1954 by 

the U.S. Department of Agriculture, Forest Service, and sold by 

the U.S. Government Printing Office is a satisfactory manual.
20 COLLECTION OF WATER SAMPLES A:4b 

A:4b COLLECTION OF REPRESENTATIVE SAMPLES

After the sampling site and frequency of sampling have been 

established, certain precautions must be exercised to insure that the 

samples collected are representative.

If only one sample of surface water is collected in the cross 

section of the stream it should be collected near the center of flow. 

The center of flow may or may not correspond with the midpoint 

of the stream. The single sample should also be depth integrated 

by lowering and raising the bottle vertically at a uniform rate 

from water surface to bottom, taking care not to get excessive 

amounts of bed material in the sample. Samples should not be 

collected from ponded or sluggish water or coves near the bank. 

Many water gages are located in pools upstream from natural or 

manmade low obstructions to flow, called controls. If there is no 

flow over the control this fact should be noted on the sample 

bottle.

A representative sample can be collected through a hole in the 

ice only with difficulty. The hole will often fill with slush or 

chips that rarely have the same dissolved-mineral characteristics 

of the water beneath the ice. Ice should be excluded from the water 

sample either by clearing the hole or by collecting the sample 

below the ice. Information on means and conditions of collection 

should be noted on the sample bottle.

When a ground-water sample is collected to determine the chemical 

characteristics of the water in the aquifer adjacent to the well, the 

water in the gravel pack, well casing, and distribution system between the aquifer and the point of collection should be displaced 

several-fold before the sample is collected. Water-quality data serve 

many other uses in ground-water studies, and the collector's techniques for special purposes may differ appreciably from the standard 

procedure. The important point is that the collector must know 

the volumes involved in the water system and have a reasonable 

idea of what the sample represents.

Precautions in the collection of special samples are discussed in 

section A: 4d.

All bottles should be rinsed with the water to be sampled before 

the sample for analysis is collected. Sufficient air space should 

be left in glass bottles to allow for expansion of the water at increased temperatures; polyethylene bottles may be completely filled.

A:4o TABULATION OF SAMPLE DATA

The importance of describing in writing the source of the sample 

and conditions under which it was collected cannot be overem-
A:4c SAMPLING INSTRUCTIONS 21

phasized. A water analysis is of limited value if unaccompanied 

by detailed information on what the sample represents. Field notes 

are extremely valuable in project-type investigations of water quality, 

but they can easily be misplaced or lost. Field notes never should 

take the place of detailed information accompanying the sample 

from point of collection to finished analysis tabulation.

The Geological Survey records the sample description in two 

ways. Small soft-glass bottles, of 12-ounce capacity and less, are 

usually etched on the outside, and the information is written directly 

on the etched surface with a lead pencil. Standard sample tags, 

which provide space for sufficient detail to describe most samples, 

are used for other types of containers (fig. 5 and 6). The tag for 

ground-water samples is numbered 9-016 and that for surface-water 

samples is numbered 9-070.

Surface water. The minimum data required for most surfacewater samples are:

Name of the water body 

Location of the station or site 

Point of collection 

Date of collection 

Time of collection 

Gage height or water discharge 

Temperature of the water 

Name of collector

Weather and other natural or manmade conditions that may assist 

in interpreting water quality

Many surface-water samples are collected at stream-gaging stations operated by the Geological Survey. The exact name used 

by the Survey for the stream-gaging station must be recorded. One 

common and extremely troublesome example of carelessness is the 

use of "at," "near," "above," or "below" in disagreement with the 

official name of the station. If the sampling station is at or within 

1 mile of a surface-water gaging station, the name of the water 

body and location can be noted simply; for example, "Colorado 

River near Glenwood Springs, Colo., at gage" or "Colorado River 

near Glenwood Springs, Colo., 100 yards above gage." If any 

tributaries or other influents enter the stream between the sampling 

point and the gage, this fact should be recorded under "Remarks." 

(See fig. 5.)

When a station is established for the collection of surface-water 

samples on a routine basis where water-discharge data are not 

compiled, this station is given an official name, in a manner similar 

to the designation of a stream-gaging station. Where the sampling 

site is more than 1 mile from the gage or routine sampling station,
22 COLLECTION OF WATER SAMPLES A:4c

WATER SAMPLE

U. S. Geological Survey 

Location

(City, at, near, or direction

(State)"

(Name afatreani)"' °"" """"~"~ ~"

Point of collection

(OVER) lft-47071-l OPO"

Collected by

Date of collection

Gage height (ft.) __________ Time 

Appearance

Remarks

LEAVE A LITTLE AIR SPACE IN BOTTLE

FIGURE 5. Tag for surface-water sample.

the information given for the location of a sHe should be complete 

and clear enough for a person to locate readily the exact sampling 

site on suitable maps and to return to the location. For example, 

if the following sites were not at regular stations they would be 

described as "Republican River at Cambridge, Nebr., at bridge 

on State Highway 47, y2 mile south of Cambridge, Furnas County, 

14 mile upstream from Medicine Creek" or, if the site is not at 

a bridge, "Supply Canal (Tri-County Diversion) near Maxwell, 

Nebr., at Parshall flume, sec. 28, T. 13 N., 29 W., Lincoln County." 

If topographic maps are available and if there are no distinguishing 

landmarks (such as a Parshall flume), the location of the site should
A:4c SAMPLING INSTRUCTIONS 23

be pinpointed to within a quarter-quarter section. Sketches of the 

sampling station in respect to roads, buildings, and the surrounding 

area are often desirable.

The quality of surface water may be far from uniform throughout the water body. For this reason the location of the sampling 

site must be clearly defined.

"Point of collection" refers to the spot within the cross-sectional 

plane of the stream where the sample is collected. When only 

one sample is collected from a stream it is usually obtained from 

the center of flow, and the sample should be so marked. If more 

than one sample is collected, all should be designated by station 

numbers. Station numbers, unless otherwise indicated, are the distance, in feet, from a predetermined reference point on one of the 

banks. Many station numbers are painted on bridge railings, but 

some are determined by a tagline or other measure stretched across 

the stream. A notation accompanying the sample should tell where 

the sample was collected in relation to the banks. The abbreviations "LEW" and "KEW" for the left and right edges of the water, 

respectively, are used to code this information. For example, the 

coded notation "Sta. 25', LEW 5', REW 65'" tells that the stream 

was 60 feet wide at the time of sampling and that the sample 

was collected 20 feet from the left edge of the water. The left edge 

is on the left side as one faces downstream.

A series of stations inr an unchanging direction on a stream, lake, 

or reservoir constitute a range line (RL). Range lines are usually 

located in respect to bench marks or readily apparent topographic 

or cultural features. In establishing a range line, consideration 

should be given to the permanence of the landmarks. It is very 

disconcerting to establish a range line in relation to a red barn 

and a tall tree and then return to the area later and find that 

someone has repainted or torn down the barn and cut down the 

tree. Stations along a range line should be noted as directed in 

the preceding paragraph wherever possible. A transit mounted on 

shore and two-way communication can be used to keep the boat 

on the range line. Orientation along a range line on a lake or 

reservoir may be more difficult than on a stream. Long taglines 

can be used but they are not practical when the distance is long 

or when they are a hazard to navigation. The position of the 

boat can be located accurately by triangulation, but this method 

requires two or more transits on shore, a shore party, and means 

of communication. Another method is to lay out 2 markers a 

known distance apart on shore and read the marker-boat-marker 

angle with a sextant; angles greater than 30° and less thanJ30°
24 COLLECTION OF WATER SAMPLES A:4c

are more accurate for plotting. Plotting of this information is 

simplified when one marker is on the range line. Less precise 

methods include time of travel at a uniform rate, stadia distances, 

and visual estimation. On estimating distances by sight it is often 

easier to record distances in reference to the nearest shore than in 

reference to the point of origin. If a bottom profile is available 

along the range line, or if soundings have been made on a previous 

trip, the sampling station can sometimes be located with sufficient 

accuracy by soundings; of course, depths must be corrected for 

changes in water-surface elevation. This identical-depth technique 

is limited to bodies of water that have had little change in bottom 

profile during the period since the last sampling. These are only 

a few of the methods available for station location in lake and 

reservoir surveys.

If more than 1 sample is collected in the vertical profile at 

1 station, the depth at which the sample is collected below the 

water surface, as well as the total depth of water, should be recorded. Therefore, to locate adequately the point of collection the 

coded notation accompanying the sample may read "RL 2, LEW 

10', Sta. 125' REW 500', depth 4' in 53' of water."

"Time of collection" should include a.m., p.m., 12:00 m (noon) 

or 12:00 p.m. (midnight), if military time is not used. The use 

of daylight saving time is always a cause of uncertainty and confusion. Consequently, notations of standard time on the samples 

are desirable. Any recorded daylight saving time must be clearly 

marked as such.

"Gage height" (water level) should always be recorded if a waterdischarge measuring station is nearby. The stream discharge can 

be computed from this gage height and suitable rating tables or 

curves. At sites other than water-discharge measuring stations it 

may be necessary to compute an approximate discharge from data 

at other stations in the general area or to make a measurement or 

estimate of discharge. Estimated, measured, or computed discharge 

entered in the space on the tag provided for gage height should 

be clearly noted as such, in order that it will not be confused with 

verified, factual information. When a measurement is made, the 

discharge is recorded on the sample bottle or tag in the field and 

noted "unchecked." A manual of methods and practices used by 

the Geological Survey to gage streams has been prepared by Corbett 

and others (1943). Water-surface elevation of impounded water 

at time of sampling should be recorded.

Descriptions of turbidity and color of sediment and water are 

particularly helpful if there is any question of chemical or phys-
A:4c SAMPLING INSTRUCTIONS 25

ical change between the time the sample was collected and when 

it was received in the laboratory. Other information pertinent to 

the physical characteristics of the sample at the time of collection 

or to conditions under which the sample was collected can be noted 

as "Kemarks." If the water has been treated, the treatment should 

also be described here, if possible.

Ground water. Ground-water sampling sites are somewhat more 

difficult to locate and to identify than surface-water sites, and 

the use of ground-water analyses requires more information. In 

quality-of-water work, springs and seeps are considered as ground 

water. Minimum data required for most ground-water samples 

include:

Geographical and legal location 

Depth of well 

Diameter of well

Length of casing and position of screens 

Method of collection (source) 

Point of collection 

Water-bearing formation (s) 

Water level

Yield of well in normal operation 

Water temperature 

Principal use of the water 

Name of collector 

Date of collection 

Appearance at time of collection

Weather or other natural or manmade factors that may assist in interpreting chemical quality

When recording the data on the sample-bottle tag (fig. 6), the 

collector should differentiate between those that he has determined 

himself and those that are reported by the owner or well driller.

"Geographical and legal location" ("location" on fig. 6) refers 

to the nearest town and the direction and distance therefrom; 

county and State should also be included. Geographical locations 

give the general area only and are of little help in locating one 

well in the presence of many.

Legal descriptions of the sampling site and well-numbering 

systems used by the Geological Survey are not standardized 

throughout the United States. In areas surveyed by the General 

Land Office the most logical designations, and those used in 

many areas, are related to townships north and south of base lines, 

ranges east and west of principal meridians, sections, and quarterquarter sections. The standard tag for ground-water samples, No. 

9-016, includes space for this type of location. It also has a space 

designated as "Field No." where other location descriptions can be
26 COLLECTION OF WATER SAMPLES A:4c

shown. Multiple wells on farms and ranches are not uncommon. 

Sketches of the location of the sampled well to fence lines, buildings, and topography are very helpful. Some owners do not object 

to numbering the wells with a permanent mark on the pump or 

structure.

"Method of collection," designated by the ambiguous term 

"Source" on standard tags now in use, refers to the description 6f 

the well and equipment relating to means of taking the sample, 

such as motor-pumped well, windmill, hand-pumped well, flowing 

well, spring, seep, and sample obtained with a bailer.

"Diameter," other than for uncased wells, refers to the diameter 

of the submerged casing and not to the diameter of the gravel

WATER SAMPLE

U. S. GEOLOGICAL SURVEY

Field No.

Location

(City, at, near or direction from)

Source

(Type of well)

Depth (ft.) Dlam. (in.)

from to

(OVER)

W

Date 

Cased to ft. drilled

Point of collection

Owner

(Address of owner)

Water-bearing formation

Water level ft. above 

below

Yield G. M. Flow

Temperature

Pump « " <"« >

Degrees F.

Circle use

Use: Dom., Ps., Stock, Irr., Ind., RR., 

Air Cond., Bottling, Condensing

Collected by

Date of Collection

Appearance

(Clear, colored, turbid, sediment, etc.)

Remarks

LEAVE A LITTLE AIR SPACE IN BOTTLE

FIGURE 6. Tag for ground-water sample.
A:4d SAMPLING INSTRUCTIONS 27

pack or hole drilled when the well was put in. However, the 

diameter of the gravel pack may be included as additional information. Diameters and lengths of telescoped casings should be given 

in order from top to bottom.

"Point of collection" concerns the location in the water distribution system where the sample was collected. Inasmuch as ground 

water can change appreciably in chemical quality as it passes 

through the distribution system, any points of collection other than 

at the pump should be described in detail under "Remarks." A 

sketch of the distribution system is helpful.

"Water-bearing formation" refers to the principal aquifer or 

aquifers that discharge the water. The geologic name of the aquifers, not just sandstone, limestone, sand, and gravel, should be 

recorded if definitely known. If a certain aquifer is believed to be 

the source of the water but some uncertainty remains, the aquifer 

name is listed and followed by a question mark. The question 

mark must be retained on all tabulations until the correctness of 

the aquifer name is substantiated. If further investigation disproves the original assumption, the original notation should be 

changed.

Water levels are commonly measured to the nearest hundredth 

of a foot from a convenient measuring point on or near the well 

structure; the same point is always used and is described in the 

well schedule. Because the measuring point may be above or 

below the land surface, it should be clearly stated on the tag 

whether the recorded water level is in relation to the land surface 

or a measuring point. When a well is pumped, the water level 

may be substantially different from the "static" or nonpumping 

level. Therefore, the reported information should indicate whether 

the water level represents pumping or nonpumping conditions. It 

may be advisable to indicate the length of time the pump has been 

operating and the discharge rate.

The turbidity, color, and sediment content of the sample at the 

time of collection should be recorded under "Appearance." Rusty 

water suggests that the casing or distribution system is the source 

of the iron. A sample containing an appreciable amount of iron 

or manganese in solution may be colorless and clear at the time 

of collection, but these metals may oxidize and precipitate after 

the sample is collected.

A:4d SPECIAL TREATMENTS, PRESERVATIVES, AND FIXATION

Many of the heavy-metal ions normally present in only trace 

amounts in natural waters may not remain in the water sample 

until it is analyzed because of such chemical and physical reactions
28 COLLECTION OF WATER SAMPLES Al4d

as oxidation, reduction, precipitation, adsorption, and ion exchange. 

Coprecipitation may also reduce the concentrations of these constituents in solution. Many commonly used methods for the separation and isolation of traces of substances involve Coprecipitation 

and adsorption processes. Iron is a particularly troublesome component in natural water because of its tendency to precipitate from 

solution and to coprecipitate other metal ions.

Although most water chemists recognize the chemical and physical 

forces that remove metals from solution, the literature provides 

very little specific information on preferable techniques for collection, storage, and preservation of samples for the determination of 

hydrolyzable metals. Most authoritative manuals recommend a 

minimum of lapsed time between sample collection and analysis. 

This is a valid recommendation but one which laboratories of the 

Geological Survey can seldom follow because of the distance between laboratories and many sampling sites. Measures other than 

speed must be taken. The Geological Survey has not conducted 

extensive research on the problem, but an analysis probably will 

be more representative of the water at the time of collection if 

a separate sample of water is taken for the determination of aluminum, chromium, copper, iron, lead, manganese, and zinc. The 

sample should be freed of its sediment and acidified to about pH 

3.5 with glacial acetic acid (0.5 ml per 100 ml maximum) at the 

time of collection. Inasmuch as acetic acid stimulates the growth 

of molds, it may be necessary to add a little formaldehyde (0.2 ml 

per 100 ml) to retard growth. The presence of the acid also tends 

to stabilize the original valence state of the metals. Although 

not an ideal preservative, this acidification minimizes precipitation as well as adsorption on the container. Even collection of a 

separate clear sample, and, later, drastic acidification in the laboratory, is preferable to determination of the trace metals in alkaline 

water that has been stored for some time.

The determination of other constituents requires specific fixatives 

or preservatives. These techniques, without exception, merely retard the natural processes of change in composition, and the sample 

should be analyzed as soon after collection as possible. Recommended treatments are described below.

Ammonia nitrogen. The percentage concentrations of the various 

components of the nitrogen cycle are rapidly changed by the biologic activity of organisms common in natural waters. Chloroform 

(5 ml per liter of sample) inactivates the organisms and fixes the 

nitrogen-cycle components.
A:4d SAMPLING INSTRUCTIONS 29

Chlorine residual. Samples to be analyzed for chlorine residual 

should not be stored, as the chlorine content of a water sample decreases rapidly. Exposure to sunlight or other strong light, and 

agitation, will further reduce the amount of chlorine present.

Dissolved oxygen. Immediately after collection, the dissolved 

oxygen in a sample should be converted to an equivalent amount 

of free iodine, which is more stable in solution than oxygen. The 

sample is normally collected in a 300-ml BOD bottle, with a minimum of aeration (see sec. A:3c). Water temperature must be 

taken at the time of collection if the percentage of saturation is 

to be computed.

Suspended solids in high concentration may adsorb appreciable 

quantities of iodine in acid solution and they should be removed 

by coagulation with aluminum sulfate (alum) as follows: Collect 

the sample in a 500-ml glass-stoppered bottle using the same precautions as for the BCD bottles; add 10 ml of 10 percent alum 

solution and 1-2 ml of concentrated ammonium hydroxide; mix 

by inversion and allow the floe and suspended material to settle; 

decant the supernatant into a BOD bottle until it overflows.

If the sample is clear, or after it has been clarified, proceed with 

the following treatment.

1. Add 1 ml KF solution below tiv liquid surface of the sample in a BOD bottle.

2. Add 2 ml MnSCU solution below the liquid surface.

3. Add 2 ml alkaline-iodide sodium azide reagent below the liquid surface.

4. Stopper and mix by inversion. Allow the precipitate to settle and then repeat 

the mixing and settling processes.

5. Add 2 ml concentrated ILSO4 by allowing the acid to run down the neck 

of the bottle.

6. Mix by gentle inversion until solution is complete.

In the absence of interferences described in section D:28a-l, titration of the released iodine can be postponed for 48-72 hours if 

the treated sample is refrigerutsd. Instructions for preparation of 

reagents are given in section D : 28a-l.

Nitrate nitrogen. Although routine nitrate determinations are 

usually made on an untreated sample at the time of the regular 

analysis, treatment with chloroform is recommended if the exact 

quantity of nitrate in the sample at the time of collection is required (see "Ammonia nitrogen").

Nitrite nitrogen. See "Ammonia nitrogen. 1"

Oils and waxes. Samples for the determination of oils and waxes 

should be collected in a separate bottle and fixed with 2 ml of 

concentrated HC1 per liter to inhibit bacterial action, which may
30' COLLECTION OF WATER SAMPLES A:4d

decompose oils and waxes. Care should be taken that the sample 

is representative. It is very difficult to obtain a representative 

sample from surface water because the material is seldom, if ever, 

uniformly distributed.

Organic nitrogen. See "Ammonia nitrogen."

Orthophosphate. The conversion of organic phosphorus to 

orthophosphate is largely the result of microbiologic action. Chloroform (5 ml per liter of sample) inactivates the organisms and stabilizes the organic phosphorus-orthophosphate system.

Phenolic material. It has been shown that phenolic materials in 

solution are rapidly destroyed by the biologic activity of organisms 

common in natural surface waters. Copper sulfate (2 g per liter 

of sample) has been found to be an effective preservative for a 

few days. Refrigeration can be used as an additional measure.

Selenium. Very little information is available on the stability 

of selenium in aqueous solutions. Selenium is much more prevalent in bottom deposits of undrained lakes and the sea than it is 

in water above these deposits. This fact leads to the conclusion 

that selenium tends to precipitate from solution, possibly by reaction 

with iron salts, or to be adsorbed by some stream sediments (Trelease 

and Beath, 1949). If appreciable delay in analysis is anticipated 

the sample should be filtered at the time of collection and acidified.

Sulfide. Water containing dissolved sulfides readily loses hydrogen sulfide, particularly if the pH of the sample is low. Oxygen 

destroys sulfides by oxidation, particularly if the pH of the sample 

is high. Aeration of the sample should therefore be avoided (see 

sec. A: 3c). Two grams of zinc acetate per liter of water will fix 

the sample for several days. Acid water must be neutralized before 

addition of the zinc acetate.

REFERENCES

American Public Health Association and others, 1955, Standard methods for the

examination of water, sewage, and industrial wastes: New York, Am. Public

Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

Corbett, D. M., and others, 1943, Stream-gaging procedure: U.S. Geol. Survey

Water-Supply Paper 888. 

Keighton, W. B., 1954, The investigation of chemical quality of water in tidal

rivers: U.S. Geological Survey open-file report. 

Trelease, S. F., and Beath, O. A., 1949, Selenium; its geological occurrence

and its biological effects in relation to botany, chemistry, agriculture,

nutrition, and medicine: New York. [Published by the authors, Box 42,

Schermerhorn Hall, Columbia Univ.]
SECTION B. HANDLING OF WATER SAMPLES BEFORE

ANALYSIS

The Geological Survey is confronted with problems of handling 

and storing water samples that do not prevail in most public and 

private laboratories. There is usually only one laboratory within 

a State of group of several States. In each area the efficient 

handling of samples should be considered both in respect to the 

validity of the analytical results and the cost of transportation 

and storage.

B:l TIME BETWEEN SAMPLE COLLECTION AND ANALYSIS

The effect of bottle composition in changing the quality of water 

samples during storage is discussed in section A: 3a. The samples 

are also subject to changes not attributable to bottle composition. 

The general relations of the time between collection and analysis 

to the validity of the analytical results as being representative of 

the water quality at the time of collection are well presented by 

the American Public Health Association and others (1955) as 

follows:

In general, the shorter the time that elapses between collection of a sample 

and its analysis, the more reliable will be the analytical results. For certain 

constituents and physical values, immediate analysis in the field is required 

in order to obtain dependable results, because the composition of the sample 

may change before the sample arrives at the laboratory.

Some determinations are more likely to be affected than others by storage 

of samples prior to analysis. Certain cations are subject to loss by adsorption 

or ion exchange upon the walls of glass containers. These include iron, copper, 

aluminum, manganese, trivalent chromium, and zinc. Temperature changes 

very quickly; pH may change significantly in a matter of minutes; dissolved 

gases may be lost (O2 , CO2 , H2S, C12 , CEU) or gained (O2, CO2 ). Determinations 

of temperature, pH, and dissolved gases should always be carried out in 

the field, because changes are almost inevitable by the time the samples 

reach the laboratory. With changes in the pH-alkalinity-carbon dioxide balance, calcium carbonate may precipitate and cause a decrease in the values 

for calcium and for total hardness. Iron and manganese form readily soluble 

compounds in their lower (reduced) valence states, and relatively insoluble 

compounds in their higher (oxidized) valence states; therefore these cations 

may precipitate out, or may dissolve out of a sediment, depending upon the 

redox potential of the sample. Microbiological activity may be responsible for 

changes in the nitrate-nitrite-ammonia balance, for decreases in phenols and 

in BOD, or for the reduction of sulfate to sulfide. Any residual chlorine is 

converted to chloride. Sulfide, sulfite, ferrous iron, iodide, and cyanide may

31
32 HANDLING OF SAMPLES BEFORE ANALYSIS B:2

be lost through oxidation. Color, odor, and turbidity may increase, decrease, 

or change in quality. Sodium, silica, and boron may be leached out of the 

glass container. Hexavalent chromium may be reduced to the trivalent state. 

This list is by no means all-inclusive. It is clearly impossible to prescribe 

absolute rules for the prevention of all possible changes ... to a large 

degree the dependability of water analyses must rest upon the experience and 

good judgment of the analyst.

Even the concentrations of the more stable dissolved constituents 

may vary during prolonged storage, and a lapse of more than 3 

months between collection and analysis is not recommended. Once 

opened, the sample should be analyzed as rapidly as possible. 

Changes in water quality are greatly accelerated after the bottle 

seal has been broken. Composite samples should be prepared immediately before analysis.

Although equilibrium between sediment and dissolved material 

can be assumed for the time of collection, this is no assurance that 

the same equilibrium will be maintained during storage. Solutions 

of the sediment, ion exchange, or other chemical reactions that 

affect the chemical composition of the sample are all possible. But 

little definite information is available on the problem; much more 

work must be done before substantial conclusions can be made.

B:2 SAMPLE-BOTTLE CASES

Sample-bottle cases commonly used by the Geological Survey are 

shown in figure 7. Cork or rubber stoppers may jar loose unless 

the bottles are packed in a shipping case that exerts pressure on 

the stopper. Containers should be marked clearly "Liquid in Glass" 

and "Keep from Freezing."

B:3 SAMPLE STORAGE

It is a wise practice to check and tighten the seals of bottles 

as the samples are put on the storage shelves. High temperatures 

should be avoided. Stored samples should be left undisturbed, 

as much of the suspended material will settle on standing, and 

they should be protected from light, which accelerates photochemical 

processes that may alter the chemical composition of the sample. 

Growth of algae in stored samples is particularly troublesome.

B:4 PREPARATION OF COMPOSITE SAMPLES

When samples are collected at frequent intervals at a station, 

analysis of each individual sample would be prohibitive in cost and 

for many streams would result in needless duplication. Consequently, the compositing of several samples into a single sample 

for chemical analysis is a common practice. This is solely a pro-
B:4 PREPARATION OF COMPOSITE SAMPLES
34 HANDLING OF SAMPLES BEFORE ANALYSIS B:4

cedure of expediency, and the chemist realizes that in so doing 

the chemical characteristics peculiar to each individual sample are 

not identifiable in the composite. Likewise, it should be recognized 

that for some streams even the analysis of daily samples would not 

adequately define the changes in stream quality. The length of the 

composite period is always a compromise between the cost of analysis and the desire to determine the quality pattern of the stream 

in minute detail. Early studies indicated that periods of approximately 10 consecutive days were optimum; small variations in 

quality were smoothed out, but major changes were still shown by 

?>6 composite samples per year. For each month, three composite 

samples were regularly prepared by mixing together equal quantities 

of daily samples collected from the 1st to the 10th, llth to the 

20th, and the remainder of the month.

In the 1930's, changes in compositing methods were brought about 

by the general acceptance and use of specific conductance as an 

approximate indication of the total concentration of ionized material 

in solution. Although generally holding to the principle of 10-day 

periods, split composites were made on the basis of total dissolved 

mineral content as indicated by measurements of conductivity of 

the daily samples, supplemented by other easily obtained information such as concentration of chloride, river stage, weather conditions, appearance of sample, and previous history of the stream. 

This modification is now used by many laboratories and is recommended in preference to a fixed time interval of compositing samples 

because it affords the chemist more discretion in deciding which 

changes in concentrations are significant and worthy of more intensive study.

Short of continuous measurement of all constituents, there are 

2 fundamental categories of data application which can be met by 

2 different methods of compositing the time-weighted metnod and 

the water-discharge-weighted method. One area of data application 

requires knowledge of the chemical composition, as concentration 

per unit volume, of only a fraction of the total flow. Analyses 

of time-weighted composites provide information on the quality 

of water available for uniform-quantity withdrawal, such as municipal or industrial diversions and domestic uses, or on the effect 

of water as an environment, such as the effect of water quality 

on growth of crops or on plant and animal life. The preparation 

of time-weighted composites simulates these uses of the water, and 

therefore chemical analyses of such composites are best suited for 

consideration of these problems.
B:4 PREPARATION OF COMPOSITE SAMPLES 35

The other field of data application requires consideration of the 

whole stream as a dynamic agent transporting and capable of 

receiving material. Whereas the total quantities of water and dissolved matter that pass the cross section on the stream are irrelevant 

to the problems of uniform withdrawal and environment, they 

are vital to problems of solute transport (dissolved load). Knowledge of solute transport of a stream is essential to most quantitative 

considerations of water quality such as the geologic processes of erosion of rocks and soils and modification of landforms, the quantity 

of pollutants entering a stream or the capacity of the stream for safely 

receiving a quantity of waste material, the salt balance on irrigated 

lands, the action of evapotranspiration on land or in streams or reservoirs, and the anticipated quality of the water if impounded. Analyses of discharge-weighted composites supply information applicable to such problems.

No single method of compositing provides data applicable to 

all problems confronting the chemical-quality hydrologist, and he 

must be cognizant of the effects of the compositing method on the 

analytical results and be so guided in the application of the data. 

When there is doubt about which method to use, the decision must 

be based on the ability to extract or compute from the data the 

type of information that may be required for understanding the 

stream's water quality and for utilization of the water.

Theoretically, for either method of compositing, all samples for a 

major period such as a year could be included in one composite. 

Actually this cannot be done because changes in quality occur during 

storage of samples and because data representative of water quality 

during much shorter periods are needed. Fluctuations in quality 

during the sampling interval influence the degree to which the analysis of a composite is representative of the stream. It is improbable 

that any one sample taken during a period of fluctuating quality can 

possibly represent the mean concentration for all dissolved constituents because the factors controlling the concentration of the individual constituents are not identical. For example, the time of mean 

calcium concentration may not coincide with that of mean sodium 

or nitrate concentration. Also, the ability to choose in advance the 

time at which the mean for the interval could be sampled is even 

less likely. Consequent!}', the probability that a composite analysis 

represents the streamflow becomes a statistical problem of sampling, 

controlled by both the range of fluctuations and the number of 

samples included in the composite.

Both methods of compositing leave appreciable latitude in the 

number of composites that can be made per year. No hard and fast
36 HANDLING OF SAMPLES BEFORE ANALYSIS B:4a

rules for compositing can be set forth because of the differences in 

range and the rapidity of fluctuations in the quality of different 

streams and because of the diversity of purposes for which chemicalquality investigations are conducted. The following suggestions are 

presented as a general guide but are by no means mandatory, and 

their application should be tempered by the chemical characteristics 

of the streams at the particular sampling station and by the significance of the observed variations in quality to present and anticipated 

use of the water-quality record. The number of composites per 

year should not ordinarily be greater than 40 or less than 20. 

Analyses of less than 20 composites per year are usually insufficient 

to define adequately the water-quality pattern of the stream unless 

the chemical characteristics are exceptionally stable, such as may be 

true below a reservoir. The cost of running more than 40 analyses 

per year increases the overall station-operation cost appreciably, and 

there is always some question as to whether the more detailed information justifies the additional expenditure. No single composite 

should cover more than a 30-day period. Suggestions regarding 

selection of compositing periods are presented in the detailed discussion of each compositing method.

B:4a TIME-WEIGHTED COMPOSITE SAMPLES

If the chemical analysis of a composite sample of a stream is 

made primarily to determine the average concentration of chemical 

constituents in the water that flowed past the sampling station 

during the period, the composite should theoretically be obtained 

by continuously diverting and thoroughly mixing a constant volume 

of streamflow. To the extent that the individual samples define 

the water quality during the compositing period, satisfactory timeweighted composites can be prepared by combining volumes that 

are proportional to the sampling interval for each individual sample. 

The sampling interval is the length of time that a sample represents; generally it is half the time between the time the preceding 

sample and the following sample were taken. At many sampling 

stations one sample is collected each day at approximately the same 

time of day; therefore, the sampling intervals and the volume of the 

samples taken are constant.

In preparation of time-weighted composites the compositing 

periods are selected on the basis of observed differences in water 

quality exhibited by some single characteristic or group of readily 

determined characteristics of the individual samples, such as specific 

conductance, pH, or chloride. The purpose of utilizing conductance 

or other characteristics in defining the limits of the compositing 

period is to provide a composite that contains only water of some-
B:4b PREPARATION OF COMPOSITE SAMPLES 37

what similar chemical characteristics. The compositing period represents a period during which it is assumed that no large changes in 

composition have occurred. As a flexible guide to the selection of 

compositing periods it is suggested that the composite include only 

those consecutive daily samples whose individual conductances are 

within ±15 percent of average conductance for the group. Any 

percentage other than 15 can be used to keep the number of composites per year within the optimum range. A knowledge of stream 

characteristics at the sampling station is essential for compositing 

and can be gained only through experience. As the chemist learns 

more of the chemical behavior of the streams in his area, his judgment in selecting compositing intervals will improve.

The degree to which time-weighted composites are representative 

of the stream during the compositing period is not dependent on 

the degree to which each sample represents the arithmetic mean 

concentration of each individual constituent for the sampling interval, but rather on the extent to which the algebraic sum of the 

positive and negative deviations of the individual sample concentrations from the true mean concentration for the individual sampling intervals approaches zero.

B:4b DISCHARGE-WEIGHTED COMPOSITE SAMPLES

If the chemical analysis of a composite sample of a stream is made 

primarily to determine the concentration of chemical constituents 

in the total volume of water that flowed past the sampling station 

during the compositing period, the composite should theoretically 

be obtained by continuously diverting and thoroughly mixing a 

constant fraction of streamflow. Practically, sampling must be periodic rather than continuous, but the volume of each individual sample that goes into the composite can be varied in proportion to the 

rate of discharge at the time of sampling. To the extent that the 

products of dissolved material in the individual samples and the discharge rates define the load of dissolved material carried by the stream 

during the compositing period, satisfactory discharge-weighted composites can be prepared from individual samples by combining volumes that are proportional to the stream discharge at the time of 

sampling. If the individual samples have different sampling intervals, the volume to be taken from an individual sample should be 

proportional to the product of the sampling interval and the water 

discharge at the time of sampling.

The preparation of discharge-weighted composites requires waterdischarge data for the time of sampling. At sampling stations 

where the relations between water stage and discharge are stable, 

the water discharge can be determined by using the gage height at
38 HANDLING OF SAMPLES BEFORE ANALYSIS B:4b

the time of sampling and a stage-discharge rating table or curve. 

More precise discharge data are desirable for a shifting stream, 

where significant error is incurred by the use of stage-discharge 

tables or curves. The required data on streams that have shifting 

controls or instructions for the computation of provisional instantaneous discharges satisfactory for compositing purposes can be 

obtained from the district offices of the Surface Water Branch of the 

Water Resources Division.

When samples are composited according to discharge, the subdivision into compositing periods usually has little effect on the 

computed total tonnages (loads) of mineral constituents for a 

period of a month or a year. The accuracy of computed tonnages 

i? not increased by the use of short compositing periods nor by 

breaking compositing periods when abrupt changes in specific conductance occur. However, the chemical-quality records serve many 

purposes, and much additional factual information concerning the 

water-quality characteristics of the stream is provided by the judicious selection of compositing periods.

Both specific conductances, or other readily determined characteristics of water quality, and water discharges representative of the 

individual samples are used as a guide in determining when to end 

one compositing period and begin another. When combined, these 

factors often aid the chemist in anticipating changes in chemical 

composition which might not be evident from either alone. Specific conductance also aids in selecting individual samples to represent high and low concentrations. However, of the two variables 

discharge and conductance, the former is given somewhat more 

consideration than the latter in the selection of compositing periods. 

A knowledge of the chemical-quality characteristics throughout the 

water-discharge range permits construction of a reliable chemicalconcentration rating curve based on water discharge as the independent variable. Inasmuch as the selected compositing period has little 

effect on the validity of computed loads, these periods are chosen 

primarily to show the relation of concentration to water discharge. 

Consequently, the compositing period should represent a limited 

range of water discharge. As a flexible guide it is suggested that 

the composite sample include only those consecutive samples whose 

corresponding instantaneous discharges are within ±25 percent of the mean instantaneous discharges for the group. Any 

percentage other than 25 can be used to limit the number of composites per year to a feasible level (see sec. B:4). If the stagedischarge relation of the stream is known to shift appreciably, and 

if the water discharge must be taken from a rating curve, one com-
B:5 REMOVAL OF TURBIDITY 39

positing period should be ended and another begun at times of 

radical change in discharge. A shift in the stage-discharge relation 

during the compositing period upsets the ratio of the computed 

provisional discharges to true discharges, and correct proportional 

quantities will not be taken for the composite. These shifts usually 

are caused by changes in stream velocities that result in scour and 

fill of the stream channel at the gaging-station control. The tonnage 

figure is computed later from the true-discharge figures supplied 

by the Surface Water Branch and not from the provisional discharges used to make the composite.

The degree to which a discharge-weighted composite represents 

both the average concentration in the total mass of water passing 

the sampling station and the load of each and all determined constituents transported during the compositing period is dependent 

on the degree to which the summation of the individual products of 

concentration of constituents and water discharge approach the 

true average concentration and loads of the whole volume of water 

that flowed past the station. Deviations of individual-sample concentrations from the mean concentration or deviations of instantaneous discharge used in the compositing process from mean discharge 

for the sampling interval do not in themselves affect the validity of 

the analysis of the composite. This is true because, as with timeweighted composites, the validity of an average is proportional to 

the extent to which the algebraic sum of the positive and negative 

errors approaches zero and not on the magnitude of the individual 

errors.

B:5 BEMOVAL OF TURBIDITY

Every water chemist has been confronted with turbid samples. 

All natural-water samples contain varying amounts of ionized and 

nonionized constituents in true solution, less finely divided material 

in colloidal suspension, and coarser readily separable material. The 

analytical statement will tend to represent the total ionic concentration in the fraction of the total sample taken for analysis regardless 

of whether the constituent is in solution or suspension. Inasmuch 

as the suspended material often contains ions determined in the 

analysis, the reported concentrations are dependent on the separation of the solution and solid phases. A filtration assembly made of 

transparent plastic and designed by the Geological Survey is shown 

in figure 8.

Most authoritative compilations of analytical procedure recognize 

the significance of finely divided material in suspension but omit 

specific instructions for its separation from the water. Comparable 

analytical results demand that the water-quality sample taken for
40 HANDLING OF SAMPLES BEFORE ANALYSIS B:5

FIGURE 8, Sample-filtration assembly.

chemical analysis be defined in terms of the maximum particle size 

of the suspended material. The specific method of separation 

should produce a minimum of change in the concentrations of the 

constituents in solution. No ideal method for removing solids exists, 

but tests have indicated that filtration through the commercially 

available cellulose acetate membrane has a minimum of undesirable 

qualities compared with other available methods. As a matter of
B:5 REMOVAL OF TURBIDITY 41

convenience, the maximum particle size of suspended material can 

be limited to that which passes by gravity through a filter whose 

retention rating is 0.5 microns. This specification does not imply 

that all settled solids should be resuspended before filtration, nor 

does it imply that obviously clear samples should be filtered. Determinations of pH, alkalinity, and (or) acidity should be made before 

filtration in the laboratory because these characteristics are related 

to the gas content of the sample.

REFERENCE

American Public Health Association and others, 1955, Standard methods for the 

examination of water, sewage, and industrial wastes: New York, Am. Public 

Health Assoc., Inc., 10th ed.

SECTION C. ANALYSIS OF WATER SAMPLES

In analysis of water the chemist is confronted with a wide variety 

of complex aqueous solutions. Most water in streams, lakes, and 

ground-water reservoirs probably contains most of the elements, 

although many are present in only extremely small quantities. The 

scope of analyses by the Geological Survey is generally limited by 

the significance of a constituent in relation to the other dissolved 

material and to the usability of the water, and by the precision of 

the analytical methods used.

C:l ANALYTICAL TECHNIQUES IN WATER ANALYSIS

Water analysis may be termed "microchemistry with macrovolumes" because all determinations require the quantitative determination of either milligram or microgram quantities or fractions 

thereof. Techniques by which the water analysis is made include 

gravimetric, volumetric, spectrophotometric, and flame-photometric 

procedures.

Gravimetric techniques in water analysis are generally tedious 

and time consuming because of the usual necessity for precipitation, 

filtration, washing, ignition, and weighing. Volumetric analysis is 

usually more rapid than gravimetric analysis if the titrant reagent 

is specific for the constituent; the sensitivity and (or) precision may 

also exceed that of a gravimetric analysis for some determinations. 

Progress in instrument development, particularly during the last 

decade, has provided the water analyst with truly precision instruments in spectrophotometers and flame photometers. Spectrophotometric methods are well suited to water analysis because small 

quantities of substances can often be determined readily and accurately. The differences in techniques of individual analysts are also 

minimized. The accuracy and precision of flame-photometric determinations of the alkali metals surpass the available gravimetric 

method, and the time required is much less.

C:la SPECTROPHOTOMETRY

Colorimetric determinations based on visual estimation of color 

have long been used in water analysis, but the widespread application of colorimetric analysis in the refined form of instrumental 

spectrophotometry is a relatively new development. Because of its 

youth, the terminology of spectrophotometry is in a confused state. 

The system of terminology used in this report follows that recom43
44 ANALYSIS OF WATER SAMPLES C:la-l

mended by the National Bureau of Standards (1947), a system 

which was derived as a compromise of several recognized terminologies.

C:la-l TERMINOLOGY AND PROCEDURES

In the practical application of spectrophotometric analysis, only 

three terms need to be considered in detail. These are the fundamental 

term "transmittance" and the two terms "transmittancy" and "absorbancy" which are derived from it. Transmittance is defined by the 

equation:

T= 

where

T= transmittance,

11 = radiant energy incident upon the first surface of the sample, and

72 =radiant energy leaving the sample.

The term "transmittance" relates to the rectilinear transmission of 

homogeneous radiation through a homogeneous, isotropic, nonmetallic 

medium having plain, smooth, parallel surfaces. In analytical practice the transmittance is not measured. The quantity that is indirectly measured is a ratio of the transmittance of the sample solution 

to the transmittance of the solvent (or blank) solution. The name 

"transmittancy" is given to this ratio. It is seen from the following 

relation that transmittancy (Ts ) can be measured directly in a spectrophotometer.

 _T sample__72 sample 

8 T solvent ~72 solvent 

where

T= transmittance, and

/2=radiant energy leaving the sample.

It should be noted that the transmittancy relation is correct only if 

the reflection and absorption losses at the cell windows are insignificant. 

These conditions are met in normal applications of visible-light spectrophotometry in water analysis.

The third term of importance and the one most generally used in 

absorptiometric work is "absorbancy," which is simply the negative 

logarithm of the transmittancy.

A 8 = log10 Ts =logio -fp

In the preparation of spectrophotometric curves of light-intensity 

ratio plotted against concentration, it is preferable, for convenience, 

to use absorbancy as the basis of the plot. Under these conditions 

a system that conforms to Beer's law gives a straight-line plot, and the 

commonly used colorimetric systems that do not conform will usually 

show only a moderate curvature. Extreme curvature, when the curve
C:la-l ANALYTICAL TECHNIQUES 45

is plotted on the basis of absorbancy data, is sometimes a sign that the 

system is not sufficiently stable for analytical purposes. Semicolloidal 

suspensions of colored substances often give extreme curvatures. 

When transmittancy data are used for plotting, a curve is always obtained unless semi logarithmic coordinates are used. The modern 

spectrophotometers have an absorbancy calibration as well as the conventional "percent transmission" or "percent transmittance," and it is 

increasingly prevalent practice to use the absorbancy scale. The relations between transmittancy and absorbancy plots for potassium permanganate solutions at three wavelengths are illustrated by Mellon 

(1950, p. 95, figs. 2, 3).

Several other terms for light absorption are given in the literature 

and are still found on the printed scales of some photometers. "Optical 

density" is often used. Sometimes optical density is the same as 

absorbancy, and sometimes it represents absorbancy multiplied by 10. 

"Extinction" is sometimes used, as is "percent extinction." Even a 

few purely arbitrary scales have been used.

The fundamental expression that correlates the light-absorption 

reading of the instrument with the concentration of the light-absorbing 

constituent in the solution is Beer's law. The law is expressed by 

several mathematical formulations, and the derivation of each may be 

found in standard texts. One of the simplest statements of Beer's 

law is in the terminology recommended by the National Bureau of 

Standards (1947) :

A s=asbc 

where

a s = "absorbancy index," 

b thickness of the cell, 

c= concentration, and 

A s scale reading.

Since the terms a,« and 5 are constant for a given analytical system and 

a given absorption cell, it is apparent that As is directly proportional 

to c.

The relation demonstrated by Beer's law is very useful and 

apparently is rather generally true for solutions of a single colored 

constituent and when no chemical reaction is involved. But in 

analytical practice many colorimetric systems show deviations from 

the law. Most apparent failures of Beer's law are due to chemical or 

instrumental properties. Some of the chemical effects responsible 

for departure are:

1. Small changes in pH acting on a pH-sensitive color system. Most organic 

indicators are weak acids or bases and consequently they are pH sensitive.

2. Temperature effect. A rise in temperature of a color system generally causes 

the absorption peak to shift toward the red.
46 ANALYSIS OF WATER SAMPLES C:la-l

3. Change in refractive index with concentration.

4. Equilibria displacement with concentration change.

o. Reciprocal interaction of the absorbing constituents. This is characteristic

of dithizone and other reagents. 

6. Salt effect. This applies mainly to absorbing solutions of electrolytes, but

the effect is sometimes present even if the absorber is not an electrolyte.

The major instrumental cause of deviations from Beer's law is 

the use of an impure spectral band for illumination of the sample. 

If the spectral band covers an area where the absorbancy curve of 

the constituent shows a sharp change with wavelength, a deviation 

will appear. By using the narrowest possible spectral band, this 

effect is minimized. The analyst is referred to standard texts for a 

complete explanation of this effect.

A second frequent source of instrumental deviations from Beer's 

law is a nonlinear photometer unit in the spectrophotometric instrument. If either the photocells, amplifying unit, or indicating unit 

do not respond linearly to a variation in light intensity, a deviation 

from Beer's law will be evidenced. Most commonly such effects 

originate in aging photocells and can be remedied by replacing the 

photocell. It should be emphasized that the common practice of 

making determinations by comparison (by comparing the sample 

against a blank) is no protection against nonlinear amplification of 

the signal. The effects do not "cancel out."

An important consideration in the use of spectrophotometric 

instruments is the absorbancy range within which the particular 

instrument will give accurate readings. The relative-error relation 

of a spectrophotometer is expressed by the curve shown in figure 9, 

from which it is apparent that the best accuracy is attained in the 

absorbancy region from approximately 0.1 to 1.0. Theoretically, 

the point of minimum relative error is approximately 37-percent 

transmittancy, which is 0.43 on the absorbancy scale. When filter 

photometers and simple spectrophotometers are used, it is necessary 

to observe the above limitations if the measurements are to be 

acceptably accurate. Absorbancy measurements above 1.0 should 

not be taken; dilution of the sample, or a similar expedient, should 

be used to bring the reading down into the reliable area. At the 

other end of the range, it is frequently impossible to avoid taking 

measurements in the area of poor accuracy; however, enrichment of 

the sample is sometimes possible and should be practiced when 

convenient.

The situation with respect to accuracy is somewhat different with 

instruments such as the Beckman B and DU spectrophotometers 

that have a step wise variable-sensitivity adjustment. By use of
C:la-2 ANALYTICAL TECHNIQUES 47

the higher sensitivities for reading dense solutions the accuracy is 

much improved, and readings up to approximately 2.5 can be taken. 

The flat part of the error curve in figure 9 is extended on the left 

side to higher absorbancy values by this device. However, the 

variable-sensitivity adjustment gives no improvement in the very 

low absorbancy region at the right side of figure 9.

3-

10 20 30 40 50 60 70 

TRANSMITTANCY,IN PERCENT

80 90

00 1.0 0.7 0.6 0.5 0.4 0.3 0.2 0.1 0

A8SORBANCY 

FIGURE 9. Relative-error curve for spectrophotometers without sensitivity adjustments.

C:la-2 APPLICATIONS OF SPECTROPHOTOMETRY TO WATER ANALYSIS

The mass application of instrumental spectrophotometry to water 

analysis has enected rapid advances in the determination of trace 

metals in water. In analyses for trace metals, a point of major 

importance is the sensitivity of determination. In this respect the 

spectrophotometer is superior to the polarograph, and for certain 

metals it is also superior to the spectrograph. Only the fluorimeter 

possesses a generally higher inherent sensitivity, but at the present 

time the application of fluorimetry to different types of analytical 

problems is limited.

In spectrophotometric determinations in water analysis it is generally desirable to attain maximum sensitivity. This confers several
48 ANALYSIS OF WATER SAMPLES Cllar-2

advantages besides the obvious one of extending the determination 

to the lower concentration ranges : The interferences of other ions can 

be minimized, the effect of natural color and turbidity in the water 

sample can often be made insignificant, and a small volume of sample 

can be used. The use of a highly sensitive reagent permits dilution 

of the sample to a point where interfering ions are rendered innocuous. The necessary dilution can be simply and automatically made 

by adding a large volume of dilute reagent to the sample, as opposed 

to the more common practice in other types of analysis of adding 

a small volume of fairly concentrated reagent. An example of this 

principle is the Eriochrome Cyanine R fluoride determination, in 

which 25 ml of a dilute reagent is added to 10 ml of the water 

sample, thus diluting the sample to minimize the interference of 

phosphate, aluminum, and other ions.

The same principle applies to minimizing the effect of turbidity 

and color. The natural color and turbidity in many water samples 

show a significant absorbancy at the wavelengths used in a number 

of determinations, and this effect requires either compensation or 

elimination. In some cases it is possible to select a spectrophotometric reagent of such high sensitivity that the absorbancy of the 

sought constituent will exceed the arbsorbancy of the natural color 

and turbidity by a very large factor. If this factor is as high as 

50 for a particular determination, the error introduced by the natural 

color and turbidity will be only 2 percent, and in routine wT ork no 

compensation would be required. Several of the spectrophotometric 

determinations described in this manual, notably the chromate and 

phosphate determinations, will normally give an absorbancy reading 

considerably higher than 50 times the absorbancy of the natural 

coloring matter in the water sample. In other procedures, such as 

the aluminum determination, the absorbancy of the color and turbidity may be as high as that contributed by the analytical system. 

Thus, correction for color and turbidity is very rarely required for 

chromate and phosphate, but correction is frequently required for 

aluminum. A knowledge of the relative sensitivity for the sought 

constituent as compared to the natural color and turbidity in the 

sample is necessary to the application of spectrophotometric methods. 

Sufficient data are given in each procedure so that the analyst can 

approximate the extent of this anticipated interference in the application of selected methods.

Sensitivity is a word that is used loosely, and it is important to 

identify the type of sensitivity under consideration. The sensitivity 

of a method depends on both the chemical phase of the determination
C:la-2 ANALYTICAL TECHNIQUES 49

and the instrumental phase. Generally, the chemical phase provides 

the major differentiation of the sought constituent against the 

"background," although the instrument used to analyze and measure 

the system also contributes. The type of sensitivity under consideration is that which improves the absorbancy of the sought constituent 

relative to the background absorbancy, which is made up of the 

absorbancy of interfering ions, natural color, and turbidity. Merely 

increasing the optical path of the absorption cell does not improve 

the relative sensitivity of the sought constituent against the background because the absorbancy of both is increased proportionately. 

To obtain a useful improvement in the sensitivity it is necessary to 

select an analytical method that will provide a greater absorbancy 

for the sought constitutent without simultaneously raising the background value.

While maximum sensitivity is usually desirable for water analysis, 

the more sensitive analytical systems are more subject to contamination effects, and their use requires careful technique. Laboratory 

glassware will contribute lead and possibly other metals. In all 

sensitive lead determinations it is necessary to boil the glassware in 

dilute nitric acid to remove surface lead. Pyrex glassware is still 

the mainstay of the trace-metals laboratory, although polyethylene, 

polystyrene, and other types of plastic equipment are available for 

reactions that do not require heat. Stainless steel, aluminum, copper, 

and nickel beakers are available and might well be used more extensively for special contamination problems. In some tests the one or 

more metals contributed by the metal beaker might be less of an 

interference than the several metals extracted from glass.

A more serious source of interference in sensitive work is the 

reagents. Only reagents of the highest purity should be used, but 

even these will show some impurity if the test is sufficiently sensitive. 

Therefore, it is always advisable to use the minimum amount of reagent that accomplishes the objective. For example, there is no 

point in using 10 g of sodium acetate to buffer a solution if 1 g will 

buffer adequately. For some determinations the special spectroscopic-purity reagents, such as those supplied by the Johnson Matthey 

Co., can be used to advantage. Distilled water is another source 

of metallic contamination, but it can usually be satisfactorily cleaned 

up by the ion-exchange resins. In general, contamination should 

not be a major difficulty in spectrophotometric work.

The end result of analysis is affected not only by changes and 

impurities in the reagents but by instrumental variations as well 

(instrument maintenance is discussed in sec. C:2d). An overall 

check on the performance of each component, instrumental and
50 ANALYSIS OF WATER SAMPLES C:la^2

chemical, that enters into the determination of the final result is 

desirable. This can best be done by running standards at frequent 

intervals. For some determinations where the variations seem to be 

particularly pronounced and uncontrollable, it may be necessary to 

run a series of standards with each set of analytical determinations, 

as with the dianthrimide determination of boron. This method will 

be of no benefit in controlling purely random variations, for which 

there is no effective control. But if the absorbancy shows a slow 

and comparatively small consistent change from day to day, it is 

valid to run standards each day as a means of correcting for small 

departures from, the calibration curve. A calibration curve is normally made up of a comparatively large number of points, with the 

mean curve draAvn through them. Thus the small random errors 

in the individual determinations tend to average out. Such a calibration curve made of many points has an inherent reliability 

superior, to that of the 1 or 2 standards that would be run with each 

set of determinations to check the curve. Therefore, these "checking" standards should be run simply to detect changes in the calibration curve. Absorbancy values should not be calculated on the basis 

of just one standard. The proper procedure is to prepare a new calibration curve if it is found that the single standard has revealed 

an actual shift in the absorbancy calibration curve.

In the consideration of the significance of sensitivity in spectrophotometric water analysis, the effect of natural color and turbidity has been mentioned. This is a recurring problem and should 

be considered in detail. Waters containing high concentrations of 

humic matter, such as swamp water, have an intense yellow color. 

Iron, especially in the form of the colloidal oxide, gives a yellowtan color. These two colors are common and cannot be distinguished 

by visual inspection. Both give similar absorbancy curves. Curve 

A in figure 10 shows the absorbancy-wavelength curve of swamp 

water from Dade County, Fla., that contained a very high concentration of organic matter. The color reading on the Hellige scale is 

140. This curve is characteristic of water that appears yellow tan 

to the eye. Curve B shows the absorbancy curve of a water that 

contained colloidal iron. The reading on one, Hellige scale was 45. 

The similarity between the curve for organic coloring matter and the 

curve for iron is obvious.

Curves A and B illustrate a very important phenomenon in spectrophotometric water analysis. The interference of the "normal" 

natural color is much greater in the blue region of the spectrum. 

Therefore, there is a very definite advantage to using analytical systems that absorb in the yellow-red region. This is one reason why
C:la-2 ANALYTICAL TECHNIQUES

WAVELENGTH OF DETERMINATION

51

350 400 450 500 550 600 650 700 

WAVELENGTH, IN MILLIMICRONS

FIGURE 10. Absorbancy curves for colored and turbid samples. Absorbancy readings 

were taken with Beckman B spectrophotometer, using 40-mm absorption cells.

most of the spectrophotometric methods in this manual utilize wavelengths of 500 m/A or greater. Methods that require measurement in 

the blue region are included only when their merits are sufficient to 

outweigh the detrimental effect of high color interef erence; the ferron 

determination of aluminum is an example.

The color curves in figure 10 permit estimation of the approximate 

interference that the "normal" type of yellow color in a water sample 

will give at the wavelength of a particular determination. When 

the effect of color is not too great, a correction could be made on this 

basis. This method could be applied only to samples with the 

familiar yellow-tan color indicating humic matter and iron oxide. 

For samples containing colored industrial wastes, petroleum products, 

and other material with "abnormal" colors, new curves would have
52 ANALYSIS OF WATER SAMPLES C:lBr-2

to be prepared. Where a laboratory is analyzing colored waters 

of consistent composition on a routine basis, it would be profitable 

to prepare color-correction tables for each spectrophotometric method, 

thereby eliminating the necessity to include the color-compensation 

step in each analysis. The approximate effect of the interference of 

the natural color represented by curve A in several spectrophotometric determinations is given below:

Wavelength Error

Al______.____________________ 370 +0.6

Cu___________________________ 425 +.3

Fe___________________________ 520 +.1

F____________________________ 540 -.15

PO4 __________________________ 700 +.001

The effect of natural color may be compensated in several ways, of 

which four are described below.

1. Subtraction of natural-color absorbancy. Determine absorbancy 

of the test sample, A t », against the blank specified for the procedure. 

Determine absorbancy of the natural-color sample, Anc8 , against distilled water using the same spectrophotometric conditions as for the 

test sample. The difference is the corrected absorbancy, Ag . Use As 

to obtain concentration values.

The test sample is the water sample with color developed as outlined in the analytical procedure.

The natural-color sample can be prepared in two ways. The most 

general method is to take the same volume of sample water as was 

used for the test sample. Treat it exactly as the test sample with 

one exception: do not add the indicator reagent. Instead add an 

equal volume of indicator solvent, usually dilution water. The 

second method uses color- and turbidity-correction solutions, which 

are simply shortcuts of the above general procedure. These solutions combine the significant reagents into one solution, so that a 

single addition of reagent suffices. The color-correction solutions are 

not applicable in every case. The full procedure is usually a little 

more reliable.

The subtraction method is the most generally applicable correction 

method. It can be applied to turbidity corrections as well as natural 

color. The method fails where the indicator reagent reacts with 

or affects the natural color or turbidity in the water sample or where 

the effects of the other reagent on the test sample and natural-color 

sample are not essentially identical. The latter qualification relates 

more to turbidity than color, and filtration of excessively turbid 

samples may be required.
C:la-2 ANALYTICAL TECHNIQUES 53

2. Direct compensation. The direct-compensation procedure is illustrated by the cuprethol copper determination. All the reagents 

except cuprethol, the indicator, are added to the sample. The sample 

is placed in the spectrophotometer and the absorbancy is set to zero 

by adjusting the slit width. This step eliminates the natural absorbancy of the water. The indicator reagent is then added. The 

color is developed and the correct absorbancy is read directly from the 

scale. This procedure is not of general utility and can be applied 

only when conditions are right. One requirement is that the absorbancy curve should have only a shallow slope in the operating 

region.

3. Bleaching. If the indicator reagent reacts with the material 

that gives the sample its natural color, the resulting color must be 

removed by bleaching. One example of such a reaction is the periodate determination of manganese. Periodate oxidizes manganese 

to permanganate, developing the familiar pink color. Periodate 

also partially oxidizes the organic coloring matter in the water 

and changes the natural color. Therefore, compensation by means 

of a natural-color sample will not be correct. In the permanganate 

test the procedure is to develop the color in the test sample and determine the absorbancy. A few crystals of sodium nitrite are then 

added to the sample. This destroys the permanganate color immediately but does not affect the organic matter. The color contribution of the organic matter can now be directly measured. It is 

subtracted from the absorbancy of the test sample to obtain the 

corrected absorbancy.

4- Removal of color material. The color material is sometimes 

removed with conventional activated-carbon or alumina treatments. 

The method is mentioned only to cite its limitations. Gross contamination from impurities in activated carbon is an ever-present 

possibility. Furthermore, the carbon will adsorb trace metals and 

can completely change the content of trace metals in the sample. 

Alumina is less likely to be a source of contamination, but it will 

take up trace metals just as avidly as activated carbon. These two 

methods cannot be used for sensitive analyses. The acceptable procedures for removing color are chemical oxidations with nitric acid, 

hydrogen peroxide, and similar agents.

Curves C and D in figure 10 show how the effect of natural turbidity varies with wavelength. Curve C represents a turbidity 

of 45 ppm SiO2 , and curve D represents 95 ppm SiO2 ; both measurements are on the Hellige scale. The suspensions used to prepare the 

curves were of barium sulfate with 10 ppm sulfate in C and 40 ppm 

in D. Suspension C was of fine particles, and D was of coarser
54 ANALYSIS OF WATER SAMPLES C:lBr-2

particles. The shape of the two curves is influenced by the particle 

size in the way that would be predicted by Rayleigh's equations for 

light scattering. A turbid suspension not only absorbs light but 

scatters it as well, and the scattered-light losses are indicated as 

absorbancy by the spectrophotometer. Suspension C shows a higher 

relative scattering in the blue region, as would be predicted, because 

of its smaller particle size. The significant observation for both 

curves is that the absorbancy effect of the relatively high turbidity 

in C is greater over most of the spectrum than the very high color 

represented by curve A. Only in the extreme blue end of the spectrum, at wavelengths less than 400-425 m/*, does the interference of 

natural color normally exceed that of turbidity. The approximate 

interferences of the 2 turbid suspensions in the 5 determinations cited 

previously are as follows:

Error

Wavelength C (45 ppm StO2) D(95ppmSiO^) 

(TOM) (ppm) (ppm)

Cu 

Fe.

370

425

520

540

700

+ 0. 3

+ .3

+ .3

-. 3

+ .08

+ 1.0

+ 1.2

+ 1.2

-1. 2

+.5

Fortunately, most natural-water samples are less turbid than the 

suspension represented by curve C; however, some streams.carry 

such a fine suspension of clay that the water appears opaque when 

collected. For these samples the spectrophotometric methods of correction for turbidity, treating it the same as natural color, are not 

satisfactory, and the turbidity must be removed. To date only one 

generally acceptable method for removing such turbidity has been 

found; this is filtration through membrane filters such as the cellulose acetate type supplied by the Millipore Co. (See sec. B: 5.) 

Centrifuging is often useful, but it is less efficient than membrane 

filters for fine particles. Good results can sometimes be obtained 

by filtering the sample through a fine filter paper, such as Whatman 

42, that will remove a large part of the turbidity. The residual turbidity in the sample is then corrected for by the same methods as 

for color.

The difficulty in removing turbidity is to avoid upsetting the tracemetal ionic relations in the sample. While it is not known that any 

single method accomplishes this objective ideally, it is known that 

some methods are better than others. The methods cited above, 

while not ideal, are acceptable. Some methods are not acceptable 

because of the high possibility of gross contamination of the sample 

or gross change of the sample by adsorption of constituents on the
Cllb ANALYTICAL TECHNIQUES 55

filtering medium. The methods that cannot be used in spectrophotometric work are:

1. Filtration through Berkefeld tubes and similar porous porcelain or diatomaceous-earth filters.

2. Filtration with the aid of charcoal or alumina.

3. Filtration with the aid of filter pulp unless it has been established that the 

pulp does not significantly adsorb the particular constituent being determined.

Filtration through fritted-glass discs is probably safe from the 

standpoint of adsorption losses, but it is always necessary to guard 

against contamination by carryover from previous filtrations with 

these units. Another method is the addition of large quantities of 

salts to the sample to neutralize the charges on the particles and 

cause their coagulation. What effect this procedure may have on 

trace metals has not been investigated, but it is obvious that contamination may occur because of the high concentration of salt that 

must be added.

It is impossible to specify procedures for the correction of turbidity that will fit all field conditions. Through familiarity with 

waters of the area the analyst will be able to identify the appropriate 

method for eliminating errors resulting from turbidity. Often no 

correction will be necessary. In other cases it will be possible to 

empirically prepare correction tables or curves for the streams. It 

should be necessary only occasionally to apply lengthy physical or 

chemical procedures.

C:ll> FLAME PHOTOMETRY

Flame photometry is closely related to spectrographic analysis, 

and many of the concepts of that field carry over into flame photometry. The essential difference between the two forms of analysis 

is the temperature of the source used to excite spectral emission of 

the sample. The gas-and-air or gas-and-oxygen flames are much 

cooler than the spark and arc sources used in spectrography. Hence, 

in this manual, flame photometry is limited to the easily excited 

metals sodium, potassium, and lithium.

Selection of the fuel to be used in flame-photometric wrork is important, as the nature of the fuel affects not only the emissivity of 

the sought metal but also affects the interferences that show up in 

the measurement of the metal. For example, in the Perkin Elmer 

flame photometer the interference of calcium in the barium determination is approximately seven times greater with the propane flame 

than with the acetylene flame. The. situation is reversed in the 

interference of sodium on potassium, where the propane flame is 

superior. Generally, the hotter flame is the more sensitive flame,
56 ANALYSIS OF WATER SAMPLES C:lb

and it will permit the determination of more elements. Strontium 

can be determined with the Perkin Elmer flame photometer using 

the acetylene flame, but not with propane flame. In a few cases, 

perhaps because of a lower flame background in the desired wavelength region, the lower temperature flame gives superior sensitivity. 

For the analysis of sodium, potassium, and lithium in the Perkin 

Elmer 52-C instrument, the propane .flame is preferred because sensitivity is adequate, and the propane flame is less noisy and gives a low^er 

background than the acetylene flame. The acetylene flame offers 

some advantage with the Beckman instrument, but the optimum flame 

temperatures for determination of different elements are not identical. 

Flame photometers can be operated by either the internal-standard 

or direct-intensity method. The former method is the "tried and 

true" standby of spectrographic work and is generally preferred, 

although accurate results can be obtained by the direct-intensity 

method if four conditions are essentially fulfilled:

1. The instrument is electronically stable and gives reproducible readings-as a 

photometer.

2. The flame can be precisely and reproducibly controlled. This means fine gas 

and air adjustments that will hold their settings.

3. The interference of foreign elements can be compensated by reference to 

graphs or other techniques inasmuch as the presence of the matrix constituent tends to reduce the mutual enhancements and quenching effects 

between elements.

4. The sample is fed to the flame in a constant and reproducible rate.

While the same four conditions must be stabilized to a reasonable 

degree in the internal-standard flame photometer, only condition 1 

is as critical as with the direct-reading type. Successful directreading flame photometers, such as the Beckman flame attachments, 

are characterized by precisely machined atomizers and flame orifices 

and fine gas controls that give more precise pressure adjustment than 

the ordinary regulator.

Although internal-standard flame photometers are more reliable for 

day-to-day operation, the direct-reading instruments are inherently 

more sensitive. In the internal-standard instruments the value indicated on the meter is the difference between the signal from the 

sought element and the internal-standard element; the difference, 

of course, is always less than the full signal measured with the directreading instrument. Most internal-standard instruments have provision to switch over to direct-intensity operation when conditions 

require.

One of the common phenomena of spectrography that is also encountered in flame photometry is the mutual enhancement and depression (more often the former) of emission between two or more
C:lb ANALYTICAL TECHNIQUES 57

metals. This phenomenon is not associated wT ith the overlapping 

of lines. The most frequent manifestation is the enhancement that 

sodium gives to potassium readings. This effect is said to be eliminated by various "radiation buffers" which are made up of a high 

concentration of salts. Whether these radiation buffers do everything claimed for them is open to doubt, and, furthermore, they 

have a bad tendency to clog the atomizer. The sodium-potassium 

enhancement is completely eliminated by using lithium as the internal 

standard. With the flame attachments for Beckman models DU and 

B, the effect of sodium enhancement of potassium emission can be 

eliminated by adding sufficient sodium to standards and samples 

to obtain the maximum, or "saturation," enhancement; for practical 

work this is attained at a sodium concentration of 500 ppm.

In general, the curves obtained with flame photometers do not 

show the day-to-day reproducibility that one expects from a spectrophotometric calibration curve. It is good practice to prepare a new 

curve each time the flame photometer is used by running 4 or 5 standards along with the samples. If experience shows that the instrument 

is relatively stable, then it should be satisfactory to simply prepare a 

curve for the day. The practice of "bracketing" the samples with 

standards has been rather extensively employed. Some of the more 

erratic instruments require this procedure as a routine matter. In 

bracketing, the approximate sodium and potassium concentration is 

estimated (by calculation, past experience with samples of similar 

water, or by a preliminary run), and standards with concentrations 

slightly less and slightly greater than the unknown are prepared. 

The readings of all three solutions are taken as nearly simultaneously 

as possible (it may be necessary to take several readings and average the results), and then the value of the unknown is found by 

simple proportion. Inasmuch as flame-photometer calibration curves 

are generally not linear, it is apparent that the standards have to be 

fairly close to the unknown if the simple proportion calculation is 

to be valid. A more suitable method, and one which does not require 

as many standards, is to prepare a series of celluloid curves representing the mean curve obtained with the instrument for each of the 

several operating ranges. These "master" curves can be arranged to 

rotate about an origin or move upscale or downscale to fit the instrument fluctuations as determined from the standards. Concentrations 

can be read directly from these curves. It should be noted that flamephotometer curves do not have a general configuration common to 

all elements. The shape of the curve varies with each element and 

may show linearity, an increase in slope at higher concentrations, or, 

most frequently, a tendency to level off at higher concentrations.
58 ANALYSIS OF WATER SAMPLES C:2

The Beckman photomultiplier attachment for the spectrophotometer has considerably extended the range of flame analysis in the 

blue and green region of the spectrum. Although the sensitivity continues down into the ultraviolet, this region is not of maximum 

utility because the flame background is likely to be excessive. However, in the region from approximately 400 to 600 m/x, which contains 

sensitive lines of strontium, calcium, and sodium, the photomultiplier 

attachment presents possibilities for expanding the scope of flamephotometric analysis.

C :2 INSTRUMENTS

Water analysis as performed today requires the use of several 

measuring devices. Some determinations, such as specific conductance, can be made only with instruments. The water analyst has a 

wide variety of measuring devices to choose from, and the choice is 

rightly guided by the scope of the analytical scheme and the accuracy 

and precision desired. In this section, discussion is limited to the 

minimum of instruments required by the Geological Survey to make 

the determinations included in this manual with the precision and 

accuracy required.

The selection of a given commercial product has been based largely 

on the adequacy of the instruments for fulfilling the needs, durability, cost, maintenance requirements, and advantages of semistandardization throughout the many laboratories of the Geological 

Survey. The ultimate in accuracy or current refinement is not 

heavily overweighted in choosing the instrument. Instead, the selection has been guided to a considerable degree by the experiences of 

the different laboratories a unit is purchased by one laboratory, 

evaluated, and the findings made known to other laboratories.

C:2a CONDUCTIVITY METERS

Specific conductance is determined by using a Wheatstone bridge 

in which a variable resistance is adjusted so that it is equal to the 

resistance of an unknown solution between two platinized electrodes. 

The null point is detected by an a-c galvanometer or a cathode-ray 

tube. Alternating current is necessary to prevent polarization of 

the electrodes. Direct current produces gas bubbles on the electrodes 

that greatly increase the resistance and change the concentration 

of the electrolyte in the vicinity of the electrodes. The electrodes 

are coated with a thin layer of amorphous platinum, which tends to 

adsorb gases and catalyzes their reunion, thereby minimizing polarization.

The electrode cell may be either the dip or pipet type. The pipet 

cells are generally more satisfactory for routine laboratory use because 

they require a smaller volume of water for the determination, the
C:2a INSTRUMENTS 59

water can be drawn directly from a narrow-mouth sample bottle without transferring it to another container, the total time for the determination is less, and there is less mechanical agitation of the water 

sample. Dip cells are preferable for fieldwork.

Conductivity meters differ in their design, construction, and suitability for waters of different ionic concentration, although they 

are fundamentally the same in operating principle. The Geological 

Survey must determine the specific conductance of waters that range 

widely in concentration. A satisfactory instrument should handle 

samples having specific conductances in the range of 0 to 100,000 

micromhos; the results should be reproducible and accurate to within 

3 percent of the true value.

Cell polarization is particularly troublesome when the specific 

conductance of highly mineralized waters is determined. Polarization 

varies directly with the amount of current that flows between the 

electrodes and inversely with the frequency of the current. High 

input voltage to the bridge or low resistance of the cell (concentrated 

water) is conducive to polarization. The cell resistance can be increased by increasing the cell constant (see sec. D: 37). The selection 

of the cell constant is, however, limited by the accuracy and sensitivity 

of the bridge for measuring very high and low resistances. In addition, the current frequency should not be excessively high because 

a-c resistance is a complex function of frequency; at frequencies 

necessary to avoid polarization completely, the differences between 

a-c resistance and d-c resistance may be appreciable unless the cell 

has been very carefully designed to minimize this difference. Therefore, the ideal single apparatus for the measurement of conductivity 

throughout a wide range would necessarily incorporate practical 

compromises between low input voltage, high cell constant, high current frequency, and (or) the accuracy and sensitivity of the bridge 

for measuring extreme resistances and at high current frequencies.

For many years a bridge fabricated within the Geological Survey 

has been used successfully for the determination of specific conductances ranging from about 25 to 3,500 micromhos. The essential 

elements of this bridge are a variable resistance, 0 to 9,999 ohms in 

steps of 1 ohm, with a multiplying factor of 0.001 to 1,000 for 

widening the range of the resistance box; an a-c pointer-type galvanometer which supplies current for the bridge circuit at 6 v and has 

a sensitivity of 1 /*a (microampere) per mm scale division; and an 

insulating transformer, 115 to 115 v, 60 cycles, 50 va. This bridge is 

used with a pipet cell (cell constant, approximately 0.3 reciprocal 

cm). Above about 3,500 micromhos the determined specific conductances are lower than the true values by more than 3 percent,
60 ANALYSIS OF WATER SAMPLES C:2b

and the percentage of error increases gradually as the ionic concentration of the water increases.

Several combinations of bridges and commercially available cells 

have been investigated. The line-operated Serf ass Model RCN 15 

bridge has proved satisfactory for routine laboratory determinations 

when used as a resistance-measuring device and with a pipet cell 

having a constant of 0.3 reciprocal cm (see "Wheatstone bridge 

method," sec. D: 37a-l). Current frequency of either 60 or 1,000 cycles per second can be selected. The lower frequency is preferable for 

measuring the high resistances of very dilute solutions such as distilled water, but the higher frequency is preferable for most natural 

water. The Serfass instrument also incorporates an a-c bridge with 

an input voltage of 3; a vacuum-tube amplifier; a cathode-ray-tube 

null detector; a scale of about 14 inches in length, logarithmically 

graduated from 0.5 to 1.5, which permits direct reading of resistances 

to 2 significant figures and interpolation of the third figure; and an 

adjustment knob that varies resistance ratios from 1 to 100,000 in 

6 steps, each representing a factor of 10. Other assemblies comparable in accuracy and reproducibility throughout the 0 to 100,000 

micromho range would be equally suitable.

Direct-reading conductivity meters are commercially available, 

but none investigated by the Geological Survey gave results comparable in accuracy to the resistance method described in section 

D: 37a-l. Two potential sources of error are inherent to the directreading meters; first, they must either have a cell constant of exactly 

1.0 or some type of cell-constant-adjustment mechanism; second, the 

temperature of the solution must either be exactly 25 °C or the temperature must be compensated for manually or automatically. However, some direct-reading conductivity meters have been used satisfactorily for tests when the accuracy requirements are not stringent.

C:2b BALANCES

Either manually operated or automatic balances that are sensitive 

and accurate to 0.1 mg in the 0- to 100-g range are suitable for water 

analysis. A large-capacity balance sensitive to 0.1 g is also useful 

for less accurate weighings.

C:2c pH METERS

pH meters measure the electrical potential between two suitable 

electrodes immersed in the solution to be tested. The reference 

electrode assumes a constant potential, and the indicating electrode assumes a potential dependent on the pH of the solution 

Electrode potential is the difference in potential between the electrode and the solution in which it is immersed. The calomel electrode,
C:2d INSTRUMENTS 61

which is a widely used reference electrode in water analysis, consists 

of a mercury-calomel rod immersed in a saturated solution of potassium chloride ; this electrode has a potential of + 0.246 v. Electrical 

connection with the sample is provided through porous fiber sealed 

into the immersion end. A glass electrode is normally used as an 

indicating electrode. The glass electrode has several features that 

recommend it for pH measurements. Among the most important are 

that it is not affected by oxidizing or reducing substances in the 

sample and. that it can be used in turbid and(or) colloidal suspensions. The basic design is a silver-silver chloride or mercurymercurous chloride electrode immersed in a solution of known pH 

and the whole completely sealed in glass.

The mechanism by which the glass membrane responds to hydrogenion activity is not thoroughly understood, but it probably involves 

absorption of hydrogen ions on both sides of the membrane proportionally to the activity of the hydrogen ions in solution. The 

cell for measuring the pH of a solution is of the following type : 

A,, Solution of Glass Solution of

Ag . known pH . membrane . unknown pH

Glass electrode

The voltage of the glass electrode is a logarithmic function of the 

difference in hydrogen-ion activity of the solutions on either side of 

the glass membrane. To measure this voltage an electron-tube voltmeter is used because the resistance of the glass membrane is so great. 

pH meters differ in their design, construction, and applicability to 

all phases of water analysis in which the instrument is used. Desired 

features in a line-operated pH meter are built-in voltage regulator, 

accuracy of ±0.05 pH, calibration for wide pH range with 1 buffer 

solution, stability of calibration, built-in temperature-compensating 

mechanism, durable electrodes, and a design that permits insertion 

of the electrodes, stirrer, and a buret into a suitable vessel for titrations. For pH determinations in the field the instrument should al§o 

be rugged and compact.

C:2d SPECTROPHOTOMETERS

Most spectrophotometric instruments used in water analysis are 

designed for operation in the visual region of the spectrum. Although 

some work has been done with phenols in the ultraviolet region, 

and a considerable amount of infrared-absorption study of organic 

pollutants is reported in the literature, the major application of 

spectrophotometers is still in the visual region, and the following discussion is limited to instruments of this type.

Every spectrophotometer contains four major components: the 

light source, the monochromator, the absorption-cell unit, and the
62 ANALYSIS OF WATER SAMPLES C:2d

photometer unit. The major factor that controls the quality and 

the price of the instrument is the monochromator system, which may 

be an optical filter, a diffraction grating, or a prism. The opticalfilter type is the simplest and is satisfactory for many routine analytical applications. In the use of filter photometers, the primary 

consideration is the quality of the filters, a factor that varies over a 

wide range. While there is no standard method of specifying the 

characteristics of optical filters, the use of "half-intensity" band 

widths is fairly common. The principle of the system is illustrated in 

figure 11. Curve A on figure 11 is for a typical broad-band glass 

filter, Corning 5874. The half-intensity band pass (a) for this filter 

is approximately 55 m/*,, which is about the maximum that can be 

tolerated in a useful filter for photometric analysis. By combining 

several filters into a single unit, the band pass can be narrowed to 

10 or 20 m/t. To select the simple filter components for such a com300 350

WAVELENGTH, IN MILLIMICRONS

400

FIGURE 11. Comparative energy transmission of filters. Curve A is for a 

Corning 5874 glass filter, which has a half-intensity band pass (a) of 55 m/i, 

and curve B is for a Photovolt interference filter with a half-intensity band 

pass (&) of 10 m/t.
C:2d INSTRUMENTS 63

bination, a manufacturer's catalog that gives the transmittance 

curves for the individual filters is very useful. Sometimes, filter 

combinations seriously attenuate the light intensity. To overcome 

this problem, interference filters can be used. These are mirror 

devices wherein the desired spectral region is isolated by multiple 

reflections that cancel out all but the required region. They can 

be made very sharp and in almost any wavelength. Curve B (fig. 11) 

is for an interference filter manufactured by the Photovolt Corp. 

The half-intensity band pass (&) is less than 10 m/t. Similar filters 

having a peak transmission accurate to ±1 m//, can be obtained in 

any wavelength from 400 to 840 m/t. Other sources for interference 

filters are the Baird Instrument Co. and Farrand Optical Co.

Diffraction gratings are extensively used in the Coleman instruments and are also used in the Bausch and Lomb spectronic photometer. The latter is a very ruggedly made and serviceable instrument and has a band pass of 20 m/t. This value is comparable to 

the isolation achieved with high-grade optical filters. The major 

recommendation of diffraction-grating instruments is their linear 

dispersion, which gives a constant band pass throughout the spectrum 

at a fixed slit width. For routine analytical applications, interest 

in this feature is mainly academic, but it has some advantages in 

research-type instruments.

A variety of spectrophotometers is used by the Geological Survey, 

but the most common is the prism type, as represented by the Beckman B spectrophotometer. The dispersion of a prism is nonlinear, 

being much greater at the blue end of the spectrum than at the red. 

It is important for every operator of this instrument to realize that 

when the slit-width setting for a particular determination is increased 

to compensate for changes in a reagent solution, a drop in output 

of a photocell, or for some'other reason, the band pass at the operating wavelength is simultaneously widened. In a few determinations 

where the wavelength of the absorption system is critical, this 

broader band pass could cause a shift in the analytical calibration 

with possibly an increase of curvature. Such an adverse effect would 

be the exception rather than the rule, because at normal operating 

wavelengths the band pass is so narrow as to be safely within the 

requirements of most colorimetric systems. For example, at 500 m/i 

with a slit width of 0.2 mm, the band pass is only 3 m/t. For some 

research applications it is necessary to know the band pass of the 

instrument throughout the wavelength spectrum for various slit 

widths. These data for the Beckman instrument may be obtained 

from the half-intensity band-width curves on page 6 of the Beckman 

Bulletin 291-A.
64 ANALYSIS OF WATER SAMPLES C:2d

In selecting the sample absorption cells for use in the photometer, 

it is necessary to consider the wavelength at which the determinations 

will be made. Almost any type of glass is suitable at wavelengths 

above 400 m/t. Pyrex cells can be used successfully for the aluminum 

determination at 370 m/*, but the transmission drops off rapidly 

below this region. Corex-window cells can be used at wavelengths 

down to approximately 300 m/n. Below this region silica-wT indow 

cells must be used.

Open-top absorption cells are preferred for routine work, as 

stoppered-orifice cells are too cumbersome and often trap air bubbles. 

Only sintered- or fused-joint cells should be used. Cement in joints 

will sometimes react with reagents to give false colors. Plastic 

cells should not be used. Rectangular cells are preferred for critical 

work because they do not require the critical centering necessary 

for cylindrical cells. Several types of rectangular open-top cells 

may be obtained from suppliers of photometric equipment. Absorption cells can be made to special order in a variety of shapes and 

dimensions. Many of the colorimetric methods described in this 

manual give data applicable to rectangular cells of 40-mm optical 

depth. The cell carriage of the Beckman B spectrophotometer will 

not accept these cells. However, it is relatively easy to make a carriage that will accommodate the 40-mm or larger cell. A drawing of 

one such carriage for the 40-mm cell is shown in figure 12.

After the light passes through the filter system and the absorption 

cell, it is received on the photoreceptor and is converted into electrical 

energy. There are three major classes of photosensitive devices, each 

with distinctive characteristics with which the operator should be 

familiar. These are the barrier-layer cell, the vacuum phototube, 

and the photomultiplier tube. Diagrams of the three types are 

shown in figure 13. The barrier-layer cell is the type generally used 

in the less expensive photometers because it requires no power supply 

and puts out a comparatively high current that will operate a meter 

without further amplification. Barrier-layer cells are relatively 

stable and rugged. The output of the barrier-layer cell cannot be 

easily amplified, and consequently it is necessary to use the second 

type of cell, the vacuum phototube, for the measurement of low light 

values. This cell consists of a semicylindrical cathode with a coating 

of cesium-cesium oxide or other photoemissive metal and a central 

wire anode. When light strikes the cathode, electrons are emitted 

from the surface and are drawn to the anode, which is maintained at 

a positive potential. The microampere current so obtained develops 

a voltage across a high resistance; the voltage is fed to the grid of 

an amplifier tube. A development of the vacuum phototube is the
C:2d INSTRUMENTS 65

TEMPLATE FOR BOTTOM PLATE OF 40-MM CELL CARRIAGE

FIGURE 12. Forty-millimeter cell carriage for Beckman Model B spectrophotometer.

photomultiplier tube, in which the current is enormously amplified 

within the tube itself. The emission from the cathode circles through 

a chain of 10 positive dynodes and is amplified at each one. The 

major disadvantages of the photomultiplier tube are the necessity
66 ANALYSIS OF WATER SAMPLES C:2d

Barrier-layer cell Vacuum phototube Photomultiplier tube

FIGURE 13. Schematic representation of photocells.

for high voltages (1,000 v) and a tendency toward instability. They 

vary greatly in individual characteristics, and in critical applications 

they must often be selected from several tested for certain specific 

characteristics.

In general, the output of all three types of photoreceptors is approximately linear within their appropriate working ranges. But 

for critical work the linearity of the tube output (and amplifier) must 

be checked and a correction curve plotted. As the tubes age, the 

output falls, and curvature may become pronounced.

All photocells show a marked variation of response with wavelength. The wavelength-response curves for the various types of 

photocells can be found in a bulletin entitled "Phototubes, Cathode 

Ray and Special Tubes" and published by the Radio Corporation 

of America.

Most of the spectrophotometric methods in this manual are designed for use with the Beckman B spectrophotometer because this 

instrument is common to all laboratories of the Geological Survey. 

Except for aluminum and fluoride, all of the colorimetric determinations described in this manual can also be made with the Bausch and 

Lomb spectronic instrument. The sensitivity of the Bausch and 

Lomb is approximately half that of the Beckman B, however, because the maximum absorption-cell depth is only 22 mm as compared to the 40-mm depth usually specified for the modified Beckman B. The Beckman instrument will accept special cells as long 

as 70 mm for determinations in which maximum sensitivity is required. For field analysis the Photovolt Lumetron 401 has been 

used. The Lumetron is a simple and fairly rugged instrument that 

will operate from a 6-v storage battery as well as 110-v a-c current. 

Many of the simpler colorimetric determinations can be performed 

with this instrument without excessive loss of sensitivity as compared to the Beckman B.
C:2d INSTRUMENTS 67

For the most successful and general application of spectrophotometry in water analysis, the instruments should meet certain requirements, some of which are more exacting than the normal requirements of colorimetric work. A list of the features of the ideal wateranalysis spectrophotometer for laboratory and field use is given 

below. No single instrument meets all the requirements, but a 

combination of instruments may approach the ideal specifications.

1. Useful-wavelength range from approximately 350 m/t to 900 m/u.

2. Provision to use long absorption cells (40 mm or longer) when necessary.

3. Suitable for field use. This means specifically a rugged construction and 

operation from a 6-v storage battery.

4. Suitable for spectrotitration.

5. Narrow band pass when required, and a light source and photocell combination of sufficient sensitivity so that dense filters or narrow slits may be used.

6. Provision to shift the absorbancy scale by stepwise adjustment so that 

solutions of high absorbancy may be set to "sero." This is a very useful 

device in many analyses.

7. Readily available electronic, optical, and mechanical components for replacement. Reasonable freedom from breakdown under laboratory conditions 

and a satisfactory repair service and policy on the part of the manufacturer.

8. Availability of a few common attachments such as fluorescence-measuring 

adaptors and turbidity adaptors is desirable but not necessary.

9. Reasonable facility of operation.

10. Good voltage regulation and general stability.

11. Sufficiently rugged and well made to hold its adjustments through the 

normal rigors of routine laboratory service.

In theory, all spectrophotometers of the same type should be 

expected to give the same reading for the absorbancy of a given 

solution under fixed conditions of wavelength and band pass. In 

actual practice this ideal is not realized. The disagreement between 

high-quality spectrophotometers in the measurement of absorbancy 

is much greater than the disagreement between analytical balances 

in the measurement of mass. The spectrophotometer is a more complex instrument. Many more components contribute to the result 

than in a simple, direct device like a balance. Spectrophotometric 

instruments require careful handling and close adjustment if they 

are to perform properly. Two major phases of spectrophotometer 

function require frequent checking: the wavelength setting and 

the absorbancy scale. In any laboratory that uses spectrophotometers on a routine basis, a regular schedule for testing these two 

scales should be set up on a definite periodic basis.

Wavelength-scale calibrations can be tested by two relatively 

simple methods:

I. Emission sources of known wavelength: Helium discharge tubes, the mercury 

arc, and the sodium-flame spectrum can be used as sources of known lines 

for testing the accuracy of a spectrophotometer wavelength scale. Inex-
68 ANALYSIS OF WATER SAMPLES C:2e

pensive discharge tubes suitable for the purpose can be obtained from 

several suppliers of photometric equipment. This is the most accurate 

method for testing a wavelength scale.

2. Didymium filter: Mellon (1950) reports that the National Bureau of Standards has tested the Corning No. 5120 didymium glass (3.0-mm thickness) 

and has never found any certain variation in the positions of the wavelength maxima in the absorption curve for this glass. Therefore, it is 

concluded that this didymium glass would make a satisfactory standard 

for checking the wavelength scale of spectrophotometers used in routine 

work. This method is inferior to the use of line sources but should be 

sufficiently accurate for instruments used in routine work. The prominent 

absorption maxima in the wavelength region between 400 and 808 m/i are 

as follows:

441.0

475.5

528.7

585.0

684.8

743.5

745

808

The second periodic check of importance is the accuracy of the 

photometric scale, which may show considerable variation depending upon the condition of the phototubes and the amplifier. Neutral 

filters of known transmittance have been used for checking the 

scales. Standards that are probably superior in a practical sense are 

the calibrated color glasses issued by the National Bureau of Standards (Gibson and Balcom, 1947, p. 601). These filters of known 

transmittance cover the wavelength region from 400 m//, to 750 m/x 

and simulate the properties of colored solutions. Tables of spectral 

absorbancy throughout the wavelength region 350 m//, to 750 m//, 

for a standard solution of copper sulf ate are given by Mellon (1950). 

Similar tables are given for cobalt ammonium sulf ate and potassium 

chromate. If it is desired simply to check the linearity of the 

absorbancy scale, this may be done by determining the absorbancy of 

successive dilutions of a colored solution known to follow Beer's 

law. Cobalt ammonium sulf ate in sulf uric acid solution gives a good 

spectral absorption range for such linearity tests.

Some of the routine aspects of spectrophotometric maintenance, 

such as cleaning the slits and keeping the optics in alinement, have 

not been considered here. These subjects are discussed in detail in 

the manufacturer's handbook supplied with each instrument, and 

reference is made to these for further information.

C:2e FLAME PHOTOMETERS

Both the internal-standard type and the direct-reading type of 

flame photometer are in use by the Geological Survey. The former 

is represented by the Perkin Elmer 52-C instrument and the latter 

by the Beckman B and DU flame attachments. These relatively 

expensive instruments were put into use by the Survey several years 

ago, and there are now some less expensive filter-type flame photom-
C:2f INSTRUMENTS 69

eters on the market that may be entirely adequate for sodium and 

potassium determinations.

Before making a final decision on the type of flame equipment to 

be purchased, it is essential to obtain information on the type of 

fuel permitted by fire-control regulations. Acetylene and hydrogen 

tanks are not allowed in some places.

Attention should be given to the quality of the auxiliary equipment 

purchased for use with the flame photometers. For example, it is 

pointless to purchase the best instrument available and use it with 

an inefficient, erratic air compressor, for example. If an air compressor is necessary, the unit should have more than adequate capacity so that a constant pressure is maintained. The manufacturers 

will supply information as to preferred air compressors to go with 

their units. Filtering of the air may be very important in some 

areas; some instruments come equipped with a trap for this purpose. 

Additional purification and control can be achieved by installing 

ahead of the trap an air-purifier cartridge that is especially effective 

for removing suspended oil droplets from the air. Purification is 

usually no problem with instruments that require only bottled 

oxygen, because the bottled gas is generally sufficiently pure for 

direct use.

Flame photometers in general are less stable than spectrophotometers because the light source in the flame photometer is much 

less stable. The only means of stabilizing this component is by careful control of the flame adjustments. This practice, of course, is not 

completely successful. To avoid introducing still further fluctuation, 

the most careful attention should be given to voltage regulation and 

the photocells, which are most likely to contribute electronic instability. Considerable change in calibration curves is not uncommon after 

the substitution of new phototubes.

C:2f WATER-COLOR COMPARATORS

Color comparators are used to measure the color of the water in 

terms of the color of platinum-cobalt solutions of known concentration. However, colored glass discs individually calibrated to correspond with the colors on the platinum scale have generally replaced 

solutions in the laboratory, and their use is recognized as standard 

practice.

A rather long optical depth of solution is desirable for making 

a comparison because the observed intensity of the color is proportional to the depth of the solution. Satisfactory comparison of the 

water sample and standard can be made if long equal-length cylinders 

of sample and distilled water are illuminated from the bottom and 

viewed from the top of a lightproof housing. The more refined
70 ANALYSIS OF WATER SAMPLES C:2g

instruments bring the beams of light to focus on a single split field 

in the eyepiece.

Commercial color comparators used in laboratories of the Geological Survey must permit color differentiation to the degree listed 

for reporting color units (see sec. D:13a-l) and must give results 

comparable to those obtained with platinum-cobalt standards.

C:2g TURBIDIMETERS

Newell (1902) denned the standard for turbidity as

... a water which contains 100 parts of silica per million in such a slate of 

fineness that a bright platinum wire 1 millimeter in diameter can just be seen 

when the center of the wire is 100 millimeters below the surface of the water 

and the eye of the observer is 1.2 meters above the wire, the observation being 

made in the middle of the day, in the open air, but not in sunlight, and in a 

vessel so large that the sides do not shut out the light so as to influence the 

results. The turbidity of such water shall be 100.

A turbidity rod calibrated for use under these conditions was designed by the Geological Survey and used for many years. However, 

because of the large number of variables that must be controlled to 

measure accurately the turbidity of a water body or sample, a simpler 

method was needed to measure the same characteristic as denned 

by the standard for turbidity. The Jackson candle turbidimeter 

made it possible to determine turbidity in the laboratory and simplified the measurement somewhat. Since shortly after the turn of the 

century the Jackson turbidi meter has been considered by most water 

chemists to be the referee instrument for the turbidity determination, 

although uniform results still required control of several variables. 

More recently, other instruments simpler in operation and more sensitive to small differences in turbidity have been marketed. The advertisement for some of these instruments has stated that the results 

are comparable to the Jackson turbidimeter. Experimentation has 

shown that the nature of the suspended material has a bearing on 

the degree of similarity of the results.

The Geological Survey now uses the Hellige turbidimeter as a 

standard turbidity-measuring instrument, not primarily because of 

the instrument's superiority but because of its sensitivity for measurements of low turbidity and because of the need for results that 

are reproducible and comparable between the different laboratories. 

Operation of the Hellige turbidimeter is based on the comparison 

of the intensity of a beam of light passing through the solution with 

the Tyndall effect produced from the lateral illumination of the 

sample by the same light source.

The slit width that controls the amount of light passing up through 

the solution is controlled by a knob on the side of the instrument. 

This knob is calibrated in arbitrary units, and the readings from the
C:2h INSTRUMENTS 71

knob are converted to parts per million turbidity by means of 

appropriate curves, which are supplied by the manufacturer. The 

transmitted light is viewed as a circle of light in a field of Tyndall 

light. (See sec. D:40a-l.)

C:2H PHOTOMETRIC TITRATION ASSEMBLIES

Interest in photometric titration devices in analytical chemistry 

is increasing. This is another manifestation of the accelerating 

movement toward eliminating subjective factors in analysis. Automatic titrators are used extensively by industry, but they have not 

proved satisfactory in the low-concentration levels characteristic of 

water analysis, and they are quite expensive. For these reasons, 

spectrotitrators standard spectrophotometers or filter photometers 

with titration adaptors are used. Spectrotitrators also probably 

permit a wider range of application than the automatic devices because the latter are based on potentiometric principles and can be 

utilized only when the analytical system gives the required potential 

difference. Probably there are more colorimetric systems of analytical value than there are potentiometric systems.

Two spectrotitration assemblies have been developed by the Geological Survey: a research type for the Beckman B spectrophotometer and a routine-analysis type for the Photovolt Lumetron 401. 

Both are shown on figure 14. The principal application of the latter 

instrument is in the titration of calcium with Na2EDTA (disodium 

dihydrogen ethylenediamine tetraacetate) where murexide is used 

for the indicator. The murexide end point, which is difficult for 

many analysts to detect visually, gives a very decisive response in 

the spectrotitrator. Spectrotitration in this instrument differs from 

visual titration technique mainly in the use of a concentrated titrant 

solution. This is necessary to minimize the volume change during 

the titration. A microburet is used. In general, the monochromaticity requirements for spectrotitrations are considerably less than 

for colorimetric measurements. Filter photometers can almost always be used.

One other type of instrumental titrator is the relatively new highfrequency titration system which requires no electrodes. The principle involves loading the tank circuit of an oscillator with the titration solution. Usually the oscillator coil surrounds the titration 

beaker. As the ionic composition of the solution varies during the 

titration, the gain of the tank circuit varies correspondingly, and in 

the end-point region a major shift takes place which is observed 

on a meter. This method of detecting the end point would seem to 

present the most general application. Although development in this 

field will probably be rapid, the present instruments do not provide 

the sensitivity necessary for general application in water analysis.
72 ANALYSIS OF WATER SAMPLES C:3

C:3

FIGURE 14. Photometric titration assemblies.

GLASSWARE AND OTHER CONTAINERS USED IN THE 

LABORATORY

Laboratory vessels serve three functions: storage of reagents, 

measurement of solution volumes, and confinement of reactions. 

Soft-glass containers are usually relatively soluble and therefore are 

not recommended for storage of many reagents. Even Pyrex, which 

is a borosilicate glass, is not completely inert, particularly to alkalis. 

The Corning Glass Co. Alkali-Resistant (Boron-Free) laboratory 

glassware and T. C. Wheaten and Co. No-Sol-Vit bottles are reported 

by the manufacturers to be more chemically durable than soft-glass 

containers. Standard solutions of silica, boron, and the alkali metals 

are usually stored in polyethylene bottles. Some strong solvents and 

strong mineral acids (such as H2SO4 ) will attack polyethylene
C:3 GLASSWARE AND OTHER CONTAINERS 73

readily. In the instructions for preparations of solutions referred 

to in this manual it is assumed that the solutions will be stored in 

Pyrex bottles unless otherwise specified.

Vessels for measuring solutions differ in their accuracy of calibration. By common usage, accurately calibrated glassware for precise 

measurements of volume has become known as volumetric glassware. 

This group encompasses volumetric flasks, volumetric pipets, and 

accurately calibrated burets. Volumetric glassware used in laboratories of the Geological Survey is calibrated to contain or deliver 

volumes accurate to within the tolerances specified by the National 

Bureau of Standards (1941). These specifications are given in the 

table on the following page. Glassware certified by manufacturers to 

meet specifications of the National Bureau of Standards has been 

found to be satisfactory. Less accurate types of glassware include 

serological pipets, graduated cylinders, and Nessler tubes.

Silicone solutions have appeared on the market that are advertised 

to coat the glass and thereby prevent wetting of the glass by aqueous 

solutions. These products have not been approved by the National 

Bureau of Standards (1941) and are not recommended for laboratory use of volumetric glassware for several reasons. Volumetric 

flasks and pipets are calculated to contain and deliver the correct 

volume when the bottom of the normal meniscus coincides with the 

calibration line; coating of the glass distorts the meniscus. Furthermore, experience has shown that deterioration of the coating causes 

excessive adherence of the solution to the side of the vessel and inaccurate delivery. Once the coating material is on the glass it is very 

difficult to remove.

Pyrex containers are more universally used than any other as containers for reactions because of their relative inertness and thermal 

properties. Nessler tubes and absorption cells (spectrophotometric) 

are often used for colorimetric reactions to eliminate the necessity 

of transferring solutions before determining the color intensity.

Treatment of the water sample or reaction of the solution that 

involves evaporation can usually be carried out in glass, porcelain, or 

platinum dishes. Platinum is superior to the other materials if the 

weight of the residue must be determined accurately because the 

weight of platinum vessels is very constant.

Glassware and porcelain can be cleaned with soap, synthetic detergents, organic solvents, dichromate cleaning solution or aqua regia 

(25 percent v/v cone HNO3 in cone HC1). Aqua regia will attack 

platinum, but 50 percent hydrochloric acid is an effective cleaning 

agent. Material that does not dissolve in 50 percent hydrochloric 

acid usually can be removed by rubbing with fine sea sand.
74 ANALYSIS OF WATER SAMPLES C:3

TABLE 1. Tolerances for volumetric glassware 

[Abridged from National Bureau of Standards, 1941]

Capacity (ml) 

less than and 

including 

Limit of error (ml)

Graduated flasks

25

50

100

200

300

500

1,000

2,000

0. 03

. 05

. 08

. 10

. 12

. 15

. 30

. 50

Transfer pipets

2

5

10

30

50

100

200

0. 006

. 01

.02

.03

.05

. 08

. 10

Burets >

2

5

10

30

50

100

0.01

.02

.03

.05

. 10

1 Limits of error are of total or partial capacity. Customary practice is to test the capacity of 5 intervals

Although platinum is one of the most resistant metals, it is not 

completely inert and is subject to embrittlement. The following 

precautions are recommended: Never put solutions containing tin, 

mercury, or lead in a reducing environment in platinum; if the free 

metal should be formed it will alloy with the platinum, especially if 

heated. Do not heat mixtures of hydrochloric acid with oxidizing 

substances such as nitrate or manganese dioxide; ferric chloride in 

hydrochloric acid attacks platinum appreciably. Place hot platinum 

vessels on a refractory material; never on a cold metal surface or on 

a dirty surface. Use clean platinum-tipped tongs to handle hot
C:5 TERMINOLOGY 75

platinum vessels. Coarse crystal growth and embrittlement can be 

caused by unnecessarily prolonged heating at high temperatures, 

heating under reducing conditions, and heating phosphates or sulfates in the presence of organic compounds. Embrittlement can be 

counteracted by rubbing the platinum ware with moistened sea 

sand. Gentle rubbing with sea sand cold-works the metal and breaks 

down the crystal structure. Detailed instructions for the care and 

use of platinum ware are distributed by manufacturers of these 

vessels and are described in textbooks of quantitative analysis.

0:4 CHOICE OF ANALYTICAL METHODS

Several analytical methods may be listed for the determination of 

a single constituent in the section of this manual that deals with 

methods of analysis. The laboratory methods are generally the more 

accurate, but useful approximations and screening tests are also included. Screening tests may be either a means of approximating 

concentrations for the selection of the proper sample volume for a 

laboratory determination or they may be tests that are sufficiently 

accurate in only a low-concentration range. For example, the tetramethyldiaminodiphenylmethane test for manganese is a screening 

test that is rapid and accurate in low-concentration ranges; if the 

concentration of manganese is within the range of this test, it is not 

then necessary to perform the more laborious regular procedure.

A choice of laboratory methods is also given for the determination 

of some constituents to provide the analyst with the means for checking dubious results and to permit a flexible laboratory scheme. Most 

procedures are subject to some interferences, but seldom are two 

methods subject to the same interferences. No method is designated 

as a referee method in which the results obtained are the ultimate 

in accuracy or by which the adequacy of other methods can be gaged. 

However, the method listed first in each group of methods is generally 

preferred because of its applicability to most water samples, its 

advantages in required time, its accuracy, or its comparative simplicity.

C:5 EXPLANATION OF TERMINOLOGY

Specific terminology is used in describing the analytical techniques. 

Some that are an integral part of the instructions are described 

below.

Dilution water for adjustment of sample volume or preparation 

of reagent solutions, unless otherwise specified, shall be distilled water 

whose pH is between 5.8 and 7.2 and whose impurity as measured, by 

specific conductance at 25 °C shall not exceed 8.0 micromhos.

Redistilled water shall be prepared by redistillation of dilution 

water from Pyrex apparatus. It shall have a specific conductance
76 ANALYSIS OF WATER SAMPLES C:5

at 25°C of 1.5 micromhos or less. Redistilled water should be stored 

in resistant glass (Corning Glass Co. Alkali-Resistant glass No. 7280) 

or polyethylene bottles and should be prepared fresh frequently.

Metal-free water shall be distilled water that has been passed 

through a suitable acid-charged cation-exchange resin and that conforms to the pH and other requirements of dilution water.

Carbon dioxide-free water shall be prepared by boiling and cooling dilution water immediately before use. It shall have a pH between 6.2 and 7.2.

Ammonia-free water shall be distilled water that has been freed of 

ammonia by shaking with Folin's ammonia permutit and that conforms to the pH and other requirements of dilution water.

In the instructions for making the analysis and preparing the 

solutions, significant figures are utilized to define the accuracy of 

weights and measures. Weighings will be accurate to the last figure 

shown. For example: A mass designated as 4.532 g must be 

weighed accurately to ±0.0005 g, whereas a mass designated as 

4.5 g must be weighed accurately to only ±0.05 g.

Required accuracy for measurement of volume in the analysis 

and preparation of reagents is shown similarly. Standard solutions 

are always prepared in and measured from volumetric glassware. 

The significant figures given for such measurements are in practical 

agreement with the tolerance limits (see sec. C:3), and the word 

"approximately" indicates that the volume or weight needs to be 

within only ±5 percent of that prescribed. For example: "Add 

2.00 ml of reagent" shows that a volumetric pipet must be used for 

the addition, but "add 2 ml" or "add 1.5 ml" shows that a serological 

pipet may be used; "dilute to 1,000 ml" shows that a volumetric flask 

is essential, but "dilute to approx 1 liter" permits the volume to be 

between 950 and 1,050 ml.

The chemistry of most water-analysis procedures is such that 

they are applicable over a restricted concentration range. To 

get a satisfactory amount of the constituent into the reaction and 

still maintain the correct proportions of sample and reagents, the 

analyst may have to concentrate a large volume of sample or dilute 

a small sample to the proper volume. Concentration limits of the 

procedures are given in the instructions to aid the analyst in the 

selection of the proper sample volume. The highest accuracy in 

most determinations is obtained if the largest convenient sample 

volume is selected, because the factor for converting the weight of the 

constituent in the test sample to parts per million is smaller if a 

larger volume of sample is used. Most of the analytical methods
Cl5 TERMINOLOGY 77

are designed for dilute solutions, and, therefore, concentration by 

evaporation is rarely required. Evaporation should be kept to a 

minimum because of the danger of contamination during the long 

exposure to the atmosphere and because sparingly soluble salts may 

precipitate.

Obviously, selection of a test-sample volume can be carried to 

extremes for the sake of ..dubious accuracy. Therefore, parenthetical . 

quantities are used to designate the accuracy normally desired. For 

example: "Measure a sample containing less than 0.5 mg SiO2 

(10.00 ml max)" indicates that a 10-ml sample will give the desired 

accuracy in the 0.00- to 0.50-mg range and that a larger sample is 

unnecessary.

All test samples and standard solutions are measured with volumetric glassware (see sec. C:3). Test-sample volumes less than 

5 ml should not be measured directly because the calibration of 1-ml 

and 2-ml pipets is not as precise as that of the larger volume pipets. 

Less error is incurred if a suitable sample dilution is prepared and 

part of this dilution taken for the test sample. Although the glassware is calibrated to deliver a specific volume at 20°C, the error 

in measurement incurred by pipeting samples at room temperature 

is insignificant for water analysis. One gram of pure water (H2O) 

is contained in 1.002 ml at 20°C and in 1.007 ml at 38°C; the maximum error in volume that will result from temperature differences is 

only 0.5 percent. Brine samples unstable at room temperature cannot 

be measured in the usual way. The brine may be heated to bring 

precipitated material into solution and the test measured approximately into a tared weighing bottle and weighed accurately. Calibrated pycnometers used with a constant-temperature bath are 

preferable for measuring test-sample volumes of heavy stable brines 

when highest accuracy is required; these bottles, normally calibrated 

at 20°C, should not be used with brines that are unstable at this 

temperature.

Chemicals used in standard and reagent solutions shall conform 

with the specifications for purity of the American Chemical Society 

(1950 and subsequent revisions) when such specifications exist. 

Chemicals not contained in this compilation shall be of commonly 

accepted reagent purity. Primary standards (solid) for acidimetry, 

alkalinity, and oxidation-reduction reactions may be obtained from 

the National Bureau of Standards or from chemical manufacturers 

marketing products of comparable purity. Primary standard solutions are prepared fresh in duplicate for each standardization of 

volumetric solutions.
78 ANALYSIS OF WATER SAMPLES C:6

C:6 SCHEME OF ANALYSIS

The validity of the analytical statement and the time required for 

analysis are affected by the order in which constituents are determined and by the manner in which the sample is treated. The following discussion concerns only the group of constituents and measurements included in the "complete analysis" normally made by most 

of the laboratories of the Geological Survey and does not include 

the determinations of volatile materials or constituents readily 

changed by aeration or bacterial action.

Constituents and measurements most susceptible to change are 

determined first. The pH and relative proportions of the alkalinity 

and acidity components are related to the gas content of the sample 

and are the least stable. Filtration may alter the gas content of the 

sample. Hence, pH and alkalinity or acidity of the unfiltered sample 

should be determined immediately after the bottle is uncapped. 

Composite samples should not be prepared far in advance of analysis.

The specific conductance gives an indication of the total ionic concentration of the sample and is, therefore, useful in the selection 

of the volume of sample to be taken for the individual determinations. The specific conductance is not affected by moderate turbidity 

and can be determined from either a raw or filtered sample.

If the sample requires clarification (see sec. B:5), a sufficient 

volume should be filtered at one time to provide samples for all remaining determinations. This has two advantages: less time is 

required for one filtration than for several individual filtrations, 

and if any solution or exchange of ions is progressing at the sedimentwater interface, an early separation is desirable. The single filtrate 

provides identical water for analysis, whereas there is no assurance 

that individual filtrations at different times would provide identical 

samples.

A loss of carbon dioxide from solution converts bicarbonate to 

carbonate. Calcium carbonate, which is only sparingly soluble, often 

precipitates soon after a bottle is uncapped. Precipitation of 

calcium carbonate affects the validity of the pH, alkalinity, specificconductance, calcium, titrated-hardness, and dissolved-solid values. 

Consequently, test-sample volumes for determination of calcium, 

hardness, and dissolved solids should be withdrawn soon after the 

bottle is opened. All samples should be inspected visually for 

evidences of calcium carbonate precipitation when received in the 

laboratory. If precipitation is evident, the analytical statement will 

not represent the chemical character of the sample at the time of 

collection unless special techniques are used. Proper handling of 

such samples requires special treatments and may involve the
C:7 SIGNIFICANT FIGURES 79

omission of some determination from the analysis. Selected constituents and measurements, even calcium and total alkalinity, can 

be determined with a reasonable degree of accuracy. The selected 

special treatment will depend on the nature of the precipitate and 

the constituents that are considered most important in the analysis. 

Other constituents and measurements can be determined in any 

order. Time is important, however, and the analysis should be 

carried to completion without undue delay. It has been observed 

that, because of changes in chemical composition of the water, 

the validity of the analytical statement often decreases as the £ime 

consumed in the analysis increases. All required determinations 

should be made within 1 week after the sample is opened.

C:7 SIGNIFICANT FIGUEES FOE EEPOETING RESULTS

The significant figures used by the Geological Survey in reporting 

the results of analysis are the result of a compromise between 

precision of the measurement, importance of this precision to the 

use of the analytical data, and the obvious advantages of some 

semblance of uniformity in tables of analyses.

One of the commonly used methods, which applies only to the 

expression of the precision of a determination, is to include all 

certain digits and the first (and only the first) doubtful digit. This 

method has one obvious disadvantage: Published data so reported 

may not be interpreted to mean the same thing by all users of the 

data. Possible confusion and misinterpretation is minimized if 

the tabulated analytical results represent the certainty of the reported 

results within reasonable limits. In analytical results published 

by the Survey, the last digit may not be absolutely accurate in all 

analyses, but it is sc figure that can be used with a good degree of 

confidence in most applications of water-analysis data.

The precision of procedures for water analysis cannot be measured in absolute terms of either concentration or percentage, nor is 

it desirable to do so for practicable purposes. Generally, more 

precise measurements are required when the concentration of a constituent is low than when the concentration is high. For example, 

in potable water supplies the difference between 1.0 and 1.5 ppm 

of fluoride is very important, but the difference between 10.0 and 

10.5 ppm is not significant, although the absolute error is the same. 

Neither is percentage of error a good gage of the precision of the 

analysis. For example, an absolute error of ±0.01 ppm is an error 

of 100 percent when the concentration of the constituent is only 

0.01 ppm.

Another factor that complicates the problem of designating 

precision of analysis is the inherent tendency of most analytical
80 ANALYSIS OF WATER SAMPLES C:7

determinations to vary in accuracy throughout the concentration 

range ~ of the constituent in a manner that is akin neither to absolute

nor percentage quantities. Consequently, the Survey's rules for 

rounding off and reporting significant figures are somewhat arbitrary. 

However, the individual procedures contain a general statement as 

to the precision of the method in the concentration range most frequently found in water analysis in routine laboratory operations.

The importance of the accurate determination of minute quantities 

of a constituent to the usability of the water has also influenced the 

selection of significant figures. Water-quality data are used for a 

vast variety of purposes. Water analysts have endeavored to 

design and use analytical methods sufficiently sensitive to meet the 

requirements of exacting users of water-quality information. The 

precision of a few methods and their ability to measure low concentrations accurately are far superior to present needs. However, 

when such precision and sensitivity are available with no additional 

effort on the part of the analyst, the results are determined and 

reported to the smallest concentration increment that can be obtained.

Instructions for reporting analytical results are given in "Analytical procedures" (sec. I)) for each constituent and property.

Chemical equivalents per million are computed by multiplying 

the reported concentration of the individual constituents in parts 

per million, by the reciprocal of their combining weights. The 

reciprocal factors,of the more commonly determined constituents are 

given in table 2. Equivalents per million (epm) as reported by 

the Geological Survey are numerical expressions of parts per million 

and for uniformity are carried to two decimal places regardless of 

the magnitude of the parts-per-million value; the significant figures 

shown in no way reflect the accuracy and precision of the measurement as do the parts-per-million values.

TABLE 2. Factors for converting parts per million to equivalents per million

[1954 atomic weights]

Sum of 

Ion atomic weights

Al+3 _ __________ 26.98

Br-'__________._ 79.916

Ca+2 ___ ________ 40.08

Gl-'_-_._______ 35.457

CO_-2 __________ 60.011

F~l ____________ 19.00

Fe+2 ___________ 55.85

Fe+3 ___________ 55.85

H+>____________ 1.0080

HCO3 - 1 -----___ 61.019

I- 1 --------.--. 126. 91

Multiply 

ppm by 

0. 11119

. 01251

. 04990

. 02820

. 03333

. 05263

. 03581

. 05372

. 99206

. 01639

. 00788

Ion

K+l ____-_.__

Li+'________

Mg+2 _ ______

Mn+2 _ ______

Mn+4 _ ______

Na+1 ___--_._.

NOs-i-_-_-_

OH-i_______.

S~2 _________.

SOr2 ------.-

Zn+2 ________

Sum of 

atomic weights

... 39. 100

6. 940

. _ 24. 32

. _ _ 54. 94

54. 94

. _ 22. 991

. _ _ 62. 008

__ 17.0080

_ _ 32. 066

_ _ 96. 066

_ . 65. 38

Multiply 

ppm by 

0. 02558

. 14409

. 08224

. 03640

. 07281

. 04350

. 01613

. 05880

. 06237

. 02082

. 03059
C:8a ACCURACY OF RESULTS 81

C:8 EVALUATION OF THE ACCURACY OF ANALYTICAL RESULTS

Some errors are practically unavoidable in analytical work. Errors 

may result from the reagents, from the limitations of the method 

or instruments employed, or even from impurities in distilled water. 

The analyst's skill and general judgment have a direct bearing 

on the accuracy of the analytical statement. After the chemical 

analysis of the water sample has been completed, there are several 

ways by which the validity of the results can be evaluated. No one 

method of checking gives conclusive proof of the accuracy of the 

determinations, but the process of checking may bring to light some 

dubious results or may suggest some additional constituents of the 

sample that were not considered in the analysis.

C:8a CHEMICAL, BALANCE

One of the most commonly used procedures for checking water 

analyses is a balancing of the chemical equivalents of the major ions. 

Since water is a chemically balanced system, the sum of the equivalents of cations in solution equals the sum of the anions. If all of 

the predominant ions have been determined, the equivalents per 

million should be in reasonable balance.

The hydrogen-ion content of acid water is included in the balance. 

The hydrogen-ion concentration is approximated from the pH of the 

sample,

pH = log -

or

where the base of hydrogen is the effective concentration (activity) 

of hydrogen ions. The calculated pH of standard solutions of sulf uric acid has been compared with determined pH, and the agreement 

is good up to 2.0 ppm H (pH 2.70). Reproducibility and accuracy 

of ±0.1 epin H+1 is the best that can be anticipated under normal 

operating conditions and with most waters. Considerable possible 

error is introduced in converting pH to hydrogen-ion concentration 

because of the effect of other ions on the activity of the hydrogen ion.

<pH 

4.25-3.85

3.80-3.60

3.55-3.50

3.45-3.40

3.35-3.30

3.25-3.20 _

3.15

3.10

3.05

3.00

epm H+1 

0.1

.2

.3

.4

.5

.6

.7

.8

.9

1.0

2.95

2.90

2.85

2.80

2.75

2.70

2 65

2.60

2.50

<pH epm H+1 

1.1

1.3

1.4

1.6

1.8

2.0

2.2

2.5

2.8

3.2
82 ANALYSIS OF WATER SAMPLES C:8a

pH meters determine the activity of the hydrogen ion as distinguished 

from concentration. The chemist is referred to the work of Kolthoff 

and Laitinen (1941) for a full discussion of the subject.

Multivalent ions present difficulties in the ionic balance unless the 

ionic states are differentiated by the analysis. Fortunately, these ions 

are present in only minor amounts in many waters and rarely influence 

the ionic balance appreciably. Orthophosphates may occur in water 

as PO4~3 , HPO4-2 , and HaPO"1 ; the proportion of each of these ions is 

related to the pH of the water. The PO4~3 ion occurs only above about 

pH 9.3. Because natural waters practically never attain this pH, a 

general assumption can be made that the PO4~3 ion is not present in 

natural waters. For the purpose of ionic balance, the proportion of 

HPO4-2 and H2PO4~1 present can be calculated from the pH relation 

as shown in table 3. The data in table 3 are based on work of Sorensen 

as reported by Clark (1928) and from Rieman, Neuss, and Naiman 

(1942, p. 323). Review and checking of this table with work by 

others is in order. However, limited use of this table has provided at 

least reasonably good analytical balance for samples with a high orthophosphate content. Other forms of phosphorus cannot be differentiated in this manner, nor are suitable methods available for the other 

multivalent ions.

TABLE 3. Orthophosphate mixtures as related to pH

pH

4.5

4.6__ _ 

4.7

4.8____ __________

4.9_ 

5.0-__ ___________

5.1 -- ----

5.2 

5.3 

5.4-. ----------

5.5 - - - -

5.6 

5.7 

5.8_ _____ -----

5.9_- _ -

6.0-_-__-_-_-_---

6.1 

6.2____ __-__ _

6.3

6.4____ __ ___ _

6.5 

6.6_ _

6.7 

6.8 

6.9__-__-_____--.

HPO4

(percent)

 __ 0

--__ .5

_-._ .5

__-_ 1.0

 1.0

_--_ 1.5

____ 1.5

2 0

2 5

-__ _ 3.0

 _ 3.5

 _ 4.5

_.-_ 6.0

_ 8.0

____ 9.5

____ 12. 5

____ 15. 5

___- 19. 0

___. 23. 0

 -_ 27. 0

____ 32. 0

_ 37.5

_ _ . _ 43. 5

_ . _ _ 49. 5

_ _ _ _ 55. 5

H2PO4

(percent)

100. 0

QQ ^

99. 5

QQ 0

QQ 0

98. 5

98. 5

98. 0

Q7 ^

Q7 0

96. 5

95. 5

94 0

no n

on K

87.5

04. c

81. 0

77. 0

73. 0

68. 0

62. 5

e« c

en ^

44. 5

pH

7.0___ _ -

7.1__ --------

7.2__ --------

7.3-_ --------

7.4._--------

7 ^

7.6-- ---------

7.7__- ___ --.

7.8__ ---------

7 Q

8.0__ _--- - 

8.1.- --------

8.2-_-_--------

8

0

8 4

8.5------------

8.6__- _---_--_--

« 7

8.8- _ ---------

8.9__-__-_-------

9 0

9 1

Q 9

9.3--__ ----------

HPO4

(percent

-_.- 61. 0

___. 66. 5

__-_ 72. 0

____ 76. 5

.___ 80.5

__._ 84. 0

.___ 87. 0

__-. 89.5

__-- 91.5

__. 93. 0

____ 94. 5

--__ 95. 5

._-_ 96. 5

____ 97. 0

_. 98. 0

___- 98.5

_-_. 98. 5

_._- 99. 0

__- 99. 0

-___ 99. 0

__-_ 99. 5

,___ 99. 5

-.__ 99. 5

-.--100. 0

H2 PO4 

(percent)

39. 0

33.5

28. 0

23. 5

19.5

16. 0

13. 0

10. 5

8. 5

7. 0

5. 5

4. 5

3.5

3. 0

2.0

1. 5

1.5

1.0

1. 0

1. 0

. 5

. 5

. 5

. 0
C:8b ACCURACY OF RESULTS 83

Dissociation also must be considered in balancing analyses. Many 

of the determinations, particularly those for the heavy metals, do 

not differentiate between dissociated and undissociated constituents. 

Constituents that hydrolyze to give undissociated products determined along with the ionized forms in the analysis cannot be included directly in the ionic balance. Published dissociation constants give some indication of the possible ionized concentration, but 

complete confidence cannot be placed in these values when considering complex solutions such as natural water.

The deviations from balance can be expressed in terms of absolute 

quantities or as a percentage of the total ionic concentration. Obviously, the analyst must use some type of sliding scale to evaluate 

the significance of the deviations for water of different concentrations. With careful work and a comprehensive analysis of the 

sample, it is customary for the deviation between equivalents per 

million of cations and anions not to exceed 1 or 2 percent of the 

total concentration for analyses of waters with more than about 

150 ppm of dissolved solids. A deviation of up to 3 percent may 

result from the accumulation of small unavoidable errors in the 

analysis of water containing about 100 ppm of dissolved solids and 

about 5 percent for waters with as little as 25 or 30 ppm.

Chemical balance is only an indication of the gross validity of 

the analysis. Very large errors in the determination of minor constituents can go unnoticed; neither are compensating errors detected. 

Conclusions that can be drawn from deviations in anion and cation 

balance are usually negative. Large deviations indicate either a 

large error in one or more determinations or the presence of some 

undetermined constituent, but a good balance is not conclusive evidence that each of the determinations is accurate nor that all constituents have been determined. Chemical balance is one tool for 

evaluating the validity and comprehensiveness of an analysis but 

must not be a goal for the analyst.

C:8to RELATION OF RESIDUE ON EVAPORATION TO CALCULATED 

DISSOLVED SOLIDS

Comparison of the residue on evaporation and dissolved solids 

calculated from the analytical statement is a rough check of the 

comprehensiveness of an analysis (see sec. D:36a-2). Several important factors have a bearing on this comparison. First, the residue on evaporation might contain appreciable amounts of organic 

and some inorganic materials that are not determined in the 

analysis. Second, water of hydration may also be contained in the 

residue. Ignition loss is sometimes determined as a rough estimate 

of the water of hydration and combustible organic materials. The
84 ANALYSIS OF WATER SAMPLES Cl8c

calculated dissolved solids is usually between the value obtained as 

residue on evaporation and the residue on evaporation minus the 

ignition loss. Third, volatile solids which are determined in the 

course of the analysis may be lost during evaporation. If volatile 

solids are lost, the residue on evaporation may be less than the calculated value. Fourth, the alkalinity determination is influenced by 

weak-acids residuals other than carbonate and bicarbonate. Phosphate, borate, and silicate, measured collectively as alkalinity and 

reported as carbonate and bicarbonate, may be redetermined individually in specific analyses. If so, the solids determined by calculation may be higher than by evaporation. Other factors such as 

the behavior of acid waters and iron compounds during evaporation 

also complicate the relation between residue on evaporation and the 

calculated dissolved solids.

C:8c SPECIFIC-CONDUCTANCE RELATION

For most natural waters of mixed type the specific conductance, 

in micromhos, multiplied by a factor of 0.65±0.1 approximates 

the residue on evaporation in parts per million. This does not 

approach an exact relation because the conductance of a solution is 

dependent on the type and total quantity of ions in solution. More 

precise relations can be developed for specific water types. The 

factor of 0.65 is applicable only with comparatively dilute solutions 

and usually increases as the total dissolved-salt content exceeds 2,000 

to 3,000 ppm. For waters that contain appreciable concentrations 

of free acid, caustic alkalinity, or sodium chloride, the factor may 

be much less than 0.65. The factor for some other specific types 

of water may be higher. Nonionized silica will also disturb the 

ratio of residue on evaporation to conductance. With similar limitations, the specific conductance divided by 100 approximates the 

equivalents per million of anions or cations. This relation is particularly helpful in detecting the location of error (in anions or 

cations) as well as for estimating the comprehensiveness of an 

analysis.

Rossum (1949, p. 631) proposed a method for checking analyses 

that is based on a comparison of the specific conductance of dilute solutions with the summation of the added increments of conductance contributed by each determined ion in solution. This procedure is known 

as the diluted-conductance method. The sample is diluted with redistilled water until the conductance of the solution is between 90 

and 120 micromhos. This observed conductance is recorded. The 

exact dilution ratio, Z>, is then computed as follows :

j^_ volume of sample + volume of redistilled water 

volume of sample
C:8d ACCURACY OF RESULTS 85

Next, the conductance of the redistilled water, Kw, is determined and 

the true diluted conductance, A'd, is calculated:

/Vd = observed diluted conductance (D-\)KW

The diluted conductance is then compared with the summation of the 

conductances of the determined ions.

Dilvted-conductance factors for ions commonly found in water 

[After Rossum, 1949]

Ion

Sulfate. _ ________---____---_-_-__-____-

Micromhos 

per meq per 

liter at %5° C.

.._______-. 43.6

___________ 52.0

__________ 84.6

__________ 75.9

__________ 46.6

__________ 71.0

__________ 72. 0

__________ 48. 9

__________ 73.9

Micromhos 

per mg per 

liter at %5° C.

0. 715

2. 60

2. 82

2. 14

3. 82

1. 15

1. 84

2. 13

1. 54

The American Public Health Association and others (1955) 

recommends that a recheck of the chemical analysis is advisable if 

the diluted conductance differs by more than 2 percent from the 

sum of the conductances of the determined ions. This degree of 

accuracy may be somewhat excessive for routine water analysis 

unless the dilutions are made with extreme care.

The diluted-conductance method of checking is not applicable to 

waters whose initial conductance are less than 90 micromhos, or to 

samples whose pH values are not between 6 and 9, or to samples 

that contain appreciable quantities of ions not listed in the above 

table. The conductance contributed by hydrogen and hydroxyl 

ions is greater than that by other common ions.

This method of checking is a little more time consuming than 

chemical balance, but it does have the advantage of considering all 

of the ions in solution. A good check by the diluted-conductance 

method is conclusive proof that no major ion. has been omitted from 

the determination.

C:8d INTERRELATION OF CONSTITUENTS OF WATER FROM THE

SAME SOURCE

As samples from the same source repeatedly come through the 

laboratory, the analyst begins to recognize patterns of water quality. These may be the relation of calcium to bicarbonate, sodium 

to chloride, or other relations that appear to hold true. When the 

determined values deviate significantly from previously observed 

relations the analyst is immediately suspicious of the analytical 

results and may elect to make some checks. The apparent departure 

from the expected may be the result of manipulative errors, such as
86 ANALYSIS OF WATER SAMPLES C:8e

recording ir correct test-sample volume, buret readings, or weights 

or transpcing figures. Checks of these and similar points are 

usually made first, followed by redetermination of the questionable 

measurement when indicated, before assuming that the unexpected 

value represents an actual change in water quality. This method 

of evaluating an analysis must be used judiciously because some 

marked deviation from normal relations can be expected occasionally, 

and an attempt to reproduce a relation may in actuality result in 

an erroneous analysis. The most valid analytical value should be 

the one finally accepted and not necessarily the one that most closely 

approaches the analysis of a previous sample.

C:8e METHOD FOR CHECKING PARTIAL ANALYSES

Much has been said here about checking the validity of comprehensive analyses, but little help has been given for verifying results 

when only a few of the constituents are determined. There is no 

simple adequate method for checking a partial analysis. A recovery 

test can be run by adding a known quantity of constituent to the 

sample and repeating the analysis. The results are not always conclusive, however; a fraction of the constituent that missed detection 

in the first analysis might also be missed in the second. Interrelations of constituents in water from the same source and relations of 

certain constituents to specific conductance or water discharge may 

be helpful as guides, but the ability of the analyst is the only good 

control.

REFERENCES

American Chemical Society, 1950, Reagent chemicals, American Chemical

Society specifications: Washington, D. C. 

American Public Health Association and others, 1955, Standard methods for the

examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

Clark, W. M., 1928, The determination of hydrogen ions: Baltimore, Md.,

Waverly Press. 

Gibson, K. S., and Balcom, M. M., 1947, Transmission measurements with the

Beckman quartz spectrophotometer: Natl. Bur. Standards Jour. Research,

v. 38. 

Kolthoff, I. M., and Laitinen, H. A., 1941, pH and electro titrations: New York,

John Wiley and Sons, Inc. 

Mellon. M. G., 1950, Analytical absorption si>ectroscop.v: New York, John Wiley

and Sons, Inc. 

National Bureau of Standards. 1941, Testing of glass volumetric apparatus:

Circ. C434.

 1947, Letter Circ. LC-857. 

Newell, F. H.. 1902, Measurement of turbidity and color: U.S. Geol. Survey Div.

Hydrography Circ. 8. 

Reiman, W., Neuss, J. D., and Naiman, B., 1942, Quantitative analysis: New

York, McGraw-Hill Book Co. 

Rossum, J. H., 1949, Conductance method for checking accuracy of water

analyses: Anal. Chemistry, v. 21.
SECTION D. ANALYTICAL PROCEDURES

D:l ACIDITY

The acidic properties are attributable to the presence of mineral 

acids, uncombined dissolved gases, organic acids, and salts of strong 

acids and weak bases. Hydrolyzable salts of iron and aluminum 

of mine and industrial origin are common sources of acidity. The 

analyst should differentiate between acidity as a property of a 

solution and as a total concentration of acids. In terms of the 

dissociation theory, an acid is any compound which on dissociation 

produces hydrogen ions. To determine the concentration of a specific acid compound requires titration with a base to a practical 

end point at which all hydrogen ions that can be produced by the 

compound have been neutralized. This end point is normally taken 

as the inflection point, or points, on the titration curve. The inflection points differ with the acid compound, hence it is impossible to 

determine accurately the acid concentration of different hydrolyzable salts or mixtures by titration to a single predetermined end 

point. Total acid compounds might be estimated by titrating to 

the last inflection point on a sodium hydroxide curve, provided that 

this point is readily recognizable. Acidity, on the other hand, is 

the property of a solution attributable to the presence of an excess 

of hydrogen ions over hydroxyl ions, and the acidity value is a 

measure of the strong base required to adjust the hydrogen and 

hydroxyl ions to equivalency. Acid compounds may exist at the 

equivalence point, pH 7.0, without producing hydrogen ions in 

solution.

Three determinations of acidity are given, each of which measures a different group of contributors to acidity. Total acidity (see 

sec. D: la) shows the total stream acidic potential. The immediateacidity (see sec. D:lb) determination measures the existing acidity 

from all causes plus that immediately available from uncombined 

dissolved gases when titrated with a standardized base. Potential 

free acidity (see sec. D: Ic) represents the existing acidity from all 

causes and includes that which may develop from acid-producing 

salts but excludes that from uncombined gases. The difference 

between the total acidity and potential free acidity is an indication 

of the acidity that can be developed from uncombined gases. Similarly, the difference between total acidity and immediate acidity is

87
88 ANALYTICAL PROCEDURES D:la

an estimate of the acidity that can be developed through oxidation 

and (or) hydrolysis of acid-producing salts. Combined, the three 

determinations are useful for approximating the nature of the acidcausing components. Determinations of acidity are among the least 

reliable in water analysis in respect to accuracy and reproducibility 

of results.

The acidity of water is significant to the operation of watertreatment plants and waste-disposal systems. High acidity, particularly that due to strong acids, contributes to the corrosiveness 

of the water.

D:la TOTAL ACIDITY

Total acidity is the capacity of a water containing a compound 

or compounds, with or without hydrolysis, for neutralizing base to 

pH 7.0. Basicity is the counterpart of total acidity.

D:la-l VOLUMETRIC METHOD

Principle of determination

The volumetric determination of total acidity measures the total 

capacity of the water to neutralize base to pH 7.0, irrespective of 

the time involved for the reactions to reach equilibrium at pH 7.0. 

This capacity includes the combined effect of free mineral acids, dissolved gases, buffer systems, and combined equilibrium products of 

hydrolysis at pH 7.0. The acidity is determined by titrating the 

water with a standardized strong base. Often, total acidity is due 

primarily to carbon dioxide or other dissolved gases. Hence, customary precautions should be taken to minimize the loss of these gases 

before they are chemically combined with the tit rant base.

If the water contains principally free mineral acids, the total 

acidity can be titrated rather accurately. If hydrolyzable salts are 

present the titratioii is complicated by the degree and rate of hydrolysis. For example :

The addition of sodium hydroxide shifts the hydrolysis equilibrium 

toward the right, thereby liberating more hydrogen ions. Hydrolysis proceeds slowly in the cold as the basic tit rant is added to the 

solution, and the time involved for the hydrolysis to reach equilibrium 

at pH 7.0 may be infinite. Heating of the sample hastens hydrolysis. 

The sample is titrated rapidly in the cold to pH 7.0 ; the free mineral acids have now been neutralized and all the dissolved gases 

that are going to react at this pH have been combined. Then the 

solution is heated. Boiling is to be avoided because of the excessive 

loss of dissolved gases. The sample is allowed to cool before com-
D: la-1 ACIDITY 89

pleting the titration to avoid upsetting the pH-temperature relations. 

Heating may increase the pH slightly if no hydrolyzable salts are 

present. Such increase in pH is probably due to the loss of uncombined carbon dioxide at pH 7.0 and hence has no bearing on the 

capacity of the water to neutralize base. When pH increases during heating, the titrant volume required by the cold solution is used 

in the computation. If the carbon dioxide concentration of the 

water is high, the end point of the cold titration will not be sharp 

because of the buffering effect of the bicarbonate-carbonic acid 

system.

The procedure given here has been used for the determination of 

total acidity in waters that contain mine-drainage products and some 

industrial pollutants. The procedure may not be universally applicable to all types. The total-acidity determination is one of the least 

reliable in water analysis in respect to accuracy and reproducibility of 

results. If the sample contains no materials acting as buffers at the 

end point, reprodncibility and accuracy of ±0.05 ppm H+1 can be expected. Unfortunately, this condition seldom prevails, and usually 

the reproducibility and accuracy are appreciably less.

Apparatus and reagents

pH meter

Electric hotplate

Sodium hydroxide, 0.0248N, 1.00 mloO.025 mg H+1

Procedure

For precise work the sample for the determination of total acidity 

should be collected as directed in sec. A: 3c and the analysis should 

proceed as soon as possible after the collection of the sample.

1. Pipet a volume of sample containing less than 1.0 mg H+1 (50.0 ml max) into 

into a 150-ml beaker.

2. Rapidly titrate the solution with 0.0248F NaOH (1.00 nilc=0.023 mg H+1 ) to

pH 7.0.

3. Heat the solution to 80°-90C C (do not boil) and maintain this temperature 

for 2 min.

4. Cool the solution to approximately room temperature.

5. Continue the titration to pH 7.0. Disregard any pH increase during the 

heating period.

Calculations

ppm Total acidity (H+')= , * .. X 1>000 X ml titrant X 0.025 density ml sample

Report total-acidity concentrations of <10 ppm to 1 decinml place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.
90 ANALYTICAL PROCEDURES D:lb

Preparation of reagents

Sodium hydroxide, 0.02482V, 1.00 mloO.025 ing H+1 : Dilute 12 nil 2N NaOH with 

carbon dioxide-free water to approx 1 liter. Standardize the solution 

against primary standard potassium acid phthalate as follows: Lightly 

crush 3-4 g of the salt to a fineness of approx 100 mesh and dry for 1 or 2 

hr at 120°C. Dissolve about 2 g, accurately weighed to the nearest milligram, in carbon dioxide-free water and dilute to 500.0 ml. Titrate 50.0 ml 

of the solution with the NaOH to pH 8.6.

v Normality ... , .. gKHC8H4O4 in 50.0 mix4.8967 of alkali =- . n .. 

J ml alkali

Store the 0.02482V base in a tightly capped polyethylene bottle. 

Sodium hydroxide, 22V: Cover approx 100 g NaOH sticks with water until the 

surface coating is dissolved. Discard the supernatant fluid and immediately dissolve the remaining NaOH (about 80 g) in approx 1 liter of carbon 

dioxide-free water. Store the base in a tightly capped polyethylene bottle.

D:lto IMMEDIATE ACIDITY

Immediate acidity is the capacity of a water containing: a compound 

or compounds, with or without hydrolysis, for rapidly neutralizing: 

base to pH 7.0. The only difference between immediate acidity and 

total acidity is the degree of oxidation and hydrolysis proceeding 

during the analysis.

D:lb-l VOLUMETRIC METHOD

The volumetric method is similar in substance to D 1067-51 T, 

ASTM 1 (1954, p. 171-173) Manual on Industrial Water.

Principle of determination

Immediate acidity is determined by rapid titration with a standardized strong base in the cold to minimize the effect of oxidation and 

hydrolysis in producing hydrogen ions.

Apparatus and reagents

pH meter

Sodium hydroxide, 0.02482V, 1.00 mloO.025 mg H+1

Procedure

For precise work the sample for the determination of immediate 

acidity should be collected as directed in sec. A: 3c and analysis should 

proceed as soon as possible after the collection of the sample.

1. Pipet a volume of sample containing less than 1.0 mg H+1 (."0.0 ml max) into 

a 150-uil beaker.

2. Rapidly titrate the solution with 0.02482V NaOH (1.00 mloO.025 mg H+1 ) to 

pH 7.0.

1 American Society for Testing Materials.
D:lc-l ACIDITY 91

Calculations

ppm Immediate acidity (H+1)=^ r X H rXml titrantXO.025 density ml sample

Report immediate-acidity concentrations of <10 ppm to 1 decimal 

place, of between 10 and 999 ppm to whole numbers, and of >999 

ppm to 3 significant figures only.

Preparation of reagents

Sodium hydroxide, 0.0248#, 1.00 mloO.025 mg H+1 : Dilute 12 ml 2N NaOH with 

carbon dioxide-free water to approx 1 liter. Standardize the solution against 

primary standard potassium acid phthalate as follows: Lightly crush 3-4 g 

of the salt to a fineness of approx 100 mesh and dry for 1 or 2 hr at 120°C. 

Dissolve about 2 g, accurately weighed to the nearest milligram, in carbon 

dioxide-free water and dilute to 500.0 ml. Titrate 50.0 ml of the solution 

with the NaOH to p,H 8.6.

AT Normality ... , .. g KHC8H4 04 in 50.0 mlX4.8967 of alkali = - , rr~p 

ml alkali

Store the 0.02482^ base in a tightly capped polythylene bottle.

Sodium hydroxide, 2#: Cover approx 100 g NaOH sticks with water until the 

surface coating is dissolved. Discard the supernatant fluid and immediately 

dissolve the remaining NaOH (about 80 g) in approx 1 liter of carbon dioxidefree water. Store the base in a tightly capped polyethylene bottle.

D:lC POTENTIAL, FREE ACIDITY

Potential free acidity is the capacity of a water containing a nonvolatile compound or compounds, with or without hydrolysis, for 

neutralizing base to pH 7.0.

D:lc-l VOLUMETRIC METHOD

Principle of determination

The sample is heated to drive off uncombiiied gases and to oxidize 

and (or) hydrolyze acid-producing salts (sec. D: la). Boiling is essential for complete expulsion of uncombiiied dissolved gases.

Apparatus and reagents

pH meter 

Hotplate 

Sodium hydroxide, 0.0248#, 1.00 mloO.025 mg H+1

Procedure

1. Pipet a volume of sample containing less than 1.0 lag H+1 (50.0 ml max) into 

a 150-ml beaker.

2. Heat the sample to boiling and continue boiling for 2 min. Vigorous stirring 

may be used to hasten the removal of dissolved gases.

3. Cool the solution to approximately room temperature.

4. Titrate the degassed sample with 0.0248A7 NaOH to pH 7.0.
92 ANALYTICAL PROCEDURES D:lc-l

Calculations

ppm Potential free aciditv (H+i) = , 1 ., X 1)00° Xml titrantXO.025 density ml sample

Report potential free-acidity concentrations of < 10 ppm to 1 decimal 

place, of between 10 and and 999 ppm to whole numbers, and of >999 

ppm to 3 significant figures only.

Preparation of reagents

Sodium hydroxide, 0.0248N, 1.00 mloO.025 nig H+1 : Dilute 12 ml 2AT NaOH with 

carbon dioxide-free water to approx 1 liter. Standardize the solution against 

primary standard potassium acid phthalate as follows: Lightly crush 3-4 

g of the salt to a fineness of approx 100 mesh and dry for 1 or 2 hr at 120°C. 

Dissolve about 2 g, accurately weighed to the nearest milligram, in carbon 

dioxide-free water and dilute to 500.0 ml. Titrate 50.0 ml of the solution 

with the NaOH to pH 8.6.

XT Normality ... . ,. g KHC8H4O4 in 50.0 mlX4.8967 of alkali = - . r. r. 

J ml alkali

Store the 0.0248N base in a tightly capped polyethylene bottle.

Sodium hydroxide, 2N: Cover approx 100 g NaOH sticks with water until the 

surface coating is dissolved. Discard the supernatant fluid and immediately 

dissolve the remaining NaOH (about 80 g) in approx 1 liter of carbon 

dioxide-free water. Store the base in a tightly capped polyethylene bottle.

REFERENCE

American Society for Testing Materials, 1954, Manual on industrial water: 

Spec. Tech. Pub. 148-A.
D :2 ALKALINITY

Alkalinity is the capacity of a water containing a compound or 

compounds, with or without hydrolysis, for neutralizing strong acid 

to pH 4.5. The determination of alkalinity is a measure of the 

excess basic constituents over the amount necessary to balance the 

strong acid constituents. Alkalinity in water is caused primarily 

by the presence of bicarbonates, carbonates, and hydroxides. The 

relative concentrations of hydroxide, carbonate, and bicarbonate are 

a function of the temperature, pH, and concentration of other dissolved solids. The chemist is referred to the work of Langelier 

(1946, p. 169) for a full discussion of the equilibria involved. Minor 

acid radicals such as borates, phosphates, and silicates also add to 

the alkalinity of the water.

Because the alkalinity of many waters is primarily a function of 

the carbonate, bicarbonate, and (or) hydroxide content, the alkalinity determination is sometimes taken as an indication of the 

concentration of these constituents. Such values are maximums and 

include titratable weak acid radicals. Perhaps the effect of phosphate can be corrected for by utilizing the relations given in sec. 

C: 8a, but the Geological Survey has not had sufficient experience 

with the correction to recommend it universally. No suitable 

methods for correcting for the other acid radicals are available.

Carbonates and bicarboiiafes are, common to most waters because 

of the abundance of carbonate minerals in nature and because carbon 

dioxide, which helps dissolve them and other minerals, is readily 

available. Direct contribution to alkalinity by hydroxides is rare 

in nature, and the presence of hydroxides can usually be attributed 

to water treatment or to contamination.

The U.S. Public Health Service (1946) recommends the following 

limitations on alkalinity of drinking and culinary water on carriers 

subject to Federal quarantine regulations:

1. The phenolphthalein (CO3 ) alkalinity (calculated as CaCOs) should not be 

greater than 15 ppm + O.-t X total alkalinity.

2. The normal carbonate alkalinity should not exceed 120 ppm.

3. If excess alkalinity is produced by chemical treatment, the total alkalinity 

(calculated as CaCO3 ) should not exceed the hardness by more than 35 ppm.

The alkalinity, in equivalents per million, in excess of the alkaline 

earths has a bearing on the suitability of a water for irrigation (Eaton, 

1954). If this excess alkalinity, termed "residual sodium carbonate,'' 

exceeds 2.5 epm the water is generally not suitable for irrigation; 

water containing 1.25 to 2.5 epm is marginal, and that containing less 

than 1.25 epm is probably safe (U.S. Salinity Laboratory Staff, p. 81).

93
94 ANALYTICAL PROCEDURES D:2a-l 

D:2a-l POTENTIOMETRIC METHOD

The potentiometric method is similar in substance to that in part I, 

APHA 2 (1955, p. 35-37) Standard Methods, and D 1067-51T, ASTM 

(1954, p. 172-175) Manual on Industrial Water.

Principle of determination

Alkalinity is determined by titrating the water sample with a 

standard solution of strong acid. The equivalency, or end points, of 

the titration are selected as the inflection points in the titration of 

Na2CO3 with H2SO4 . The carbonate end point is taken as pH 8.2 and 

the bicarbonate as pH 4.5. The following reactions occur:

(titration to pH 8.2) 

(from CO3 -2 )

, , ... ,, TT . ~,~ i' (titration to pH 4.5) (natural) -fH 

The presence of hydroxide is indicated if the carbonate titrant volume 

exceeds the bicarbonate titrant volume.

For waters that contain only small quantities of dissolved mineral 

matter, the determination of alkalinity is likely to introduce the 

largest error in the analysis. Reproducibility of results between duplicate samples cannot be expected to be better than about 2 percent. 

As discussed in sec. C : 6, alkalinity is very susceptible to change 

between time of collection and analysis. Changes occur more rapidly 

after the sample bottle is opened. The total-alkalinity value is probably somewhat more stable than the relative values of the common 

alkalinity components. Unless a gross error is made in the initial 

determination of alkalinity, it is seldom advisable to try to check the 

results if several days have elapsed since the bottle was first opened. 

In some water the alkalinity may change appreciably in a few hours.

Additional information on the principle of the determination is 

given by Collins ( 1928 ) .

Apparatus and reagents

Titration assembly, consisting of pH meter, medium-speed mechanical stirrer,

and 50-ml buret. 

Sulfuric acid, 0.01639#, 1.00 mlol.OO mg HCCV1

Procedure

Water samples for the determination of alkalinity should not be 

filtered, diluted, concentrated, or altered in any way.

1. Pipet a volume of sample containing less than 40 mg alkalinity as HCOs'1 

(50.0 ml max) into a suitable beaker.

2. Insert beaker in titration assembly and record the pH.

3. Start the stirrer and proceed immediately with the titration.

4. Record the titrant volume at pH 8.2 and 4.5.

2 American Public Health Association and others.
D:2a-l ALKALINITY 95

Calculations

ppm Alkalinity as CaCO3=j r X , , X0.82Xml titrant density ml sample

tO pH 8-2) ~ (ral to

Ppm C03"2=d^X ml1sam°Ple X [M to pH 8.2X0.9835) -(ppm OH- X 3.527)]

ppm HCOa-^XXlCml pH 8.2 to 4.5) -(ml to pH 8.2)]

Report alkalinity concentrations of < 999 ppm to whole numbers and 

of >999 ppm to 3 significant figures only.

Preparation of reagents

Sulfuric acid, 0.01639#, 1.00 mlol.OO mg HCHV1 : Add 0.5 inl cone H2 SO* (sp gr 

1.84) to 950 inl water. (The titrant is stable for several months if protected 

from ammonia fumes and is usually prepared in larger quantities.) After 

the solution has been thoroughly mixed, standardize it by titrating 25.00 ml 

Na2CO3 (1.00 inlol.OO mg HC(Vl) to pH 4.5.

Sodium carbonate, 1.00 mlol.OO mg HCOs'1 : Dissolve 0.8686 g primary standard 

Na2CO3 in carbon dioxide-free water and dilute to 1,000 ml.

REFERENCES

American Public Health Association and others, 1955, Standard methods for the

examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

Collins, W. D., 1928, Notes on practical water analysis: U.S. Geol. Survey

Water-Supply Paper 596-H. 

Eaton, F. M., 1954, Formulas for estimating the drainage and gypsum requirements of irrigation waters : Texas Agr. Expt. Sta. Misc. Pub. 111. 

Langelier, W. F., 1946, Chemical equations in water treatment: Am. Water

Works Assoc. Jour., v. 38. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts., v. 61, No. 11. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline and

alkali soils: U.S. Dept. Agriculture, Agriculture Handb. 60.

D :3 ALUMINUM

Aluminum forms complex alumino silicates, and combined with 

other metals these alumino silicates are among the most abundant 

minerals in the earth's crust. Yet aluminum is usually only a minor 

constituent in water. It is highly resistant to removal by solution 

during weathering and remains behind persistently during the 

process of rock decomposition in the form of clay minerals in the 

soil and anhydrous minerals in shale and similar sediments. Aluminum is amphoteric and can exist as the natural cation or aluminate 

anion. However, the normal buffer system in natural water tends 

to maintain the pH between 5 and 9, in which range disassociated 

aluminum or aluminate ions will usually not be present in appreciable quantities. Waters of normal pH may contain as much as 

10 ppm of aluminum in the colloidal form. Aluminum ions form 

strong complexes with organic matter, sulfate, and fluoride. This 

tendency stabilizes high concentrations of aluminum in some waters.

No evidence has been found to prove that aluminum is harmful 

to human beings (California State Water Pollution Control Board, 

1952, 1954, p. 174). It is of little importance in irrigation waters, 

although very high concentrations could be toxic to some crops. 

Aluminum compounds in water in concentrations as low as 0.05 ppm 

may cause trouble in industries such as laundries and mineral-water 

plants.

D:3a-l FERRON-ORTHOPHENANTHROLINE METHOD

Principle of determination

Ferron (8-hydroxy-7-iodo-5-quinoline sulfonic acid) reacts with 

aluminum to give a complex that absorbs light in the ultraviolet 

range. A true solution rather than an absorption lake is involved. 

Color development is complete immediately. Normal temperature 

variations do not affect the reaction.

The method is subject to significant interference from iron, but 

this effect can be greatly minimized by adding an iron-complexing 

agent. Orthophenanthroline is ideal for the purpose and has the 

additional advantage that iron may be simultaneously determined. 

A correction for iron must still be made, but it is comparatively small.

Several other metals and anions show small interference effects, 

but only manganese, lead, cobalt, and fluoride show sufficiently

97
98 ANALYTICAI PROCEDURES D:3a-l

pronounced effects to require correction. Beryllium minimizes the 

interference of fluoride. The relative interfering effects of certain 

constituents are indicated by the following results obtained in solutions containing 1.00 ppm of aluminum:

Parts per Al found 

Constituent million (ppm)

Mg_. ..___.__. -

Mg__--------_-

Zn.. -----------

Mn--_-___-____

Mn_____.______

F__-_-_________

F _____________

F .

________ 40

_-_.--__ 80

5

--_.--__ 5

________ 10

---_--_- 1

________ 2

5

1. 04

1.09

1. 05

1. 17

1. 28

. 94

. 90

. 80

Orthophosphate up to 5 ppm and residual chlorine up to 5 ppm 

do not interfere. Natural color and turbidity interfere in the aluminum determination, and a correction is usually required.

Interferences are relatively rare in the associated orthophenanthroline iron determination. Orthophenanthroline is generally considered specific for iron. Copper may interfere if the pH goes 

above 6, but this will not happen if the pH buffer is in good condition. Natural color and turbidity have a measurable but not excessive effect.

Occasionally, small negative absorbancy values are obtained in 

the aluminum determination due to colorless-complex formation with 

calcium. This is harmless, as a reading is always obtained if aluminum is present.

In the analytical sequence, manganese and fluoride should be 

determined before aluminum so that the necessary corrections may 

be applied. These corrections are derived from the data in the 

above table of interferences. They apply with accuracy only at 

the 1-ppm-aluminum level. However, because concentrations of 

interfering ions are generally not greatly in excess of aluminum in 

most waters, the equation holds with satisfactory accuracy for general analytical purposes. If a highly precise aluminum determination is desired in waters containing high concentrations of interfering ions, it is suggested that empirical corrections be made by 

adding similar concentrations of the interfering ions to an appropriate aluminum standard and measuring the resulting absorbancy 

change.

Additional information on the principle of the determination is 

given by Davenport (1949) and by Smith and Eichter (1944).
D:3a-l ALUMINUM

Apparatus and reagents 

Spectrophotometer, Beckman Model B:

Al Fe

Wavelength.. _ _ _ _ _m/i- - 370 _ __-____-_-_--____-

Cells___ _______ _mm_ . 40______-___--_-_--__-

Phototube. ____________ Blue-sensitive. _________

Filter. ___--____-__-___ Blue_._ _______________

Blank.. __..-._ ________ Metal-free water, plus

	reagents. 

Initial sensitivity setting. 2_ ____-___-__-___-___- 2.

Slit width (approx). .mm __ 1.0. __________________ 0.1.

520. 

40.

Blue-sensitive. 

None.

Metal-free water, 

reagents.

plus

The following absorbancies have been observed for iron and aluminum:

mg Fe

0.025. 

.050. 

.075. 

.100. 

.125.

Absorbancy

0. 07 0. 51 

. 125 1. 02 

. 185 1. 53 

. 25 1. 95 

. 305 2. 33

0.025-.-- 

.050__-_ 

.075-.--

Absorbancy 

(370 TOM)

0.64

1. 14

1. 60

Potassium alum, 1.00 ml=0.010 mg Al+3 

Iron chloride, 1.00 ml=0.004 mg Fe 

Hydroxylamine-hydrochloric acid reagent 

Ferron-orthophenanthroline reagent 

Sodium acetate, 35 percent

Procedure

Samples for the determination of aluminum should be collected in 

accordance with directions given in sec. A: 4d.

1. Pipet a volume of sample containing not more than 0.075 mg Al+a or 0.10 mg 

Fe (25.00 ml max) into a 50-ml beaker and adjust the volume to 25.0 ml.

2. Prepare a metal-free water blank and sufficient standards, and adjust the 

volumes to 25.0 nil.

3. Add 2.0 ml NHzOH-HCl reagent and let stand 30 min to permit complete 

reduction of the iron to ferrous iron. If precipitated iron is to be determined, the standing time is prolonged until all iron is in solution (overnight 

standing is a good practice).

4. Add 5.00 ml ferron-orthophenanthroline reagent and stir.

5. Add 2.0 ml NaC*H3O2 . Stir and let stand at least 10 min but not more than 30 

min before taking a reading of the color.

6. Determine the absorbancy of the test sample and standards against the blank 

at 370 nifji and 520 m/u and, when necessary, make corrections for water color 

as directed in sec. C: la-2, method 1; 2.0 ml NaC2H3O2 is added to the colorcorrection blank.
100 ANALYTICAL PROCEDURES D:3a-l

Calculations

1. Determine the mg Fe in the test sample from a plot of absorbancies at 520 m/i 

of standards containing known amounts of the constituent.

2. ppm Fe=j X H :-Xmg Fe in sample density ml sample

3. Determine the apparent mg Al+3 in the test sample from a plot of absorbencies 

at 370 m/i of standards containing known amounts of the constituent.

4. Apparent ppm Al+3=- r X H r-X apparent mg A1+3 in sample * density ml sample *^

5. ppm Al+3=apparent ppm Al+3 0.12 ppm Fe 0.04 ppm Mn+0.05 ppm F'1.

Report aluminum concentrations of <10 ppm to 1 decimal place 

and of > 10 ppm to 2 significant figures only.

Preparation of reagents

Potassium alum, 1.00 ml=0.010 mg Al+3 : Dissolve 1.758 g A1K(SO4 ) 2 -12H2O in 

metal-free water. Add 0.5 ml CHCla and dilute to 1,000 ml with metal-free 

water. Dilute 25.00 ml of this solution to 250.0 ml with metal-free water.

Iron chloride, 1.00 ml=0.004 mg Fe: Dilute 10.00 ml FeCl3 (1.00 ml=0.400 mg Fe) 

to 1,000 ml with metal-free water, containing 1 or 2 drops of cone HC1 (sp 

grl.19).

Iron chloride, 1.00 ml=0.400 mg Fe: Weigh out 0.400 g analytical-grade iron wir> 

which has been cleaned in dilute HC1, rinsed, and dried. Dissolve in a minimum of dilute HC1 and dilute to 1,000 ml.

Hydroxylamine-hydrochloric acid reagent: Dissolve 100 g NH2OH HC1 in metalfree water. Add 40 ml cone HC1 (sp gr 1.19). Add 1 g BeSOv2H2O. Dilute 

to approx 1 liter with metal-free water.

Ferron-orthophenanthroline reagent: Add 0.5 g ferron and 1.0 g of orthophenanthroline to approx 1 liter metal-free water. Stir for several hours until the 

maximum solution is obtained. The ferron will not always dissolve completely. Allow any solids to settle out and decant the clear supernate for use.

Sodium acetate, 35 percent: Dissolve 350 g anhydrous NaCaHsOa in metal-free 

water and dilute to approx 1 liter.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1. 

Davenport, W. H., 1949, Determination of aluminum in presence of iron: Anal.

Chemistry, v. 21. 

Smith, G. F., and Richter, F. P., 1944, Phenanthroline and substituted phenanthroline indicators: G. F. Smith Chemical Co.
D: 4 ARSENIC

Arsenic compounds are present naturally in some waters, but the 

occurrence of quantities detrimental to health is rare. Weed killers, 

insecticides, and many industrial effluents contain arsenic and are 

potential sources of water pollution. The U.S. Public Health Service 

(1946) states that the maximum concentration of arsenic in drinking 

and culinary water on carriers subject to Federal quarantine regulations must not exceed 0.05 ppm. The lethal dose for animals is believed 

to be about 20 mg per animal pound (Miller and Byers, 1935, p. 456). 

Concentrations of 2-4 ppm of arsenic are reported not to interfere 

with the self-purification of streams (Rudolphs and others, 1944, 

p.222).

D:4a-l GUTZEIT-VOLUMETRIC METHOD 

Principle of determination

The arsenic is distilled from the sample as arsine (fig. 15), and its 

concentration determined by oxidation-reduction reactions and volumetric titration. The optimum pH for the reaction in the receiver 

tube is 7.0-9.5. The reactions that occur in the determination are as 

follows: 

In distillation flask 

As2O3+6 H2->2 AsHs+3 H2O

In receiving tube 

2 AsHs+12 HgCl2+3 H2O->12 HgCl+As2Os+12 HC1 

On addition of 20 percent KI 

12 HgCl+24 KI->6 Hg°+6 K2HgI4+12 KC1

On addition of 0.00IN I2

6Hg°+As2Os+8 I2 +12 KI+2 H2O->6 K2HgI4+As2O5+4 HI 

On back titration of excess iodine with standard As2O3 

As2O3+2 I2+5 H2O-»2 HsAsO4+4 IJ+4 H+1

Lead acetate in the distillation apparatus removes all sulfides and 

sulfur dioxide, which would otherwise interfere with the determination. Antimony is not separated from arsenic and may interfere. 

The dimensions of the apparatus (fig. 15) should conform closely with 

that prescribed by Cassil and Wickmann (1939); otherwise, additional refluxing may be required to carry the arsine over quantitatively. The apparatus should be dismantled and cleaned and new lead 

acetate plugs inserted at the beginning of each day's work, but a large 

number of determinations can be made consecutively without cleansing 

or flushing the distillation apparatus. Excess lead acetate should be

101
102 ANALYTICAL PROCEDURES D:4a-l

r-GLASS-WOOL PLUGS IMPREGNATED

5 

J

D

O

WITH LEAD ACETATE-

/<

f 24/40 1

I 1

o

1 ^

! '

S I SLEEVE 

"~~ ADAPTER

 RECEIVING TUE

J (SMALL- DIAME 

OR CENTRIFUGI

J

24/40

v 200-ML 

FLASK

FIGURE 15. Arsine generator for the isolation of arsenic.

blown out of the plugs with air before making the determinations. The 

thermometer may be suspended within the condenser, or a flask with 

a thermometer well may be used.

With the listed apparatus, results are reproducible and accurate to 

±0.002 mg.

Apparatus and reagents

Arsine generator (fig. 15)

Hotplate

Buret, 5-ml, graduated in 0.01 ml

Glass wool

Lead acetate, 10 percent

Mercuric chloride, 1.6 percent

Hydrochloric acid, cone (spgr 1.10)

Potassium iodide, 15 percent

Stannous chloride, 40 percent in cone HC1

Zinc, granulated, low in As

Potassium iodide, 20 percent

Iodine, O.OOltf

Buffer solution

Starch indicator, 1 percent

Arsenous oxide, 1.00 ml=0.050 mg As2O3
D:4a-l ARSENIC 103

Procedure

1. Pipet 1.0 ml HgCL. into the receiving tube of the generator and insert the 

tube in the reflux assembly so that the tip of the delivery tube is below the 

surface of the solution.

2. Allow the hotplate to heat.

3. Pipet a volume of sample containing less than 0.07 mg As (100.0 ml max) 

into the 250-ml liberation flask.

4. Add rapidly, and with mixing, 10 ml cone HC1, 5 ml 15 percent KI, 1 ml 

SnCL, and 4-5 g Zn. Connect the flask immediately to the reflux condenser.

5. Continue refluxing 5 min after the temperature of the solution has reached 

95°C.

6. Disconnect the receiving tube and adapter.

7. Precipitate and redissolve the red Hgl2 with a few drops of 20 percent KI. 

Note : If much As is present, the solution will contain a white or pale-yellow 

precipitate. This precipitate will not be dissolved with the KI.

8. Add 5.00 ml O.OOltf I2 through the adapter and mix well. All precipitates 

should dissolve and the solution still retain the iodine color.

9. Add 2.0 ml buffer solution through the adapter and mix. Connect a small 

rubber bulb to the adapter and wash down the sides of the tube.

10. Add a few drops of 1 percent starch and titrate the excess I2 with As2O3 

( 1.00 ml=0.050 mg As2O3 ) . Record the titrant volume to the nearest 0.01 ml.

11. Determine a reagent blank by taking 100 ml dilution water through the 

procedure given above.

Calculations

Calculations take into consideration the fact that in the AsH3 liberation and subsequent reactions 

6 Hg°

and on titration of excess I2 

1 As2Os02 L

1. Determine the arsenic and iodine equivalents by titrating in the receiver tube 

3.00 ml O.OOltf I2 with As2O3 (1.00 ml=0.05 mg As2O3 ) in the presence of 2.0 

ml buffer and 0.5 ml starch. t

. . ' = iodine equivalent M in titration ml As2O3

 , T" 3 * -1-; =mg As2O3Ol.OO ml I2 in reaction (arsenic equivalent) ml 1. 2 4

2. Calculate ml I2 not consumed in back titration:

ml As2O3 X iodine equivalent = I2 (o

3. Calculate ml I2 used in the reaction:

I2(i) l2(«) = I2 ( r)

4. Calculate mg As distilled:

I2 (r)XAs equivalent = mg As2Os

5. ppm As= . * X . 1>00° . (mg As2O3 -mg blank) X 0.7574 ^ density ml sample v 6 e
104 ANALYTICAL PROCEDURES D:4ar-l

Report arsenic concentrations of <1.0 ppm to 2 decimal places and 

of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Lead acetate, 10 percent : Dissolve 10 g Pb2 (C2H3O2 ) 3OH in water, make just acid 

to litmus, and dilute to approx 100 ml.

Mercuric chloride, 1.6 percent : Dissolve 1.6 g HgCl2 and 0.0.5 g gum arabic in water 

and dilute to 100 ml.

Potassium iodide, 15 percent : Dissolve 15 g KI in water and dilute to approx 

100 ml.

Stannous chloride, 40 percent in cone HC1 : Dissolve 40 g SnCl2 -2H2O in approx 

100 ml cone HC1 (sp gr 1.19) .

Potassium iodide, 20 percent : Saturate approx 75 ml boiling water with KI. 

Add a crystal of I2 and allow KI to recrystallize in the refrigerator. Filter 

off the crystals with suction and wash once with alcohol. Dry the crystals 

overnight in a 180° C oven to sublime the excess I2 . Dissolve 20 g recrystallized KI in water and dilute to approx 100 ml.

Iodine, 0.0012W Dilute 5.00 ml 0.02N I2 to 100.0 ml. This dilution 6f reagent is 

relatively unstable, and its arsenic equivalent in the reaction and iodine 

equi * alent in the backtitration should be determined each day. To determine 

these equivalents add 3.00 ml 0.001# I2 , 2.0 ml buffer, and 0.5 ml starch solution to a receiving tube; titrate with As2O3 (1.00 ml=0.050 mg As2O3 ) (see 

calculation 1 ) . Store the 0.001N I2 in a glass-stoppered bottle protected from 

light.

Iodine, 0.02N: Dissolve 2.54 g I2 and 12 g KI in water and dilute to 950 ml before 

standardizing. Standardize against primary standard As2O3 as follows : 

Dissolve approx 0.04 g As2O3 weighed accurately to 0.0001 g in 10 ml IN 

NaOH. Add 15 ml IN H2 SO4 and mix. Add 50 ml 4 percent NaHCO3 . Titrate slowly with I2 solution, maintaining constant agitation until most of 

the I2 has been reacted (0.04 g As2O3 requires approx 40 ml 0.02N I2 ). Add 

starch solution for indicator and continue the titration until the initial pink 

coloration just passes to clear blue. Deduct from the volume of iodine 

consumed the amount required to produce the same color in a solution composed of the reagents added to 40 ml of freshly boiled and cooled water in 

which 5 g KI has been dissolved.

As2O3 X 20.220 xr Normality ,-. , T of I2= ml

Store in a glass-stoppered bottle protected from light. 

Arsenous oxide, 1.00 ml=0.050 mg As2O3 : Dilute 50.00 ml As2O3 (1.00 inl=1.00

mg As2O3 ) to 1,000 ml. This solution is the basic standard for the arsenic

determination. 

Arsenous oxide, 1.00 nil=1.00 mg As2O3 : Dry 1-2 g Mallinckrodt AR Primary

Standard As2O3 at least 1 hr at 110°C. Dissolve 1.000 g As»O3 in 25 ml 20

percent NaOH saturated with CO2 ; dilute to 1,000 ml with water. Prepare

the 20 percent NaOH by dissolving 20 g NaOH pellets in approx 100 ml water.

Saturate the NaOH solution with CO2 from a CO2 generator : filter off the

Xa2CO3 , if any. 

Starch indicator, 1 percent : Triturate 1 g soluble starch with a little water and

add the suspension slowly to approx 100 ml boiling water. Store in a glassstopptred bottle in a refrigerator. 

Buffer solution: Dissolve 10 g Xa2HPO4-12H2O in water and dilute to approx

100 ml.
D:4a-l ARSENIC 105

REFERENCES

Cassil, C. C., and Wickmann, H. J., 1939, A rapid volumetric micro method for 

determining arsenic: Assoc. Official Agr. Chemists Jour., v. 22.

Miller, J. T., and Byers, H. G., 1935, Selenium spring: Indus, and Eng. Chemistry, News Ed., v. 13.

Rudolfs, Willem, and others, 1944, Critical review of the literature of 1943: 

Sewage Works Jour., v. 16.

U. S. Public Health Service, 1946, Drinking water standards: IJ. S. Public 

Health Service Repts, v. 61, no. 11.

D:5 BARIUM

Many mineral salts of barium are soluble, but barium sulf ate is 

very insoluble. Consequently barium ions will be quickly precipitated from waters containing sulfate. Barium salts are used in 

dyeing, paint, and some explosives manufacture, and in tanning 

operations. Barium in measurable quantities is generally limited to 

certain types of brines and industrial wastes and to normal waters 

whose sulfate content is extremely low. Barium is not determined 

in normal waters having a sulfate concentration greater than 2 ppm.

The U.S. Public Health Service (1946) states that salts of barium, 

which have a deleterious physiological effect, must not be added to 

drinking and culinary water on carriers subject to Federal quarantine regulations.

D:5a-l GRAVIMETRIC METHOD

The gravimetric method is recommended for waters whose barium 

content is greater than 10 ppm and for brines and industrial wastes 

when the complexometric method fails (see sec. D: 5a-2).

Principle of determination

Barium is precipitated from acid solution as the sulfate, ignited, 

and weighed. Calcium and strontium also form slightly soluble 

sulfates and may precipitate with the barium sulfate. Iron, aluminum, and manganese may be occluded by the barium sulfate and 

contaminate the precipitate. There is always the possibility of silica 

coprecipitating. These possible interferences are particularly significant when the amount of barium is small, but they are eliminated 

or sufficiently minimized if the gravimetric barium determination 

is made on the filtrate from the gravimetric calcium determination 

(see sec. D: 8a-3, procedure 13). Coprecipitation of other ions in the 

filtrate with barium sulfate are other sources of error in the determination; they are classics in quantitative analysis and are discussed in detail by Kolthoff and Sandell (1952) and Hillebrand 

and Lundell (1929, p. 486-505).

When the barium concentration in the original water samples is 

less than 10 ppm, the quantity of barium sulfate precipitate is so 

small that results obtained by the gravimetric method may be in 

appreciable error. The accuracy and reproducibility of results are 

dependent on the quantitative removal of strontium and on the 

amount of coprecipitation of foreign ions with barium sulfate. With 

careful work the results may be accurate and reproducible to 0.25 mg.

107
108 ANALYTICAL PROCEDURES D:5a-l

Apparatus and reagents

Steam bath

Porcelain crucibles

Muffle furnace

Hydrochloric acid, cone (sp gr 1.19)

Sulfuric acid, 5 percent v/v

Sulfuric acid, cone (sp gr 1.84)

Procedure

1. Dilute or concentrate the filtrate from the gravimetric calcium determination 

(see sec. D : 8a-3, procedure 13) to a convenient volume. Take an aliquot 

containing more than 1 mg Ba+3 for the determination.

2. Concentrate the aliquot to about 50 ml by evaporation.

3. Transfer the concentrate to a casserole and evaporate to dryness.

4. Decompose the ammonium salts by heating over an open flame.

5. Take up the residue with hot water.

6. Add 5-10 drops cone HC1.

7. Filter off any insoluble material through Whatman No. 42 paper.

8. Wash the insoluble material 3 times with hot water and combine the washings with the filtrate.

9. Dilute to approx 100 ml if necessary.

10. Heat the filtrate to boiling and add sufficient 5 percent H2 SO4 to precipitate the barium and to provide a small excess.

11. Digest the precipitate overnight on a steam bath.

12. Quantitatively transfer the precipitate to Whatman No. 42 filter paper. The 

beaker should be scrubbed at least 3 times with a rubber policeman.

13. Wash the beaker and precipitate with hot water until a sample of the 

filtrate will develop no turbidity with AgNOs.

14. Slowly ignite the precipitate in a tared porcelain crucible over a low, 

oxidizing flame until the filter paper is reduced to white ash.

15. Transfer the crucible to the muffle furnace and ignite at 800°G for 1 hr.

16. Cool the crucible in a desiccator and weigh.

17. After weighing, add 1 drop cone H2 SO4 . Fume off and reweigh. The 

weighings in steps 16 and 17 should be identical. A discrepancy indicates 

probable coprecipitation of chlorides, and the weight taken in step 17 

should be used in the calculation.

18. If the purity of the BaSO4 in respect to calcium and strontium is questionable, dissolve the precipitate completely in the crucible in 5 ml hot cone 

H2 SO4 . Cool and pour into 50 ml water and dilute to 100 ml. Digest for 

1 hr on the steam bath, filter and wash with a little hot water, ignite, and 

weigh (Hillebrand and Lundell, 1929).

Calculations

ppm Ba+2= 1 . X *'°°Q , Xmg BaSO4 X0.5884 density ml sample

Report barium concentrations of <1() ppm to 1 decimal place of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 8 

significant figures only.
D:5a-2 BARIUM 109

Preparation of reagents

Sulfuric acid, 5 percent v/v: Mix 50 ml cone H2 SO4 (sp gr 1.84) with water and 

dilute to approx 1 liter.

D:5a-2 COMPLEXOMETRIC METHOD

The complexometric method is applicable to most natural and treated 

waters whose barium content ranges from 0 to 10 ppm. The method 

fails conspicuously at times with acid or polluted waters that contain 

excessive amounts of heavy metals.

Principle of determination

Barium is determined by the difference between the titrant volume 

of disodium dihydrogen ethylenediamine tetraacetate (Na2EDTA) required by an untreated sample and a sample from which barium has 

been precipitated as barium sulfate. Na2EDTA forms a slightly ionized colorless stable complex with the alkaline earths. A solution of 

Eriochrome Black T dye is bright blue in the absence of alkaline earths 

but with them forms a deep red complex which has a higher ionization 

constant than the Na2EDTA complex. Hence by using Eriochrome 

Black T as an indicator, the alkaline earths can be titrated with 

Na2EDTA. For example, with calcium:

o o

NaO C CH2 CH2 C ONa

N CH2 CH2 N +Ca+2

HO C CH2 CH2 C OH

II II

O I O

NaO C CH2 CH2 C ONa

II \ /I!

O N CH2 CH2 N O

CH2 CH2 +2H+2 

\ Ca /

Barium titrates approximately stoichiometrically with calcium. The 

optimum pH for the reaction is 10.4 or higher.

The determination is not strictly quantitative because coprecipitation errors are involved in the precipitation of barium, especially in 

the presence of strontium. Interferences of heavy metals with the 

Na,EDTA titration (see sec. D:l7a-l) are not significant to the 

barium determination because their net effect may be assumed to be 

identical in both samples. Interference of heavy metals with detec-
110 ANALYTICAL PROCEDURES

tion of the end point, however, may cause an error in the determination 

of barium unless this interference is removed. Precision is improved 

if the end point is compared with a color blank of dilution water plus 

reagents.

Results obtainable are generally accurate and reproducible to ±0.1 

mg in the 0-10 ppm range.

Apparatus and reagents

Titration assembly: Some analysts prefer to use conventional lighting and hand 

stirring. Others have reported better results by using a visual titration 

assembly consisting of motor-driven stirrer, 25-ml buret, white-porcelain-base 

buret holder, and shaded incandescent lamp. The sample beaker is placed 

near the front of the porcelain base and the reaction is viewed diagonally 

downward through the side of the beaker and against the white background. 

Illumination is from behind the beaker and in the plane of vision.

Photometric-titration assembly, optional (see sec. C: 2h)

Hydroxylamine hydrochloride, 3 percent

Ammonium hydroxide, cone (sp gr 0.900)

Sodium cyanide, 2.5 percent

Eriochrome Black T indicator solution

Na2EDTA, 1.00 mlol.OO mg CaCO3

Sulfuric acid, 5 percent v/v

Sodium sulfate, 10 percent

Procedure

1. Pipet a volume of sample containing less than 0.5 mg Ba+2 (50.00 ml max) 

into a 150-ml beaker and adjust the volume to approx 50 ml.

2. Insert the beaker in the titration assembly and start stirrer.

3. Add 1 ml 3 percent NH2OH-HC1.

4. Add 1 ml cone NH4OH. (If NH4OH not tightly stoppered it tends to lose 

Strength; weak NH4OH will not buffer the solution to the necessary pH.)

5. Add 2 ml 2.5 percent NaCN (CAUTION deadly poison). NaCN may be 

eliminated unless the presence of interfering metals decreases the sharpness of the end point.

6. Add 2.0 ml Eriochrome Black T indicator.

7. Titrate with NaJJDTA U-00 mlol.OO mg CaCoa) until blue or purple swirls 

begin to show. The end point is reached when all traces of red and purple 

have disappeared and the solution is clear blue in color. The detection of 

the end point may be facilitated by comparison of the titration solution 

with a color blank prepared of metal-free water and the reagents.

8. Record the titrant volume (mlt) to the nearest 0.05 ml.

9. Measure approx 75 ml sample into a 150-ml beaker.

10. Add 1.5 ml 5 percent H2 SO4.

11. Add 1.5 ml 10 percent Na2SO4 .

12. Mix and allow the sample to stand several hours for the BaSO4 to settle.

13. Pipet a volume of clear supernatant identical with that taken in step 1 and 

adjust the volume to approx 50 ml.

14. Treat the sample as directed in steps 2-8, with the exception of adding 2 ml 

NH4OH in step 4. The additional NH4OH is required to neutralize the acid 

added in step 10.
D:5a-2 BARIUM 111

Calculations

ppm Ba^-p-Vx i 1 ' 000 . 

density ml sample

where ml»=titrant volume for the first titration and 

volume after barium is removed.

Report barium concentrations of <10 ppm to 0.1 ppm, of between 

10 and 999 ppm to whole numbers, and of >999 ppm to 3 significant 

figures only.

Preparation of reagents

Hydroxylamine hydrochloride, 3 percent : Dissolve 30 g NH2OH-HC1 in metal-free 

water and dilute to approx 1 liter.

Sodium cyanide, 2.5 percent (CAUTION NaCN is a deadly poison and the 

reagent solution must be so marked) : Dissolve 2.5 g NaCN in metal-free 

water and dilute to approx 100 ml.

Eriochrome Black T indicator solution : Dissolve 0.40 g Eriochrome Black T in 

100 ml metal-free water and dilute to approx 1 liter with 95 percent ethyl 

alcohol. This indicator is stable for at least 2 months. The Eastman 

reagent product has been found to be satisfactory.

Na2EDTA, 1.00 mlol.OO mg CaCO3 : Dissolve 3.720 g Na2EDTA, which has been 

dried overnight in an H2 SO4 desiccator, in metal-free water and dilute to 

1,000 ml. Check the titer of the reagent by titrating 25.00 ml CaCl2 (1.00 

mlol.OO mg CaCOs) as described in the procedure for sample analysis. This 

Na2EDTA solution is identical to that used in the hardness determination 

(Sec. D: 17a-l)

Calcium chloride, 1.00 mlol.OO mg CaCO3 : Suspend 1.000 g CaCO3 , dried at 

180°C for 1.0 hr before weighing, in approx 600 ml metal-free water and 

dissolve cautiously with a minimum of dilute HC1. Dilute to 1,000 ml.

Sulfuric acid, 5 percent v/v : Mix 5 ml cone H2SO4 (sp gr 1.84) with metal-freee 

water and dilute to approx 100 ml.

Sodium sulfate, 10 percent : Dissolve 10 g Na2SO4 in metal-free water and dilute 

to approx 100 ml.

REFERENCES

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis : New

York, John Wiley and Sons, Inc. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis : New York, Macmillan Co., 3d ed. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public

Health Service Repts., v. 61, no. 11.

D :6 BORON

The boron content in most surface waters seldom exceeds 1 ppm, 

but some ground waters, particularly those affected by igneous rocks, 

may contain appreciably more boron. The form or forms in which 

boron exists in natural waters is not known with certainty and may 

not be uniform for all waters; hence, the boron content is reported 

as elemental boron.

In normal concentrations, boron in drinking water is not regarded 

as a hazard to humans, but the boron content of water to be used 

for irrigation is highly important. Boron is an essential plant 

nutrient, but the presence of as little as 1 ppm has been known to 

injure citrus trees and walnut and bean crops, while waters containing more than 2 ppm will in time cause trouble with many 

common crops (U.S. Salinity Laboratory Staff, 1954, p. 81).

D:6a-l DIANTHRIMIDE METHOD

The dianthrimide method is susceptible to high concentrations of 

oxidizing or reducing material and dissolved organic matter but 

is not affected by buffering solutions or high concentrations of total 

salts. It is satisfactory for most natural waters.

Principle of determination

Boron when heated with l,l'-dianthrimide in concentrated sulfuric acid gives a colored complex (Ellis, Zook, and Baudisch, 1949, 

p. 1345-1348). The color change ranges from greenish-yellow to 

blue. The reaction producing the blue color depends on the nature 

of the vessel in which the reaction occurs, the temperature and 

duration of heating, and the concentration of reagent and of boron. 

In platinum vessels, the reagent is unstable and darkens with no 

blue-color formation. Maximum color development is obtained 

after the reaction has proceeded for 3 hours at 90°C.

Traces of moisture precipitate the reagent and interfere in the 

determination; therefore, precautionary measures given in the procedure must be followed explicitly. Nitrate and bicarbonate interfere with color development unless removed. Removal as nitric 

acid and carbon dioxide by evaporation of the sample in the presence 

of sulfuric acid is employed. Organic matter in high concentrations chars and causes a discoloration of the complex, but this interference is easily recognized; small quantities of organic material 

cause no trouble. Some success in removal of the organic-material 

interference has been obtained by heating the sample in the presence 

of hydrogen peroxide for 1-2 hours, but it is essential that all nascent 

oxygen be volatilized before the dianthrimide is added to the sample.

113
114 ANALYTICAL PROCEDURES D:6a-l

When peroxide digestion is used, the final color complex should be 

compared with standard boron solutions similarly treated. Oxidizing and reducing constituents also interfere. In cleaning glassware 

avoid the use of chromic-sulfuric acid.

Some boric acid is probably volatilized during the evaporation of 

the sample in the presence of sulfuric acid. Prolonged heating, or 

temperatures higher than that recommended, volatilize an excessive 

amount of boron and decrease the sensitivity of the test. The loss of 

boron is proportional to the boron content of the sample or standard, 

hence such loss in no way affects the linearity of the color development 

if the heating is uniform. Nonlinearity of the concentration-versusabsorbancy curve can result from weak reagents. The two standards 

in step 2 of the procedure act as a check on linearity of the reaction 

and suitability of the working reagent.

With the listed apparatus, results are generally accurate and reproducible to ±0.0001 mg.

Apparatus and reagents

Oven, 90°C: Uniformity of temperature throughout the oven is imperative.

Fisher Isotemp ovens have been found to be satisfactory. 

Stoppers to fit absorption cells: Teflon or polyethylene are satisfactory, rubber

can be used with caution, glass-stoppered cells are preferable. 

Spectrophotometer, Becknian Model B : 

Wavelength: 620 mju 

Cells : 23-mm or 25-mm optical depth 

Phototube: Blue-sensitive

Blank : Dilution water carried through the procedure with the sample 

Initial sensitivity setting: 2 

Slit width : 0.6 mm (approx) 

The following absorbancies have been observed: 

mg B Absorbancy

0. 001 0. 26

. 002 . 52

. 005 1.30

Sodium tetraborate, 1.00ml=0.001 mg B. 

Sulfuric acid, cone (sp gr 1.84) 

l,l'-dianthrimide, working reagent

Procedure

1. Pipet a volume of sample containing less than 0.005 mg B (5.00 ml max) into 

the absorption cell and adjust volume to 5.0.ml.

2. Prepare a blank of redistilled water and sufficient standards and adjust 

volumes to 5.0 ml.

3. Add 1.0 ml cone H2 SO4 and mix by swirling the contents of the cell.

4. Evaporate overnight in an oven at 90°C. At the end of the evaporation, the

solution volume should be between 1.0 and 0..1 ml. 

r>. Add 5.0 ml dianthrimide working reagent, and mix by swirling the contents

of the cell.
D:6a-2 BORON 115

6. Incubate for 3.0 hr at 90°C.

7. Cool and immediately dilute with 10.0 inl cone HaSO*

8. Mix thoroughly but carefully with a stirring rod. The contents must not be 

spattered on the upper walls of the cell.

9. Stopper the cells but do not let the stopper come in contact with the acid 

contents.

10. Remove all traces of acid, reagent, and fingerprints from the exterior surface 

of the cell.

11. Determine the absorbancy of the sample and standards against the blank.

Calculations

1. Determine the ing B in the sample from a plot of absorbancies of standards 

containing known amounts of the constituent.

2. ppm B=-3 X r-^ r-X mg B in sample 

density ml sample

Report boron concentrations of <1.0 ppm to 2 decimal places and 

of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Sodium tetraborate, 1.00 ml=0.001 mg B : Dilute 10.00 ml Na^O (1.00 ml=0.100 

nig B) to 1,000 ml with redistilled water. Store in a plastic bottle.

Sodium tetraborate, 1.00 ml=0.100 mg B : Dissolve approx 10 g NajB4<V10H2O in 

redistilled water at 50°-60°C. Recrystallize by placing in refrigerator for 

several hours. Dry by removing the water with suction and washing with 

alcohol followed by ether. Do not dry in oven. Dissolve 0.8813 g in redistilled water and dilute to 1,000 nil. Store in plastic bottle.

l,l'-dianthrimide, working reagent: Dilute 1 volume of stock reagent to 20 

volumes with cone H-SO* (sp gr 1.84). The reagent is stable for long periods 

of time if the container is sealed with sealing wax and is kept in a refrigerator.

l,l'-dianthrimide, stock reagent: Dissolve 200 mg in 50 ml cone H-SO* (sp gr 

1.84). The reagent is stable for long periods if the container is sealed 

with sealing wax and is kept in a refrigerator.

D:6a-2 POTENTIOMETRIC METHOD

The potentiometric method is similar in substance to part IB, APHA 

(1955, p. 46-49) Standard Methods, and method 73a, U.S. Salinity 

Laboratory Staff (1954, p. 140-141) Handbook 60.

The method is susceptible to buffering agents and high concentrations 

of dissolved minerals but is not particularly affected by the presence of 

organic material.

Principle of determination

Boric acid with an ionization constant of 5.5 X10"10 cannot be titrated 

directly with base. However, certain polyhydroxy organic compounds, such as mannitol, form complex acids with boric acid that are 

much stronger than boric acid itself. The quantity of alkali required 

to titrate the complex acid back to the initial pH of the boric acid is a
116 ANALYTICAL PROCEDURES D:6a-2

measure of the boron present (Foote, 1932). The following reactions 

are involved :

O,H14O.+H3BO3-» Cell I2O4 BO3H+2H2O 

C6H12O4-BO3H+NaOH-»CeH12O--BO3Na-|-H2O

The pH of the water is adjusted to 7.00 with sulfuric acid and (or) 

sodium hydroxide, and mannitol is then added. The liberated hydrogen ion is backtitrated to 7.00 pH with standard sodium hydroxide 

(Wilcox and Hatcher, 1947). The pH meter with glass electrode is 

used.

The method is particularly susceptible to the effects of buffering 

systems. The water is freed of carbon dioxide by acidification and 

boiling. Iron, aluminum, and probably most of the weak ions that 

tend to buffer the solution interfere to some extent unless removed. 

Phosphate reacts with mannitol, though not quantitatively. The reliability of the results generally decreases in proportion to the total 

concentration of the dissolved material.

With careful treatment, results with dilute alkaline waters and 

in the absence of buffering action may be accurate and reproducible 

to ±0.005 mg but are appreciably less with most natural waters. 

Results are not reliable with acid waters that contain hydrolyzable 

salts.

Apparatus and reagents

Water bath for cooling

pH meter

Electric stirrer: The stirrer should adequately mix the solution but should not

aerate it excessively. Magnet-type stirrers are recommended. 

Buret, 10-ml 

Methyl orange indicator 

Sulfuric acid IN

Sodium hydroxide, 0.5A7 , carbonate-free 

Sulfuric acid, 0.02N 

Mannitol 

Sodium hydroxide, 0.0231A7, 1.00 ml=0.25 mg B

Procedure

1. Measure a volume of sample containing less than 1 mg B (250 ml max) into 

a 600-ml beaker and adjust the volume to 250 ml.

2. Add 2-3 drops methyl orange.

3. Acidify with 12V H2 SO4 .

4. Bring to boil and stir cautiously for 3 or 4 inin, then vigorously i'or 1 niin to 

expel CO2 . If the sample reverts to the alkaline color during the boiling 

period, adjust the pH with IN H2SO4 to the acid color.

5. Cover the Sample with a watch glass and rapidly cool to room temperature in 

a water bath. The cooling should not require more than 30 min.

6. Insert the beaker in the titration apparatus and start the stirrer.

7. Adjust the pH to 7.00 with 0.5A7 NaOH, 0.002AT H,SO4 , and 0.0231IV XaOH. 

The indicator needle of the pH meter should be steady and not drift from 

the reading of pH 7.00.
D:6a-2 BORON 11,7

8. Record the buret reading to 0.01 ml.

9. Add approximately 5 g mannitol. If boron in the form of boric acid is present, 

the pH will drop to some value below 7.00.

10. Titrate rapidly back to the initial pH of 7.00 with 0.02312V7 NaOH. Use caution near the end point to permit a slight lag in the pH-ineter response.

11. When the indicator needle remains steady at pH 7.00 for at least 10 sec, record 

the volume of XaOH used to the nearest 0.01 ml.

12. Determine a blank correction for the reagents by carrying 2.10 ml carbon 

dioxide-free water through the complete procedure. The normal blank is 

0.06 ±0.01 ml.

Calculations

ppm B = 1 .. X 1 1 '0°° . X[/(ml NaOH-ml blank)]X0.25 density ml sairple

where/=factor for 0.02312V7 NaOH.

Report boron concentrations of < 1.0 ppm to 2 decimal places and of 

> 1.0 ppm to 2 significant figures only.

Preparation of reagents

Methyl orange indicator: Dissolve 0.05 g methyl orange in water and dilute to 

approx 100 ml.

Sulfuric acid, 0.022V7 ; Dilute 20 ml IN H2 S(X to approx 1 liter.

Sulfuric acid, 12V7 : Mix 28 ml cone H2 SO4 (sp gr 1.84) with water and dilute to 

approx 1 liter.

Sodium hydroxide, 0.52V7 , carbonate-free: Prepare a saturated solution of NaOH 

in alkali-resistant flask (Corning No. 728). Stopper and let stand until 

the Na2CO3 precipitates. Titrate a portion of supernatant liquid with 1N 

H2 SO4 and dilute to proper strength with carbon dioxide-free water. Store 

the solution in a stoppered polyethylene bottle.

Sodium hydroxide, 0.02317V7 , 1.00 mloO.25 mg B: Dilute 46 ml 0.52V' NaOH to 950 

ml with carbon dioxide-free water. Standardize by titrating solutions containing 0.25, 0.50, and 1.00 mg B which have been carried through the complete procedure for the determination of boron. Despite the usual precautions, the normality of this dilute base is not very stable and a factor should 

be computed every 2 or 3 weeks. Use the following equation to calculate 

the factor (/) for the 0.02317V7 NaOH :

ml theoretical titration 

J -iml actual titration ml blank

The normal blank is O.()6±0.01 ml. Accepted / values are 0.9-1.1. Protect 

the solution from carbon dioxide by means of a soda lime tube.

Sodium tetraborate, 1.00 ml = 0.010 mg B : Dilute 100.0 ml Na2B4O7 (1.00 ml = 

0.100 mg B) to 1,000 ml with redistilled water. Store in plastic bottle.

Sodium tetraborate, 1.00 ml=0.100 mg B: Dissolve approx ,10 g Na2B4O7 -loaO 

in redistilled water at 50°-60°C. Recrystallize by placing in refrigerator 

for several hours. Dry by removing the water with suction and washing 

with alcohol followed by ether. Do not dry in oven, -Dissolve 0.8813 g in 

redistilled water and dilute to 1,000 ml. Store in plastic bottle.
118 ANALYTICAL PROCEDURES D:6a-3 

D:6a-3 CARMINE METHOD

The carmine method is similar in substance to part IA, APHA 

(1955, p. 45-46) Standard Methods, and method 73b, U.S. Salinity 

Laboratory Staff (1954, p. 142) Handbook 60.

The method finds maximum utility for waters whose boron content exceeds 1 ppm or when the precision required does not exceed 

0.1 ppm.

Principle of determination

In acid solution, boron forms a colored complex with carmine. 

The color change is from red to blue. Maximum color development 

requires approximately 1 hr, but the color intensity decreases thereafter. The fading is proportional to the boron content and is 

therefore not consequential if standards are run simultaneously 

with the sample. Special attention must be given to the time 

element because the fading is rather rapid, approx 25 percent or 

more reduction in intensity between 1.0 and 2.0 hr. The optimum 

range for the procedure on undiluted or unconcentrated samples is 

0.5-10 ppm of boron. Samples containing less than 0.5 ppm must be 

concentrated by special evaporation. The concentration is carried 

out in strong alkaline solution to prevent the loss of volatile boric 

acid.

Ammonium, molybdenum, cerium, chloride, calcium, magnesium, 

sodium, and potassium are reported not to interfere. Strong ammonia fumes affect the reagent. Silica interferes, but the interference is independent of the silica concentration and dependent on 

the boron concentration. With 0.5 ppm of boron the determined 

value may be about 20 percent high in the presence of 5 or 30 ppm 

of silica. At lower boron levels the effect is variable. Fluoride, 

nitrate, and phosphate contribute some interference but to a lesser 

degree.

With the listed apparatus, results are reproducible to ±0.0005 nig 

in the 0.002- to 0.020-mg range. Accuracy is dependent on effect 

of interfering substances.

Additional information on the principle of determination is given 

by Hatcher and Wilcox (1950).

Apparatus and reagents

Steam bath

Spectrophotometer, Beckman Model B : Wavelength: 600 m/i

Cells : 25-mm optical depth

Phototube: Blue-sensitive

Blank: Dilution water carried through the procedure with the sample
D:6a-3 BORON 119

Initial sensitivity setting: 2

Slit width: O.lmm (approx)

The following absorbancies have been observed :

mg B Absorbancy

0. 005 0. 24

. 010 .48

. 015 . 70

. 020 .93

Sodium tetraborate, 1.00 ml=0.010 mg B 

Sodium hydroxide, 2 percent 

Hydrochloric acid, 5 percent v/v 

Hydrochloric acid, cone (sp gr 1.19) 

Sulfuric acid, cone (sp gr 1.84) 

Carmine solution, 0.05 percent

Procedure

1. Pipet a volume of sample containing between 0.001 and 0.020 mg B (2.00 ml 

max) into a 150-ml flask. If the sample contains less than 0.5 ppm B, take 

a 2.00-ml aliquot of sample concentrated as follows :

a. Pipet an accurately measured volume of sample (100 ml max) into

an evaporating dish, 

b. Add 1 ml 2 percent NaOH. 

c. Evaporate on a steam bath to dryness.

d. Take up the residue, triturate in 5.00 ml 5 percent HC1 with a 

rubber policeman, and centrifuge to obtain a clear solution.

2. Prepare a blank and sufficient standard and adjust volume to 2.0 ml.

3. Add 2 drops cone HC1.

4. Add 12 ml cone H2 SO4 , mix, and allow to cool for at least 30 min.

5. Add 10 ml 0.05 percent carmine solution and mix.

6. Allow the solution to stand approx 1 hr.

7. Determine the absorbancies of the sample and standards against the blank 

and when necessary make corrections for water color as directed in sec. 

C : la-2, method 1.

Calculations

1. Determine the mg B in the sample from a plot of absorbancies of standards 

containing known amounts of the constituent.

2. ppm B = - X ; ^ ^-Xrng B in sample density ml sample

Keport boron concentrations of <1.0 ppm to 2 decimal places and 

of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Sodium hydroxide, 2 percent: Dissolve 2 g NaOH in water and dilute to approx

100 ml. 

Hydrochloric acid, 5 percent v/v: Mix 50 ml cone HC1 (sp gr 1.19) with water

and dilute to approx 1 liter. 

Sodium tetraborate, 1.00 ml=0.010 mg B : Dilute 100.0 ml Na2B4O7 (1.00 ml=0.100

mg B) to 1,000 ml with redistilled water. Store in plastic bottle.
120 ANALYTICAL PROCEDURES D:6a-3

Sodium tetraborate, 1.00 ml=0.100 mg B: Dissolve approx 10 g Na2B4O7-10H2O 

in redistilled water at 50°-60°C\ Recrystallize by placing in a refrigerator 

for several hours. Dry by removing water with suction and washing with 

alcohol followed by ether. Do not dry in oven. Dissolve 0.8813 g in redistilled water and dilute to 1,000 ml. Store in plastic bottle.

Carmine solution, 0.05 percent: Suspend 0.50 g carmine in 1,000 ml cone H2 SO4 

and mix with a mechanical stirrer until solution is complete. The stain 

certified by the Biological Stain Commission has been used satisfactorily.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

Ellis, G. H., Zook, E. G., and Baudisch, Oskar, 1949, Colorimetric determination

of boron using l,l'-dianthrimide: Anal. Chemistry, v. 21. 

Poote, F. J., 1932, Determination of boron in water: Indus, and Eng. Chemistry,

Anal. Ed., v. 4. 

Hatcher. J. T., and Wilcox, L. V., 1950, Colorimetric determination of boron

using carmine: Anal. Chemistry, v. 22. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline and

alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60. 

Wilcox, L. V., and Hatcher, J. T., 1947, Methods of analysis used in the

Rubidoux Laboratory: U. S. Dept. Agriculture Bur. Plant Industry, 5th ed.
D:7 BBOMIDE

Bromide is a very minor element in the earth's crust and is normally 

present in natural waters in only minute quantities. Measurable 

amounts may be found in some streams that receive industrial wastes, 

and some natural brines may contain rather high concentrations.

D:7a-l OXIDATION METHOD

The oxidation method is similar in substance to D 1246-53 T, ASTM 

(1954, p. 260-263) Manual on Industrial Water.

Principle of determination

The determination of bromide involves the determination of bromide 

and iodide collectively and then the determination of iodide alone; 

bromide is calculated by difference. Bromide and iodide are oxidized 

to bromate and iodate, respectively, by hypochlorite, and the excess 

hypochlorite is subsequently decomposed with sodium formate (Kolthoff and Sandell, 1952, p. 585-605) .

3 OCl-'+ 

3 OCl-'

Iodine equivalent to the combined iodate and bromate is then liberated 

by addition of potassium iodide to an acid solution.

The liberated iodine is then titrated with standard thiosulfate using 

starch as the indicator.

Iodide alone is determined by oxidation to iodate with bromine in 

a buffered solution. Iodine equivalent to the iodate is then liberated 

from added potassium iodide and titrated with thiosulfate as in the 

combined determination.

Iron, manganese, and organic material interfere with the basic reactions of the method, but their interferences are removed by preliminary treatment with calcium oxide.

Results are accurate and reproducible to ±0.05 mg.

Apparatus and reagents

Iodine flasks, 250-ml

Buret. 10-ml

Calcium oxide, anhydrous powder

Sodium chloride, crystals

Methyl red indicator solution, 0.01 percent

Hydrochloric acid, 50 percent v/v

121
122 ANALYTICAL PROCEDURES D:7a-l

Potassium hypochlorite, 4.4 percent 

Calcium carbonate, powder 

Sodium formate, 50 percent 

Sodium molybdate, 0.85 percent 

Potassium fluoride, crystals 

Potassium iodide, crystals 

Sulf uric acid, 20 percent v/v 

Sodium thiosulfate solution, 0.0102V 

Starch indicator, stable 

Sodium acetate solution, 16.5 percent 

Acetic acid, 12.5 percent v/v 

Bromine water, saturated

Procedure

1. Remove soluble iron manganese, and organic matter by adding a slight excess 

of CaO to approx 400 ml of sample, shake, let stand about 1 hr, and filter 

through dry paper. Discard the first 75 ml of filtrate.

2. For the combined Br"1 and I"1 determination pipet a volume of the filtrate containing less than 5.0 mg Br-1 and I'1 (100.0 ml max) into a 250-ml iodine 

flask and adjust the volume to approx 100 ml.

3. Prepare a blank of approx 100 ml water and carry it through the procedure 

with the sample.

4. Add sufficient NaCl to produce a 3.0-g Cl'1 content in the sample.

5. Add a drop of 0.01 percent methyl red indicator, and neutralize the solution 

with 50 percent HC1.

6. Add 10 ml 4.4 percent KC1O.

7. Add 0.5 ml 50 percent HC1.

8. Add sufficient CaCOs to produce an excess of approx 0.1 g.

9. Heat the solution to boiling and maintain this temperature for about 8 min.

10. Reduce the excess KC1O by adding 2 ml 50 percent NaCHO2 , taking precautions to wash down the sides of the flask with a small amount of hot water. 

Keep the solution hot for an additional 8 min.

11. Cool and add several drops of 0.85 percent Na2MoO4.

12. If any iron precipitates at this point, add 0.5 g KF-2H2O.

13. Add approx 1 g KI.

14. Add 10 ml 20 percent H2 SO4 .

15. Let stand 5 min in the dark.

16. Titrate the liberated I2 with O.OWN Na2 S2O3 , adding 2-3 ml starch indicator 

as the end point is approached. Disregard the return of the blue color after 

the end point has been reached.

17. For the I"1 determination, pipet a volume of filtrate (step 1) containing less 

than 5.0 mg I-1 (100.0 ml max) into a 250-ml iodine flask and adjust the 

volume to approx 100 ml.

18. Prepare a blank of 100.0 ml water and carry through the procedure with

the sample. 

15). Add a drop of 0.01 percent methyl red indicator and make the solution just

acid with 20 percent H2 SO4 .

20. Add 15.0 ml 16.5 percent NaC2H3O2 .

21. Add 5.0 ml 12.5 percent HC2H3O,.

22. Add sufficient bromine water to produce a light-yellow color, mix, and allow 

to stand 5 min.
D:7a-l BROMIDE 123

23. Reduce the excess Brs by adding 50 percent NaCHOa until the yellow tinge 

in the sample disappears, then add an excess of 1 ml.

24. Wash down the sides of the flask with a small amount of water and blow 

out Bra vapors with a syringe and a glass tube inserted through the mouth 

of the flask.

25. If any iron precipitates at this point, add 0.5 g KF-2H2O.

26. Add approx 1 g KI.

27. Add 10 ml 20 percent H^SO*.

28. Let stand 5 min in the dark.

29. Titrate the liberated I2 with O.OlOtf Na2S2O3, adding 2-3 ml starch indicator 

as the end point is approached. Disregard the return of the blue color after 

the end point has been reached.

Calculations

1. epl Br-i+I-^ 1 1 '000 . X^X (ml,-ml blank) 

r ml sample 6 v

o 2. epl ITT MOO vx°-01 ^/ i i ui i\ I-1 = r-2 r X -£- X (ml/-i ml blank) 

r ml sample 6

3. ppm Br-1 =

where epl= equivalents per liter,

ml*=titrant volume for combined determination, 

ml/-i=titrant volume for I"1 determination.

Report bromide concentrations of < 10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.

Preparation of reagents

Sodium chloride crystals: In addition to conforming to American Chemical 

Society specifications, these shall also be free from I"1 , IOT1 , Br"1 , and 

BrOs-1. The NaCl can be tested for ICV1 and BrOr1 by dissolving about 0.1 

g in 5 ml water, acidifying with 1 or 2 drops cone H2 SO4 (sp gr 1.84), and 

adding 2-3 ml starch solution. Immediate appearance of blue color indicates 

presence of ICV1 or BrCV1 , slow color formation is caused by atmospheric 

oxidation.

Methyl red indicator, 0.01 percent : Dissolve 0.01 g water-soluble methyl red in 

approx 100 ml water.

Hydrochloric acid, 50 percent v/v : Mix 50 ml cone HC1 (sp gr 1.19) with water 

and dilute to approx 100 ml.

Potassium hypochlorite, 4.4 percent : Dissolve 6.2 g KOH in approx 100 ml water, 

then saturate the solution with bromine-free CU while continually cooling 

and stirring. Store in glass-stoppered actinic-glass bottle. Prepare fresh 

daily.

Sodium formate, 50 percent : Dissolve .~>0 R XaCIlO, in hot water and dilute to 

approx 100 ml. Prepare fresh daily.

Sodium molybdate, 0.85 percent : Dissolve 1.0 g Na=MoO*2H2O in water and dilute 

to approx 100 ml.

Potassium fiuoride crystals, KF-2H2O.
124 ANALYTICAL PROCEDURES D:7a-2

Potassium iodide crystals, lOs^-f ree: The KI can be tested for ICV1 by dissolving 

about 0.1 g in 5 ml water, acidifying with 1 or 2 drops cone H2SC>4, and adding 

2-3 ml starch indicator. Immediate appearance of a blue color indicates 

the presence of ICV1 ; slow color formation is due to atmospheric oxidation.

Sulfuric acid, 20 percent v/v: Mix 200 ml cone H2SO4 (sp gr 1.84) with water 

and dilute to approx 1 liter.

Sodium thiosulfate solution, O.OION: Dilute 100.0 ml O.IN Xa2 S2O3 to 950 ml with 

carbon dioxide-free water and standardize against KlO3 as follows: Dry 

approx 0.5 g KIO3 for 2 hr at 180°C. Dissolve 0.3567 g in water and dilute 

to 1,000 ml. Pipet 25.00 ml of the K1O3 into a 250-ml iodine flask, then 

add successively 75 ml water and 0.5 g KI crystals. After solution is complete add 10 ml 20 percent H2 SO4 . Allow the stoppered flask to stand 5 

min in the dark, then titrate with Xa2 S2O3 using 2 ml starch indicator as 

the end point is approached (light straw color).

Normality of Na2S2O3 =;ml Na2S2O3

Sodium thiosulfate, 0.1^T : Dissolve 25.0 g Xa2 S2O3-5H2O in carbon dioxide-free

water, add 1 g Na2CO3 , and dilute to 1,000 ml. 

Sodium acetate solution, 16.5 percent : Dissolve 273.5 g NaC2H3O2-3H2O in water

and dilute to 1,000 ml. 

Acetic acid, 12.5 percent v/v: Mix 125 ml glacial HC2H3O2 (sp gr 1.049) with

water and dilute to 1,000 ml. 

Bromine water, saturated : Add to approx 250 ml of water slightly more liquid

Br2 than will dissolve when shaken. Store in a glass-stoppered actinicglass bottle.

D:7a-2 BROMINE DISPLACEMENT METHOD

The bromine-displacement method is only semiquantitative and finds 

maximum utility as a screening or field test.

Principle of determination

Chlorine water oxidizes the bromide to elemental bromine, which is 

concentrated and extracted with carbon disulfide or carbon tetrachloride(ASTM,1954,p.93).

The amount of color developed in the "'lobule of carbon disulfide 

( or tetrachloride) depends on the relation between the amount of chlorine water added and the amount of bromide present. For each concentration of bromide, there is an optimum amount of chlorine water 

that develops the deepest color. Therefore, to develop the deepest 

color in the unknown, it is necessary to add the chlorine water stepwise, 

0.10 ml at a time, until the maximum color in the samples lias been 

developed. The point of maximum color development is best detected 

if duplicate sets of unknowns and standards are prepared and one 

set is kept one step behind the other.
D:7ar-2 BROMIDE 125

Apparatus and reagents

Potassium bromide solution, 1.0 ml 1.0 mg Br~l 

Carbon disulflde or carbon tetraehloride 

Chlorine water

Procedure

1. Measure 10 ml of the sample into a test tube.

2. Prepare stepwise standards containing, in a 10.0-ml volume, from 0.5-4.0 mg 

of Br"1 as KBr (1.00 ml KBr solution 1.00 mg Br"1 ), in increments of 0.5 mg.

3. Add 1.0 ml CS2 or CCh and 1 ml chlorine water to liberate the bromine.

4. Shake well.

-"». After 15 min compare the color of the globule of CS2 or CCU with that of the 

globules in standards.

Calculations

ppm Bi-1 ^ 1)OUQ Xl.Oxml standard 

ml sample

Preparation of reagents

Chlorine water: Saturate a small volume of water with C12 .

Potassium bromide solution, 1.0 ml=1.0 mg Br"1 : Dissolve 1.49 g KBr in water 

and dilute to 1,000 ml.

REFERENCES

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

Kolthoff, I. M., and Sandell, B. B., 1952, Textbook of quantitative inorganic

analysis: New York, MacMillan Co.

D:8 CALCIUM

Calcium is dissolved from practically all rocks but is usually 

found in greater quantities in waters leaching deposits of limestone, 

dolomite, gypsum, or gypsiferous shale. Waters associated with 

granite or silicious sand may contain less than 10 ppm of calcium. 

Many waters from limestone areas contain 30 to 100 ppm, and 

waters that traverse gypsiferous shale may contain several hundred ppm.

Calcium imparts the property of hardness to water and when 

present with alkalinity or sulfate may cause boiler scale. A little 

calcium carbonate is desirable in water used domestically because 

of the protective coating that it may form in the pipes. A high 

ratio of calcium to sodium is desirable in water used for irrigation because calcium flocculates the soil colloids and tends to maintain good soil structure and permeability.

D:8a-l COMPLEXOMETRIC METHOD

The complexometric method is similar in substance to part IC.4, 

APHA (1955, p. 118) Standard Methods; D 1126-53 T, ASTM 

(1954, p. 254) Manual on Industrial Water; and method 7, U.S. 

Salinity Laboratory Staff (1954, p. 94) Handbook 60.

The method is applicable to most natural waters but may fail in 

the analysis of brines or some acid or polluted waters that contain 

excessive amounts of heavy metals.

Principle of determination

Disodium dihydrogen ethylenediamine tetraacetate (Na2EDTA) 

forms a slightly ionized, colorless, stable complex with calcium 

ions. Murexide is dark purple in the absence of calcium, but with 

calcium forms a light-salmon-colored complex which has an ionization constant higher than the Na2EDTA complex. Hence, by using 

murexide as an indicator, a solution containing calcium ions may

o o

II II NaO C CH2 CHj C ONa

N CH2 CHj N +Ca«

HO-C CH2 CH2 C-OH

4 I *

NaO C CH2 CH2-C ONa 

O N CHj-CHr-N O

0^\ V^Xj2 \ / /\H v^ .tJ 2 +2-r~Zfi H«£

\ Ca /

\_</ \_c/

II II O O

127
128 ANALYTICAL PROCEDURES

be titrated with Na2EDTA. The optimum pH of the titration is 

10.4 or above.

The salt, Na2EDTA, reacts with iron, manganese, copper, zinc, 

lead, cobalt, nickel, barium, strontium, calcium, magnesium, and 

several other metals. Murexide reacts with strontium but not with 

magnesium or barium; however, the end point in the presence of 

strontium is sluggish, and the titration is not strictly stoichiometric. 

Barium does not titrate as calcium but affects the indicator in some 

unknown way so that no end point, or at best a poor end point is 

obtained. Barium can be removed by prior precipitation with sulfuric acid. The interference of heavy metals is minimized by the 

addition of hydroxylamine and cyanide, which reduce and (or) 

complex the metals. Concentrations of 5 ppm of iron and 10 ppm 

of manganese can be tolerated.

The interference of heavy metals is relatively easy to detect 

because of the typical end point. Conventional methods of hydroxide 

and sulfide treatment can be used, if necessary, to remove these 

metals from solution before titration. Magnesium in high concentrations may precipitate as magnesium hydroxide, but the precipitation is not significant unless the photometric-titration assembly 

is used to determine the end point.

The results are generally accurate and reproducible to ±0.025 

mg in the 1-mg range and to ±0.05 mg in higher concentrations.

Additional information on the principle of determination is given 

by Banks (1952, p. 484).

Apparatus and reagents

Visual titration apparatus consisting of a motor-driven stirrer, 25-ml buret, 

white-porcelain-base buret holder, and shaded incandescent lamp. The 

sample beaker is placed near the front of the porcelain base and the reaction is viewed diagonally downward through the side of the beaker and 

against the white background. Illumination is from the side and at about 

90° from the line of vision.

Photometric-titration assembly, optional (see sec. C: 2h)

Buret, 25-ml

Hydroxylamine hydrochloride, 3 percent

Sodium hydroxide, 2.2V

Sodium cyanide, 2.5 percent

Murexide, dry mixture

NaaEDTA, 1.00 ml=0.50 mg Ca+2

Procedure

1. Pipet a volume of sample containing less than 10 mg Ca"2 (50.0 ml max.) into 

a 150-ml beaker and adjust the volume to approx 50 ml.

2. Insert beaker in titration assembly and start stirrer.

3. Add 1 ml 3 percent NH2OH-HC1.

4. Add 1 ml 2N NaOH.
D:8a-2 CALCIUM 129

5. Add 1 ml 2.5 percent NaCN (CAUTION deadly poison).

6. Add 0.2 g murexide indicator and proceed immediately with the titration.

7. Titrate with Na2EDTA (1.00 mloO.50 mg Ca+2 ) until purple swirls begin to 

show. The end point is reached when the sample color changes from salmon 

to orchid purple. At the end point, the addition of a small increment of 

titrant will not cause a deepening of the purple color.

8. Determine a blank correction by similarly treating 50 ml metal-free water. 

The normal blank correction is 0.05 or 0.10 ml.

Calculations

ppm Ca+2 = 1 .. X 1 '000 x(ml titrant-ml blank) X0.5 density ml sample

Report calcium concentrations of <10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.

Preparation of reagents

Hydroxylamine hydrochloride, 3 percent: Dissolve 30 g NH2OH-H 1 in metalfree water and dilute to approx 1 liter. 

Sodium hydroxide, 2N: Dissolve 80 g NaOH in 800 ml water. Cool and dilute

to approx 1 liter. 

Sodium cyanide, 2.5 percent (CAUTION NaCN is a deadly poison and the

reagent must be so marked.) Dissolve 25 g NaCN in water and dilute to

approx 1 liter. 

Murexide, dry mixture : Mix thoroughly 1.0 g ammonium purpurate with 200 g

sucrose. Eastman reagent-grade product has been found to be satisfactory.

Provide the bottle with a dispensing spoon of 0.2-g capacity. 

Na2EDTA, 1.00 mloO.50 mg Ca+2 : Dissolve 4.65 g Na2EDTA, which has been

dried overnight in an H2 SO« desiccator, in metal-free water and dilute to

1,000 ml. (The titrant is stable for several months and is usually prepared

in larger quantities.) Standardize the titrant by titrating 25.00-ml CaCU

solutions (1.00 ml=0.40 mg Ca+2 ) as described in the procedure for sample

analysis. 

Calcium chloride, 1.00 ml=0.40 mg Ca+2 : Suspend 1.000 g CaCo», dried at 180°C

for 1.0 hr before weighing, in approx 600 ml metal-free water and dissolve

cautiously with a minimum of dilute HC1. Dilute to 1,000 ml.

D:8a-2 PERMANGANIMETRIC METHOD

The permanganimetric method is similiar in substance to part IB, 

APHA (1955, p. 51-52) Standard Methods.

Principle of determination

Calcium is precipitated as the oxalate ; the calcium oxalate is then 

dissolved and titrated in acid media with a standard solution of 

potassium permanganate. The following reactions are involved :

5C.,Or2 +2MnOr1 + 16H+1-> 2Mn+2+ 10CO2 +8H2O

The presence of magnesium increases the solubility of calcium oxalate 

appreciably, as a consequence of the formation of a complex mag-
130 ANALYTICAL PROCEDURES D:8a^-2

nesium oxalate which withdraws oxalate ions. The effect may be 

counteracted by the addition of a large excess of oxalate. Heavy 

metals interfere unless removed by hydroxide and (or) sulfide precipitation (see sec. D:8a-3). Barium and strontium also form insoluble oxalates. Calcium can be separated from small quantities of 

barium (3-4 nig or more) by double precipitation, but strontium 

interference cannot be eliminated in this manner. Precipitation of 

calcium oxalate from acid solution followed by hot digestion minimizes 

the effects of most coprecipitation and provides a precipitate of nearly 

theoretical composition.

The first few drops of permanganate react slowly with oxalic acid, 

but after a small amount of manganous salt has been formed the reaction proceeds almost instantaneously. If the water is low in calcium 

the initial ^increments of permanganate should be added with care 

to prevent overshooting the end point. In hot acid solutions permanganate is slowly decomposed with the evolution of oxygen; therefore, too rapid titration with insufficient stirring may cause an error. 

In addition, the permanganate tint at the end point can be expected 

to fade in time.

Results are generally accurate and reproducible to ±0.1 mg.

Apparatus and reagents

Steam bath

Hotplate

Buret, 50-ml

Methyl orange indicator solution, 0.05 percent

Hydrochloric acid, cone (sp gr 1.19)

Ammonium oxalate, 5 percent

Ammonium hydroxide, cone (sp gr 0.900)

Ammonium hydroxide, 1 percent v/v

Sulf uric acid, 25 percent v/v

Potassium permanganate, 0.0499#, 1.00 mlol.OO mg Ca1

Procedure

1. Pipet a volume of sample containing less than 40 mg Ca (100.0 ml max) into 

a 250-ml beaker and adjust the volume to approx 100 ml.

2. Acidify to methyl orange with cone HC1.

3. Heat the solution nearly to boiling. While hot, add 10 ml 5 percent 

(NH4 ) 2C2O4 slowly and with constant stirring.

4. Add cone NH4OH with constant stirring until CaC2C>4 begins to precipitate. 

Do not make the solution alkaline at this point.

5. While the mixture is still acid to methyl orange, digest for 0.5 hr on a steam 

bath.

6. Slowly add cone NH4OH to make the solution just alkaline to methyl orange 

and then 2-5 drops in excess.

7. Allow the precipitate to digest on the steam bath for another 2 hr.
D:8a-3 CALCIUM 131

8. By decantation rinse the beaker and precipitate at least 4 times with 5- to 

10-ml portions of hot 1 percent NH4OH, pouring the rinses through Whatman 

No. 40 filter pai>er.

9. Carefully unfold the filter paper and quantitatively wash any precipitate back 

into the original beaker with a minimum of hot water.

10. Dissolve and rinse any remaining precipitate on the filter paper into the 

beaker with 2 successive 10-ml portions of warm 25 percent H2SO4. Rinse 

the paper 3 times with hot water and collect the rinsings in the original 

beaker.

11. Heat the solution 55°-60°C and, while hot, titrate to the first pink tint that 

persists for 30 sec.

Calculations

ppm Ca+s= , 1 ., X 1 1 '°00 . Xml titrant density ml sample

Report calcium concentrations of<10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 significant figures only.

Preparation of reagents

Methyl orange indicator solution 0.05 percent: Dissolve 0.50 g methyl orange 

in water and dilute to approx 1 liter.

Ammonium oxalate, 5 percent: Dissolve 50 g (NHOsCsO* in water and dilute 

to approx 1 liter.

Ammonium hydroxide, 1 percent v/v: Mix 10 ml cone NHiOH (sp gr 0.900) with 

water and dilute to approx 1 liter.

Sulfuric acid, 25 percent v/v: Mix 250 ml cone H2SO4 (sp gr 1.84) with water 

and dilute to approx 1 liter.

Potassium permanganate, 0.04992^, 1.00 mlol.OO mg Ca+2 : Heat 500 ml KMnGu 

stock solution to boiling and filter it through asbestos fiber or fritted glass. 

Dilute to 950 ml and standardize against primary standard Na2C2 04 as 

follows: Dry 0.5 g of the salt at 105 °C for 1 hr. Dissolve approx 0.1 g, accurately weighed to 0.1 mg, in 50 ml water and 20 ml 25 percent H»SO4.

ml i KMn0 T * r\ j m 4 required 

Add, from a buret, 90 percent of required KMnO4 to the oxalate solution and 

heat to 5o°-60°C. Continue the titration slowly and with constant stirring 

until a faint pink color persists for 30 sec. For practical work, no blank 

correction is required. Store the standardized KMnC-4 solution in a lightproof bottle.

Potassium permanganate, stock solution: Dissolve 3.25 g KMnO4 in water and 

dilute to 1,000 ml. Store the solution in a dark place for at least 1 week to 

permit precipitation of manganous salts.

D:8a-3 GRAVIMETRIC METHOD

The gravimetric method is similar in substance to D 511-52, ASTM 

(1954, p. 245) Manual on Industrial Water.

This method is normally used in the analysis of brines, some acid 

waters, and concentrated industrial wastes.
132 ANALYTICAL PROCEDURES D:8a-3

Principle of determination

Calcium is precipitated as calcium oxalate, ignited to calcium oxide, 

and weighed. The following reactions are involved:

The possible error caused by coprecipitation of silica is eliminated 

if the calcium is determined on the filtrate from the gravimetric 

silica determination.

The presence of magnesium increases the solubility of calcium 

oxalate very appreciably as a consequence of the formation of a 

complex of magnesium oxalate which withdraws oxalate ions. The 

effect may be counteracted by the addition of a large excess of 

oxalate. Heavy-metal ions which also form slightly soluble salts 

are removed by precipitation as the hydroxides. Manganese is precipitated as the sulfide. When only a few milligrams of barium 

is present in the sample volume, double precipitation suffices to 

separate the calcium from it. Strontium also precipitates as the 

oxalate, and its interference is not eliminated by double or triple 

precipitation. Strontium is frequently present in brines, and the 

completeness of its precipitation as strontium oxalate requires further study. Probably the strontium should be subtracted from the 

gravi metrically determined calcium value, but the Geological Survey has not yet found a completely satisfactory method for determining strontium in all types of water. Until such a method is 

developed, the reported value for calcium may also include some 

strontium. Recognized coprecipitations with calcium oxalate are 

minimized by diluting a sample about fifteenfold. Interferences 

in some brines and industrial wastes may present special problems 

of separation. The analyst is referred to Hillebrand and Lundell 

(1929, p. 486-505) for detailed discussion of chemistry and technique of the determination.

Calcium oxide is very hygroscopic and special attention should 

be given to the desiccant over which the ignited residue is cooled. 

Freshly charged alminum oxide, freshly ignited calcium oxide, or 

concentrated sulfuric acid are satisfactory drying agents. Calcium 

oxide is a better drying agent than calcium chloride, hence the latter 

is unsatisfactory. To insure a completely anhydrous residue, the 

calcium oxide is weighed, ignited, and reweighed, until 2 successive 

weighings are comparable.

If suitable precautions are taken to minimize the effect of interferences, the results may be accurate and reproducible to ±0.5 mg.
D:8a-3 CALCIUM 133

Apparatus and reagents

Hotplate

Platinum crucible 

Muffle furnace, 1,000° C 

Bromthymol blue, 1 percent 

Ammonium hydroxide, cone (sp gr 0.900) 

Ammonium sulflde, 20 percent 

Hydrochloric acid, cone (sp gr 1.19) 

Ammonium oxalate, 4 percent 

Ammonium hydroxide, 50 percent v/v 

Ammonium oxalate, 0.1 percent 

Hydrochloric acid, 20 percent v/v

Procedure

1. Dilute or concentrate the nitrate from the gravimetric silica determination 

(see sec. D: 34a-2, procedure 8) to a convenient volume. Pipet an aliquot 

containing less than 250 mg Ca into a 1,000 ml beaker and adjust the volume 

to approximately 800 ml.

2. Add 2 drops 1 percent bromthymol blue.

3. Neutralize the solution to the green color with cone NHtOH. Some silica may 

precipitate along with the iron and aluminum.

4. Add 1 ml 20 percent (XH4 ) a S to precipitate the manganese and allow the mixture to stand 1-2 hr.

5. Filter off the precipitates through Whatman No. 40 filter paper, retaining the 

nitrate.

6. Wash the beaker and precipitate several times with distilled water that contains a few drops of cone NH4OH.

7. Discard the residue and combine the filtrate and washings.

8. Acidify the filtrate with cone HC1 and add a 10-drop excess.

9. Warm on a hotplate but do not bring solution to a boil.

10. While hot, add with constant stirring approx 100 ml warm 4 percent 

(NH4 ).C204.

11. Add with constant stirring 50 percent NH4OH dropwise until the color changes 

from yellow to blue.

12. Remove the mixture from the hotplate and allow it to digest for 1 hr.

13. Quantitatively collect the precipitate on Whatman No. 40 filter paper and 

rinse the beaker and precipitate 4 times with cold 0.1 percent (NH4 ) 2C2O4 . 

Retain the combined filtrate and washings for the gravimetric barium and 

magnesium determinations.

14. Carefully unfold the filter paper and wash the precipitate back into the original beaker with hot water. Dissolve any precipitate remaining on the paper 

with 50 ml 20 percent HC1, collecting the washings in the original beaker. 

Rinse the paper 3 times with hot water, collecting the rinsings in the original 

beaker. When dissolving the precipitate give particular attention to the 

folds in the filter paper.

! ">. Repeat steps 9 through 13.

16. Wash the precipitate 3 times with cold water containing 3 drops cone NH4OH.

17. Slowly ignite the precipitate in a tared platinum crucible over a low oxidizing flame until the filter paper is reduced to white ash.

18. Transfer the crucible to the muffle furnace and ignite at approx 1,000°C 

for 0.5 hr or until the precipitate is white.
134 ANALYTICAL PROCEDURES D:8a-4

19. Cool the crucible in a desiccator and weigh. Ignite and weigh to constant 

weight.

Calculations

ppm Ca+2= * X I'000 X mg CaOXO.7147 density ml sample

Report calcium concentrations of <999 ppm to whole numbers and 

of > 999 ppm to 3 significant figures only.

Preparation of reagents

Bromthymol blue, 1 percent: Triturate 1.0 g bromthymol blue with 1.6 ml O.IN

NaOH and dilute to approx 100 ml with water. 

Ammonium sulfide, 20 percent: Dissolve 20 g (NH4 ) 2 S ,in water and dilute to

approx 100 ml. 

Ammonium oxalate, 4 percent: Dissolve 40 g (XH4)^C2O4 in water and dilute to

approx 1 liter. 

Ammonium hydroxide, 50 percent v/v: Mix 50 ml cone NH4OH (sp gr 0.900)

with water and dilute to approx 100 ml. 

Ammonium oxalate, 0.1 percent: Dissolve 1 g (NHO-CaO* in water and dilute

to approx 1 liter. 

Hydrochloric acid, 20 percent v/v: Mix 200 ml cone HC1 (sp gr 1.19) with water

and dilute to approx 1 liter.

D:8a-4 TURBIDIMETRIC METHOD

The turbidimetric estimation of calcium is a semiquantitative test 

useful in fieldwork and in selecting the proper sample volume for 

quantitative determinations.

Principle of determination

The turbidity of precipitated calcium oxalate is compared with the 

turbidity of standard calcium solutions similarly treated.

Apparatus and reagents

Test tubes, 10-ml graduations 

Calcium chloride, 1.0 ml=0.1 mg Oa+"' 

Acetic acid, 50 percent v/v 

Potassium oxalate, 5 percent

Procedure

1. Pipet a volume of sample containing less than 0.2 mg Ca (10.0 ml max) into 

a graduated test tube and adjust the volume to 10 ml.

2. Prepare a blank and sufficient standards and adjust the volumes to 10 ml.

3. Add 1 ml 50 percent HC2H3O2.

4. Add 1 ml 5 percent K2C2O4 and mix well.

5. Allow the mixture to stand for 10 min : then resuspend the precipitate.

6. Compare the turbidity of the sample with the turbidities of the standards.

Calculations

ppm Ca+2= ~ Xmg Ca standard 

ml sample
D:8a-4 CALCIUM 135

Preparation of reagents

Calcium chloride, 1.0 ml=0.1 mg Ca+2 : Suspend 0.2497 g CaCOs in 600 ml water.

Dissolve with a minimum of dilute HC1 and dilute to 1,000 ml. 

Acetic acid, 50 percent v/v: Mix 50 ml glacial HC2H3O2 (sp gr 1.049) with water

and dilute to approx 100 ml. 

Potassium oxalate, 5 percent: Dissolve 5.0 g K2C2O4 in water and dilute to

approx 100 ml.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 14&-A. 

Banks, J., 1952, The volumetric determination of calcium and magnesium by

the ethylenediamine tetra-acetate method: Analyst, v. 77. 

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis:

New York, John Wiley and Sons, Inc. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline

and alkali soils: U.S. Dept. Agriculture, Agriculture Handb. 60.

D:9 CARBON DIOXIDE

Carbon dioxide is a respiration product of aquatic plants and 

animals and is one of the byproducts of aerobic and anaerobic decomposition of organic matter. Acids from natural sources or pollution 

liberate carbon dioxide from bicarbonate. Most of the carbon 

dioxide in water is from these sources. Air contains only about 

0.04 percent carbon dioxide, and the concentration in water in equilibrium with air will approach O.T ppm. Streams are normally 

rather low in carbon dioxide, but some ground waters contain large 

quantities in their natural environment.

Carbon dioxide in drinking water is not physiologically important to man and livestock, but it has a marked effect on fish. Tolerances differ widely with the species, dissolved oxygen, and other 

chemical properties of the water. Of waters of the United States 

that support good fish fauna, 50 percent have less than 1.5 ppm 

and 95 percent have less than 5 ppm (California State Water Pollution Control Board, 1952, 1954, p. 205-206). Free carbon dioxide 

contributes to the corrosiveness of the water and is likely to damage 

calcareous building material such as cement. It has been recommended that concrete be coated if it is in contact with water containing 20 ppm carbon dioxide (Terzaghi, 1949, p. 136).

D:9a-l CALCULATION METHOD

The calculation method is applicable to most unpolluted water 

whose dissolved-solids content does not exceed 500 ppm, or at the 

most 800 ppm. References are given to ways of extending the 

range of application somewhat. Carbon dioxide should not be 

calculated for waters whose pH is not essentially a function of the 

carbonate-carbon dioxide system only; this would probably exclude 

some highly colored waters and others containing appreciable quantities of organic or inorganic acids from natural or cultural sources.

Principle of determination

Equations for the calculation of carbon dioxide are derived from 

3 well-known mass-law equations and 1 stoichiometric equation as 

follows :

[H+J]XHC03 -

H2CO3 

[H+']XC03 -2

-XK,'=4. 54X10-7 (1)

[H+J] X [OH-1 ] = # ,'= 1.00 X10-" (3)

137
138 ANALYTICAL PROCEDURES D:9a-l

KI and KZ are apparent constants applicable for normal water having 

a dissolved-solids concentration of less than 500 ppm.

Alkalinity+ [H+1 ] =2C<VS+HCO3-1+ [OH'1 ] (4)

where alkalinity is titratable equivalents of base per liter (to pH 4.5). 

From equation 2:

ro -2 _^2> XHCO8~1 .,.) 

3 [H+i] () 

From equations 3 and 4:

Alkalinity+ [H+1 ]-HCO3-1 - K"

Equating equations 5 and 6:

Alkalinity - HCO3-' + [H+1 ] -:

2 ~ [H+i] 

Solving for HCCV1 :

and

alkalinity + [H+1 ] 

1 I ^ 

From equation 1:

alkalinity + [R+l]

rw K

l

 V ' *>V ' 

1 1 I ^A2 ^~[H+']

In order that CO2 may be expressed in parts per million: 

ppm CO2 =moles H2CO3 X44,011 (or 4.4XH)4)

equiv. per liter alkalinity =^7-7^0 X ppm alkalinity as bijUiy

= 1.64X10-5 alkalinity as HCOr1 

Evaluating the constants and expressing alkalinity as ppm

PPm alkalinity as HCOs-1 X 1.64X lQ

(8)

H.22X10-

ppm alkalinity as HCO3-i X 1 .64 X 10-5 + [H+1 ] - 

= 9.70X10'o[H+i]X

11 ' (10)
D:9a-l CARBON DIOXIDE 139

The values for constants K^', K2 ', and Kw' are affected by other dissolved minerals and become greater as the total salt concentrations 

increase because of the diminution of activities of the various ions 

(Moore, 1939, p. 51-65) . The numerical values of the constants given 

are satisfactory for practical purposes for waters having less than 

800 ppm of dissolved solids ; 500 ppm has been recommended by some 

investigators. To compute apparent constants from the constants at 

zero ionic strength (the values given in handbooks), the analyst is 

referred to the works of Langelier ( 1936, p. 1500-1521) and Larson and 

Buswell (1942, p. 1667-1678).

Between pH 6 and 9 some of the terms of equation 10 become insignificant and can be eliminated for practical purposes. The [H+1] in 

the numerator of the second half of the equation becomes relatively 

insignificant above a pH of about 5. The 10~14/[H+1 ] is also insignificant at a pH of less than 9 ; however, the possible carbon dioxide concentration at a pH of 9 and higher is so small that little error is 

incurred by eliminating this term from any computations. As pH 

decreases, the denominator of the second half of the equation approaches 1 and is 1 for all practical purposes in the acid and neutral 

range. The 2 Kz' /[H+1] term becomes significant at elevated pH's, 

and its elimination would result in a rather high percentage error at 

and above about pH 9. Nevertheless, it can be eliminated because the 

absolute error (ppm) is very small. Therefore, in a pH range of 6 to 

9, equation 10 may be simplified to :

ppm CO2=j -7[H+»]X1.64X10-5 ppm alkalinity as HCOa"1

4.04: /\ 1U

= 1.589X108 [H+i]Xppm alkalinity as HCOa-1

Additional information on the principle of the determination is 

given by De Martini (1938, p. 85) .

Apparatus and reagents

Titration assembly consisting of pH meter and medium-speed stirrer 

Sulfuric acid, 0.01639#: 1.00 mlol.OO mg HCOa'1

Procedure

Samples for the determination of carbon dioxide should be collected 

as directed in sec. A : 3c, and the determination should be made as soon 

after collection as possible. (See discussion of changes of pH-alkalinity-carbon dioxide balance in sec. B : 1.) .

1. Pipet a volume of sample containing less than 40 mg alkalinity as HCOr1 (50.0 

ml max) into a suitable beaker.

2. Insert the beaker in the titration assembly and record the pH to the nearest 0.1.

3. Start the stirrer and titrate immediately to pH 4.5.
140 ANALYTICAL PROCEDURES

-x 1,000 -Xml titrant

Calculations

1. ppm Alkalinity as HCO3~1= ;j- -± ^ i i '

density ml sample

2. ppm CO2 can be calculated directly from the complete equation 10 under 

"Principle of determination". Between pH 6.0 and 9.0:

ppm CO2 =1.589X106 [H+1 ] Xppm alkalinity as HCCV1.

For convenience a table has been prepared in which the term 1.589X10" [H+1 ] 

has been evaluated for each 0.1 pH unit between 6.0 and 9.0

PH

6.0 

6.1- 

6.2 

0.3 _

6.4 

6.5 

6.6 

6.7 

6.8 

6.9 

1.589X10»[H-"]

1. 589

1. 262

1.003

. 796

. 633

. 503

. 399

.317

. 252

. 200

pH

7.0 

7.1 _

7.2 _

7.3 -

7.4 -

7.5 -

7.6 _

7.7 

7.8 

7.9 _

1.589X10«[H+i]

0. 159

. 126

. 100

.080

.063

.050

. 040

.032

.025

. 020

pH

8.0 ___-

8.1 

8.2

8.3 

8.4 _

8.5 

8.6 __-_

8.7 _

8.8- 

8.9 

9.0 _ 

1.589X10«[H-H]

0.016

.013

. 010

. 008

.006

.005

.004

.003

.003

. 002

.002

Report carbon dioxide concentrations of <10 ppm to 1 decimal 

place, of between 10 and 999 ppm to whole numbers, and of >999 

ppm to 3 significant figures only.

Preparation of reagents

Sulfuric acid, 0.01639#, 1.00 mlol.OO mg HCOr1 : Add 0.5 ml cone H2SO4 (sp gr 

1.84) to 950 ml water. (The titrant is stable for several months if protected 

from ammonia fumes and is usually prepared in larger quantities.) After 

the solution has been thoroughly mixed, standardize it by titrating 25.00 

mlNa^COs (1.00 mlol.OO mgHCCV1 ) topH4.5.

Sodium carbonate, 1.00 mlol.OO mg HCOr1 : Dissolve 0.8686 g primary standard 

Na2CO3 in carbon dioxide-free water and dilute to 1,000 ml.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1.

De Martini, F. E., 1938, Corrosion and the Langelier calcium carbonate saturation index: Am. Water Works Assoc. Jour., v. 30.

Langelier, W. F., 1936, The analytical control of anti-corrosion water treatment: 

Am. Water Works Assoc. Jour., v. 28.

Larson, T. E., and Buswell, A. N., 1942, Calcium carbonate saturation and 

alkalinity interpretations: Am. Water Works Assoc. Jour., v. 34.

Moore, E. W., 1939, Graphic determination of carbon dioxide and three forms 

of alkalinity: Am. Water Works Assoc. Jour., v. 31.

Terzaghi, R. D., 1949, Concrete deterioration due to carbonic acid: Boston 

Soc. Civil Engineers Jour., v. 36.
D:10 CHLORIDE

Most naturally occurring chlorides are very soluble. Chloride 

concentration in natural waters ranges widely from less than 1 ppm 

in some waters to many thousand ppm in some brine. Chloride is 

the major anion in most brines of the United States. The discharge 

of some industrial wastes into streams or ground-water reservoirs 

may considerably increase their chloride content. Human and animal excreta are high in chloride and nitrogenous material. In 

water supplies, the presence of abnormal concentrations of the two 

together is indicative of possible pollution by human or animal 

wastes.

A high concentration of chloride imparts a salty taste to the 

water but the threshold of detection A^aries with individuals. Although for drinking purposes, water with a chloride content of 

1,000 ppm may be physiologically safe, the U.S. Public Health 

Service (1946) recommends that the concentration not exceed 250 

ppm in water on carriers subject to Federal quarantine regulations.

Chlorides may accelerate corrosion - in pipes, boilers, and other 

fixtures (Taylor, 1949). Magnesium chloride when heated releases 

hydrochloric acid which is highly corrosive. Many crops may be 

injured by waters containing excessive quantities of chloride. Chloride is generally about twice as toxic to crops as sulfate (Eaton, 

1954).

D:10a-l VOLUMETRIC METHOD

The volumetric method is similar in substance to part IB, APHA 

(1955, p. 60) Standard Methods, and D 512-49, ASTM (1954, p. 

220-222) Manual on Industrial Water.

The procedure is recommended for waters whose chloride concentration is less than 2,000 ppm and can be used satisfactorily for 

measuring chloride concentrations up to 5,000 ppm.

Principle of determination

In the well-known Mohr method for determination of chloride, 

use is made of the fact that in the titration of sodium chloride with 

silver nitrate, the solution is saturated with silver chloride at the 

equivalence point and contains equal concentrations of silver and 

chloride ions. Addition of an excess of silver precipitates silver 

chloride, whose solubility decreases with an excess of either silver 

or chloride ions. When potassium chromate is used as an indicator, the chromate ion combines with the excess silver to form very 

slightly soluble red silver chromate. The following reactions occur:

Ag+' + Cl-^AgCl 

2 Ag+H Cr04 -2 -+Ag2 CrO4

141
142 ANALYTICAL PROCEDURES D:10a-l

The pH for the titration should be between 7.0 and 10.5. In an acid 

medium, the sensitivity of the method is decreased; the second ionization constant of chromic acid is small, and therefore the chromate ion 

reacts with hydrogen ions.

CrOr'+H+^HCrOr1

The solution should not be too alkaline because silver hydroxide might 

then precipitate before the silver chromate (Collins, 1928). Calcium 

carbonate can be used to adjust the pH of acid waters without danger 

of making the solution too alkaline. Detection of the end point is 

facilitated by illuminating the titration with yellow light or by viewing the titration through yellow goggles or a filter.

Iodide and bromide titrate stoichiometrically as chloride. Phosphate, sulfide, and cyanide interfere. Sulfide and cyanide can be 

removed by acidifying and boiling the sample, then adjusting the pH 

with calcium carbonate. Hydrogen sulfide can often be removed 

simply by passing pure air through the sample. Sulfite interferes but 

can be oxidized readily to sulf ate with hydrogen peroxide.

Two strengths of silver nitrate are provided. The dilute titrant 

is recommended if the chloride concentration is less than 5 mg in a 

50- or 25-ml sample. The end point with the dilute silver nitrate is 

not as sharp as with the concentrated titrant; therefore, the concentrated silver nitrate is recommended if the chlpride concentration is 

greater than 200 ppm. In high-chloride waters, the voluminous precipitate masks the end point, and the maximum chloride concentration that can be titrated satisfactorily is about 50 mg. Sample dilution can be carried only so far before the dilution factor decreases 

the precision and accuracy considerably. If a 10-ml sample were 

taken, 0.1 ml silver nitrate (1.00 ml=c=5.00 mg CH) is equivalent to 

50 ppm of chloride. Sample volumes less than 10 ml are not 

recommended.

The results are usually accurate and reproducible to ±0.05 mg when 

the dilute silver nitrate (1.00 mloO.50 mg CH) is used. When the 

concentrated silver nitrate (1.00 ml=c=5.00 mg CH) is used, accuracy 

and reproducibility of ±2-3 percent can be expected.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 303-309).

Apparatus and reagents

Yellow light (a yellow bulb or yellow filter is suitable)

Buret, 25-ml

Potassium chromate indicator solution

Silver nitrate, 1.00 mlc=5.00 mg Cl"1

Silver nitrate, 1.00 mloO.50 mg Cr1
D:10a-2 CHLORIDE 143

Procedure

1. Pipet a volume of sample containing less than 50 mg Cl"1 (50.0 ml max) into 

a porcelain evaporating dish and adjust the volume to approx 50 ml.

2. Add 10 drops K2CrO4 indicator.

3. With constant stirring, titrate with AgNOa until the pink-red Ag2CrO4 persists 

for 10-15 sec.

4. Determine a blank correction by similarly treating 50 ml dilution water. 

The normal blank correction with AgNO3 (1.00 mloO.50 mg Cl"1 ) is 0.05 or 

0.10 ml. No blank correction is required with the stronger titrant.

Calculations

ppm Gl-1 =, 1 .. X 1>00° x (ml titrant ml blank) X (mg Cl'1 per ml titrant) density ml sample

Keport chloride concentrations of <10 ppm to 1 decimal place of between 10 and 999 ppm to whole numbers and of >999 ppm to 3 significant figures only.

Preparation of reagents

Potassium chromate indicator solution: Dissolve 50 g KsCrO* in approx 1,000 ml 

water. Add AgNO3 (1.00 ml=c=0.50 mg Cl"1 ) until a small amount of red 

AgzCrO* precipitates. Put aside in dark for 24 hr and filter to remove the 

AgjCrO*.

Silver nitrate, 1.00 mloO.50 mg Cl"1 : Dilute 100 ml AgNO3 (1.00 mlo5.00 mg Cl'1 ) 

to 1,000 ml. Check the titer of the reagent by titrating 10.00 ml NaCl (1.00 

ml=1.00 mg Cr1 ). Store in a lightproof bottle.

Silver nitrate, 1.00 mlc=5.00 mg Cl"1 : Pulverize approximately 30 g AgNO3 crystals 

in a mortar and dry at 105°-120°C. Browning of the crystals indicates 

reduction to Ag° owing to the high heating temperatures or presence of 

impurities. Dissolve 23.96 g dried AgNO3 in water and dilute to 950 ml 

before standardizing. Standardize by titrating 25.00 ml NaCl (1.00 ml= 

1.00 mg CP) diluted to approx 50 ml. Store in a lightproof bottle.

Sodium chloride, 1.00 ml=1.00 mg Cl"1 : Fuse NaCl in a platinum dish and cool. 

Dissolve 1.6484 g of the fused NaCl in water and dilute to 1,000 ml.

D:10a-2 GRAVIMETRIC METHOD

The gravimetric procedure is recommended for industrial wastes 

that contain substances that interfere with the Mohr titration and 

for waters whose chloride content exceeds 5,000 ppm. Although 

the volumetric method can be used in the range of 2,000-5,000 ppm, 

results obtained by the gravimetric procedure are usually more 

accurate and reproducible than those obtained by the volumetric 

method.

Principle of determination

Chloride is precipitated as silver chloride. Best results are obtained if the precipitate weighs between 250 and 500 mg (approx 

65-125 mg Cl"1 ). The optimum silver nitrate concentration in the 

precipitation reaction is 0.05 g per liter. For routine analysis a
144 ANALYTICAL PROCEDURES D:10a-2

moderate excess of silver nitrate can be tolerated, but an unnecessarily large excess should be avoided.

Precipitation is complete almost immediately after addition c>f 

the silver nitrate, but digestion is needed to flocculate the colloidal 

silver chloride into a filterable mass. The presence of nitric acid 

and heat hastens the flocculation. When exposed to light, silver 

chloride slowly decomposes to silver and chlorine, the silver remaining dispersed in the precipitate and the chlorine escaping from the 

crystals. Excessive exposure to light during precipitation and 

digestion phases of the determination may also precipitate free 

silver from the silver nitrate. The precipitate collected on filter 

paper cannot be ignited satisfactorily because the carbon and reducing gases formed in burning off the paper reduce some of the silver 

chloride to free silver.

Anions that form silver salts insoluble in nitric acid, such as 

bromide, iodide, cyanide, and sulfide, also precipitate with the silver 

chloride. Cyanide and sulfide can be removed from the water by 

acidifying and boiling. High concentrations of heavy metals may 

interfere, but such concentrations are seldom found except in some 

industrial wastes.

The results are generally accurate and reproducible to ±0.5 mg.

Apparatus and reagents

Steam bath

Gooch crucible with asbestos- or glass-fiber mat

Suction filtration apparatus

Oven, 180°C

Nitric acid, cone (sp gr 1.42), CP-free

Nitric acid, 0.1 percent v/v

Silver nitrate, 5 percent

Procedure

1. Pipet a volume of sample containing between 65 and 125 mg Cl"1 into a 250-ml 

beaker and adjust the volume to approx 100 ml.

2. Add 5-6 drops cone HNO3 .

3. Heat solution nearly to boiling.

4. With constant stirring, slowing add 5 percent AgNO3 until no more Cl"1 is 

precipitated.

5. Digest the mixture on a steam bath until the precipitate flocculates and 

settles, leaving a clear supernatant solution.

6. Test the supernatant solution with 1 drop 5 percent AgNO3 to insure complete 

precipitation.

7. Place the mixture in the dark and allow it to stand overnight.

8. Collect the precipitate quantitatively in a tared Gooch crucible.

9. Wash the beaker and precipitate with small amounts of 0.1 percent HNO^

until the washings give no test for silver with dilute HjOl.

10. Dry the crucible and precipitate at 180°C to constant weight. Record 

the precipitate weight to the nearest 0.0001 g.
D: 10a-2 CHLORIDE 145

Calculations

ppm Cl-' = 1 .. X 1 » 000 xmg AgClX0.24737 density ml sample s 6

Report chloride concentrations >999 ppm to 3 significant figures 

only.

Preparation of reagents

Nitric acid, 0.1 percent v/v : Mix 1 ml cone HNO3 with water and dilute to approx

1,000 ml. 

Silver nitrate, 5 percent: Dissolve 5.0 g AgNO3 in water and dilute to approx

100ml.

REFERENCES

American Public Health Association and others, 1955, Standard methods for 

the examination of water, sewage, and industrial wastes: New York, Am. 

Public Health Assoc., Inc., 10th ed.

American Society for Testing Materials, 1954, Manual on industrial water: 

Spec. Tech. Pub. 148-A.

Collins, W. D., 1928, Notes on practical water analysis: U.S. Geol. Survey WaterSupply Paper 596-H.

Eaton, F. M., 1954, Formulas for estimating the drainage and gypsum requirements of irrigation waters: Texas Agr. Expt. Sta. Misc. Kept. 111.

Kolthoff, I. M., and Sandell E. B., 1952, Textbook of quantitative inorganic 

analysis : New York, Macmillan Co., 3d ed.

Taylor, E. W., 1949, The examination of waters and water supplies: Philadelphia, Pa., P. Blakiston's Son and Co.

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health 

Service Repts., v. 61, no. 11.

D:ll CHLORINE RESIDUAL

Residual chlorine includes "free available chlorine" and "combined 

chlorine." Chlorine hydrolyzes immediately in water.

'+Cr1

Through common usage the term "free available chlorine" is recognized to include HOC1 and OC1'1 and "combined available chlorine" 

to include chloramines and other chloro derivatives (APHA, 1955, 

p. 62-81) . Both free and combined chlorine may be present together. 

Results are reported in terms of chlorine (C12 ).

Chlorination is used in the treatment of many public water supplies 

and sewage effluents. Industrial processes, particularly those that use 

bleaching operations, may discharge free or combined chlorine to 

watercourses. Chlorine is unstable in solution ; strong light accelerates 

its dissipation.

It is generally agreed that the small amounts of residual chlorine 

that are present in palatable drinking water are not harmful; low 

concentrations of chlorine are reported to be toxic to fish, but the 

degree to which chlorine is harmful is dependent on pH, temperature, 

dissolved oxygen and synergism and antagonism of other pollutants 

(California State Water Pollution Control Board, 1952, 1954, p. 212).

Industrial water supplies are often chlorinated to control bacteria 

and other slime-producing organisms. Residual chlorine may cause 

taste in processed foods (Kohman, 1923). Concentrations not exceeding 2.0 ppm are recommended for the manufacture of fine paper. 

In concentrations normally found, chlorine is not reported to be injurious to land crops.

D:lla-l ORTHOTOLIDINE-ARSENITE METHOD

The orthotolidine-arsenite method is similar in substance to part 

ID, APHA (1955, p. 72-73) Standard Methods.

Principle of determination

In dilute solution, hydrolyzed chlorine oxidizes orthotolidine to 

give a yellowish-brown complex. The color developed is compared 

with that of standards or standardized colored-glass discs. Other 

free halogens react quantitatively. The reaction is practically instantaneous with free available chlorine but proceeds more slowly 

with the combined forms. This differential in reaction rates is utilized 

in differentiating between the two. Arsenite inhibits the reaction but 

does not reverse it.

The pH must be below 1.3 for proper color development, and the 

ratio of orthotolidine to chlorine should be at least 1 to 3. The reac147
148 ANALYTICAL PROCEDURES Dllla-l

tion is temperature sensitive, the precision of the test increasing with 

decreasing temperature. A reaction temperature of less than 20° C 

is recommended.

Iron, manganese, nitrite, algae, and lignocellulose interfere and 

increase the color. The effect of these interferences is evaluated by 

development of color in presence of arsenite and compensated for in 

the calculations.

Reproducibility of results is dependent on close adherence to prescribed procedure, temperature, and relative concentrations of free 

and combined available chlorine. Reproducibility of ±0.0001 mg 

chlorine may be achieved with less than 0.0005 mg and ±0.0003 mg 

in the 0.005- to 0.01-mg range.

Apparatus and reagents

Hellige Aqua Tester with standard color discs covering range of 0.00-0.10 ppm,

0.1-1.0 ppm, and 1.0-2.0 ppm. 

Sodium arsenite, 5 percent 

Orthotolidine reagent

Procedure

Because of its instability, residual chlorine should be determined 

immediately after collection of the sample. (See sec. A:4d.)

1. Pipet equal volumes of sample containing less than 0.01 mg residual chlorine 

as C12 (50.0 ml max) into 3 flasks and adjust the volumes to 50 ml.

2. To the first flask add 0.5 ml 5 percent NaAsO2 and mix. Add 5.0 ml orthotolidine, mix rapidly, and immediately compare the color in the "Aquatester." The value observed is the blank correction (Bi) for the free 

available chlorine.

3. Make a second observation on the contents of the first flask 5.0 min after 

addition of the Orthotolidine. This value is the blank correction (B2 ) for 

the total available chlorine.

4. To the second flask add 5.0 ml Orthotolidine. Mix quickly and immediately 

add 0.5 ml NaAsOs. Immediately compare the color. The observed value 

(A) is the free available chlorine plus effect of interfering substances that 

react rapidly.

5. To the third flask add 5.0 ml Orthotolidine, mix, and compare the color after 

5.0 min. The observed value (OT) is the total residual chlorine plus the 

effect of interfering substances that react in 5.0 min.

Calculations

ppm Total residual chlorine =. *. X , 5° , (OT-B2~) density ml sample

1 50 

ppm Free available chlorine =-; : : 

 X ; ^- i 

density ml sample

1 CQ

ppm Combined available chlorine =. V X r - r-[(OT-B,) (A-Bi)] density ml sample

Report chlorine concentrations of <0.10 ppm to 2 decimal places 

and of >0.10 ppm to 2 significant figures only.
D:lla-l CHLORINE RESIDUAL 149

Preparation of reagents

Sodium arsenite, 5 percent: Dissolve 5.0 g NaAsO2 in water and dilute to approx 

100 ml.

Orthotolidine reagent: Dissolve 1.35 g orthotolidine hydrochloride in approx 450 

ml water. Add this solution with constant stirring to a solution of 375 ml 

cone HC1 (sp gr 1.19) and 125 ml water. Dilute to 1,000 ml. Store in 

amber or actinic bottle and protect from light. The reagent is stable for 

6 months.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1. 

Kohman, E. F., 1923, The effect on canned foods of industrial wastes in the

water supply: Natl. Canners Assoc. Circ. 4-L. 

Technical Association of the Paper and Pulp Industry, 1948, Specifications for

chemical process water for fine paper manufacture: Tech. Assoc. Paper

and Pulp Industry Standards, B-60D, s-48, corrected.

D:12 CHROMIUM

Few if any waters contain chromium from natural sources. Hexavalent chromium salts are used in metal pickling and plating, 

anodizing aluminum, and in the manufacture of paints, dyes, explosives, ceramics, paper, and many other substances. Trivalent chromium salts are used as mordants in textile dyeing, in the ceramic 

and glass industries, and in photography. Chromium is a corrosion 

inhibitor and may be present in treated cooling waters. Waste 

products from many of these activities may contain chromium.

The U.S. Public Health Service (1946) states that the concentration of hexavalent chromium shall not exceed 0.05 ppm in drinking and culinary water on carriers subject to Federal quarantine 

regulations; no limit is given for the trivalent form. The toxicity 

of chromium salts to aquatic life differs widely with the species, 

temperature, pH, valence of chromium, and other factors. (California State Water Pollution Control Board, 1952, 1954.)

D:12a HEXAVALENT CHROMIUM

D:12a-l DIPHENYLCARBAZIDE METHOD

The diphenylcarbazide method determines only the hexavalent 

chromium.

Principle of determination

In acid solution diphenylcarbazide forms with hexavalent chromium a soluble red-violet product that absorbs light at 540 mju,. 

The formula for the colored substance is not known. For all practical purposes the reaction is specific for chromium; metallic interference almost never occurs. Iron, mercury, and molybdenum in 

concentrations as high as 100 ppm show only a small effect. Vanadium should not be present in concentrations exceeding 4 ppm. 

The effect ( f water color is small, and color up to 50 can be tolerated. 

The pH of the reaction is not critical; solutions differing in pH 

from 0.7 to 1.3 give identical colors. The color of the chromiumdiphenylcarbazide product changes slightly with time but for practical purposes it can be considered to be stable. The chromium color 

develops almost instantly and is stable, whereas vanadium color 

develops instantly and then fades rapidly. If the original vanadium 

concentration is less than 4 ppm, no color persists after 10 min.

Results are reproducible and accurate to ±0.0002 mg.

Additional information on the principle of the determination is 

given by Sandell (1950, p. 260-268).

151
152 ANALYTICAL PROCEDURES D:12a-l

Apparatus and reagents

Spectrophotometer, Beckman Model B: 

Wavelength: 540m/* 

Cells : 10-mm optical depth 

Phototube: Blue-sensitive 

Blank: Metal-free water plus reagents 

Initial sensitivity setting: 1 

Slit width : 0.1 mm (approx) 

The following absorbancies have been observed :

mg Cr Abaorbancy

0. 01 0. 7

.02 1.2

.03 1. 7

. 045 2. 5

Potassium chromate, 1.00 ml=0.10 mg Cr+6. 

Sulf uric acid, 6.5 percent v/v 

Diphenylcarbazide reagent

Procedure

Samples for the determination of chromium should be collected and 

treated as directed in sec. A :4d.

1. Pipet a volume containing less than 0.03 mg Cr+8 (10.00 ml max) into a 50-ml 

beaker and adjust the volume to 10.0 ml with metal-free water.

2. Prepare a blank of metal-free water and sufficient standards, and adjust the 

volumes to 10.0 ml with metal-free water.

3. Add 1.0 ml 6.5 percent H2 S(>4 and mix.

4. Add 0.5 ml diphenylcarbazide reagent and mix.

5. Allow to stand 10 min.

6. Determine the absorbancy of the sample and standards against the blank and 

when necessary make correction for water color as directed in sec. C: la-2.

Calculations

1. Determine the mg Cr+8 in the sample from a plot of absorbancies of standards 

containing known amounts of the constituent.

2. ppm Cr+6 = -3 X H r-Xnig Cr+6 in sample 

density ml sample e

Report hexavalent chromium concentrations of < 1.0 ppm to 2 decimal places and of >1.0 ppm to 2 significant figures only.

Preparation of reagents

Potassium chromate, 1.00 ml=0.10 mg Cr+8 : Dissolve 0.3734 g K2CrO4 dried overnight in H2 SO4 desiccator in metal-free water and dilute to 1,000 ml.

Sulfuric acid, 6.5 percent v/v: Add 6.5 ml cone H2 SO4 (sp gr 1.84) to metal-free 

water and dilute to 100 ml.

Diphenylcarbazide reagent: Dissolve 0.2 g diphenylcarbazide and 1.0 g phthalic 

anhydride in approx 200 ml ethyl alcohol. This reagent is stable for several 

weeks; slight discoloration may be noted but this does not impair the usefulness of the reagent.
D: 12b-l CHROMIUM 153

I>:121> CHROMIUM

D:12b-l PERMANGANATE-AZIDE METHOD

Principle of determination

Total chromium is determined by oxidizing trivalent chromium to 

the hexavalent state with potassium permanganate prior to diphenylcarbazide color development (see sec. D: 12a-l). The excess oxidant 

is destroyed with sodium azide.

Attention should be given to the glassware used, because scratched 

glassware may adsorb chromium. Any glassware cleaned with 

chromic acid cleaning solution should be recleansed with hydrochloric 

acid to remove the last traces of chromium.

Additional information on the principle of the determination is 

given by Saltzman (1952, p. 1016) and Lieber (1956, p. 295-299).

Apparatus and reagents

Steam bath

Spectrophotometer, Beckman Model B :

Wavelength: 540 in/*

Cells: 10-inm optical depth

Phototube: Blue-sensitive

Blank: Metal-free water plus reagents

Initial sensitivity setting: 1

Slit width : 0.1 mm (approx)

The following absorbancies have been observed : 

mg Cr Abaorbunoy

0. 01 0. 7

.02 1.2

.03 1.7

. 045 2. 5

Potassium chromate, 1.00 ml=0.10 nig Cr+s

Sulfuric acid, 6.5 percent v/v

Diphenylcarbazide reagent

Trivalent chromium, 1.00 ml=0.002 mg Cr+s

Sulfuric acid, Q.5N

Potassium permanganate, Q.1N

Sodium azide, 5 percent

Ammonium hydroxide, 50 percent v/v

Procedure

Samples for the determination of chromium should be collected and 

treated as directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.30 mg Cr (50.0 ml max) into 

a 250-ml Erlenmeyer flask and adjust the volume to approx 50 ml.

2. Prepare a blank of metal-free water and sufficient standards and adjust the 

volumes to approx 50 ml with metal-free water.

3. Add 10 ml 0.5N H2 SO4 and mix.

4. Add approx 0.5 ml 0.1N KMnO4.

5. Heat on steam bath for 20 min. If color disappears, add more KMnO« to 

maintain a slight excess.
154 ANALYTICAL PROCEDURES D:12b-l

6. Remove from steam bath, and while warm add 5 percent NaN3 dropwise until 

KMnO4 color disappears. Allow about 10 sec between each drop of NaN3 ; 

3 to 5 drops are usually sufficient.

7. Promptly cool to room temperature in a cold-water bath.

8. Add 0.75 ml 50 percent NEUOH and dilute to 100.0 ml.

9. Pipet a 10.00-ml aliquot into a 50-ml beaker and proceed as directed in sec. 

D : 12a-l. The color correction is not made.

Calculations

1. Determine the mg Cr in the aliquot from a plot of absorbancies of standards 

containing known amounts of the constituent.

0 1 1.000 ml sample ^, . .. , 2. ppm Cr=-3 TT X j-^ i~X ; ..- y . Xmg Cr in aliquot density ml sample ml aliquot

Keport chromium concentrations of <1.0 ppm to 2 decimal places 

and of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Potassium chromate, 1.00 ml=0.10 mg Cr+8 : Dissolve 0.3734 g KsCrO* dried overnight in HjSO* desiccator in metal-free water and dilute to 1,000 ml.

Sulfuric acid, 6.5 percent v/v: Add 6.5 ml cone H2SC>4 (sp gr 1.84) to metal-free 

water and dilute to 100 ml.

Diphenylcarbazide reagent: Dissolve 0.2 g diphenylcarbazide and 1.0 g phthalic 

anhydride in approx 200 ml ethyl alcohol. This reagent is stable for several 

weeks; slight discoloration may be noted but this does not impair the 

usefulness of the reagent.

Trivalent chromium, 1.00 ml=0.002 mg Cr+3 : Dissolve 0.2263 g K2Cr2O7 in metalfree water and dilute to 1,000 ml. Pipet 5.00 ml into an Erlenmeyer flask. 

Add approx 15 mg Na2 SO3 and 0.5 ml cone HNO3 (sp gr 1.42). Evaporate 

to dryness gently; strong heating reoxidizes the Cr. Add 0.5 ml cone HNOs 

and again evaporate to dryness to destroy any excess sulfite. Take up in 

1 ml cone HNO3 with warming and dilute to 200.0 ml with metal-free water.

Sulfuric acid, 0.5^: Mix 13.9 ml cone H2SO4 (sp gr 1.84) with metal-free water 

and dilute to approx 1 liter.

Potassium permanganate, O.IN: Dissolve 0.32 g KMnO4 in approx 100 ml metalfree water. Allow to stand several days and decant or filter through 

medium-porosity fritted glass.

Sodium azide, 5 percent: Dissolve 5 g NaN3 in approx 100 ml metal-free water.

Ammonium hydroxide, 50 percent v/v: Mix 50 ml cone NH4OH (sp gr 0.900) 

with water and dilute to 100 ml.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1. 

Lieber, Maxim, 1956, Permanganate-azide test for total chromium in water:

Am. Water Works Assoc. Jour., v. 48. 

Saltzman, B. E., 1952, Microdetermination of chromium with diphenyl-carbazide

by permanganate oxidation: Anal. Chemistry, v. 24. 

Sandell, E. B., 1950, Colorimetric determination of traces of metals: New

York, Interscience Publishers Inc., 2d ed. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public

Health Service Repts., v. 61, no. 11.
D:13 COLOR

The color of water as considered herein is that due only to substances in solution. Color in water may be of natural mineral, animal, or vegetable origin. It may be caused by metallic substances, 

humus material, peat, algae, weeds, or protozoa. Industrial wastes 

may also color water. Color may range from zero to several hundred units.

In domestic water, color is undesirable from aesthetic considerations and because it may dull clothes or stain food and fixtures. 

The U.S. Public Health Service (1946) states that the color shall not 

exceed 20 units in drinking and culinary water on carriers subject 

to Federal quarantine regulations. Color is undesirable in water 

for many industries, particularly food processing, laundering, ice 

manufacturing, bottled beverage, photographic and textile (California State Water Pollution Control Board, 1952, 1954, p. 226-228).

D:13a-l COMPARISON METHOD

Principle of determination

The color of the water is compared to that of glass colored discs 

which have been calibrated to correspond to the platinum-cobalt 

scale of Hazen (1892, p. 427-428). The unit of color is that produced by 1 mg of platinum per liter. The Hazen scale is usually 

satisfactory for most waters, but the hues and shades of some waters 

may not easily be compared with standards. If the hue of the 

water does not compare with that of the standard there is very 

little that can be done about it, other than to visually compare the 

optical densities of the sample and standard. A minimum of dilution of highly colored water is recommended because the color of 

the diluted sample often is not proportional to the dilution.

Turbidity causes the observed color to be higher than the true 

color, but there apparently is some disagreement as to the magnitude of the effect of turbidity. One authoritative source states 

that "Even a slight turbidity causes the apparent color to be noticeably higher than the true color" (APHA, 1955, p. 87), whereas 

another states that "The color of water with low turbidity is substantially the same as that of clear water" (California State Water 

Pollution Control Board, 1952, 1954, p. 226-228). The removal 

of turbidity is the most recurrent problem in the determination of 

color. Color is removed by adsorption on suspended material, and 

filtration of the sample through paper and other media which pass 

the water through a membrane of concentrated sedimentary material tends to decolorize the water. Centrifiiging is preferable to 

filtration through any media for turbidity removal, but centrifuging

155
156 ANALYTICAL PROCEDURES D:13ar-l

is not completely effective if the water contains very finely divided 

particles. Flocculation of the dispersed particles with a strong 

electrolyte has been proposed (Lamar, 1949, p. 726). The electrolytic method is effective for removing turbidity, but the process 

of flocculation is also often used to decolorize some waters. The 

flocculation method may have some application but is probably not 

suitable for every type of color in water.

The color in some highly colored waters may have a tendency to 

fade and (or) precipitate on standing. Biological changes during 

storage may also affect the color. Consequently the color determination should be made within a reasonable period of time after the 

sample is collected.

Because of the many complicating factors involved, the determination of color is one of the least precise in water analysis. No statements on the accuracy and reproducibility of the tests can be made.

Apparatus

Hellige Aqua Tester (see sec. C : 2f).

Procedure

1. Fill one instrument tube with the sample of water, level, insert the glass 

plug making sure that no air bubbles are trapped, and insert the tube into 

the Aqua Tester.

2. Use distilled water in the second tube as a blank.

3. The color comparison is made by revolving the disc until the color of the 

two tubes matches. Waters that have colors greater than 70 should be 

diluted.

Calculations

The color is read directly from the matching color standard, and 

the proper dilution factor is applied. Report color as follows:

Record units 

Color unit to nearest 

1-50 1

51-100 5

101-250 10

250-500 20

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1. 

Hazen, Alien, 1892, A new folor standard for natural waters: Am. Chem. Soc.

Jour., v. 12. 

Lamar, W. L., 1949, Determination of color of turbid waters: Anal. Chemistry,

v. 21. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public

Health Service Repts., v. 61, no. 11.
D:14 COPPER

Most copper minerals are relatively insoluble, and little copper 

found in water is of natural origin. The presence of copper in more 

than trace amounts can usually be attributed to corrosive action of 

water on copper pipes, to industrial wastes, or to the use of copper 

salts for the control of algae and other aquatic growths.

Copper imparts a disagreeable metallic taste to water. As little 

as 1.5 ppm can usually be detected, and 5 ppm can render the 

water unpalatable. Copper is not considered to be a cumulative 

systemic poison like lead and mercury; most copper ingested is 

excreted by the body and very little is retained. The pathological 

effects of copper are controversial, but it is generally believed very 

unlikely that humans could unknowingly ingest toxic quantities 

from palatable drinking water. The U. S. Public Health Service 

(1946) recommends that copper should not exceed 3.0 ppm in drinking and culinary water on carriers subject to Federal quarantine 

regulations. The toxicity to aquatic organisms differs significantly 

not only with the species but also with the chemical and physical 

characteristics of the water, such as temperature, hardness, turbidity, and carbon dioxide content (California State Water Pollution Control Board, 1952, p. 230).

Copper is undesirable in water used for canning foods (Weckel, 

1942) and for metal-plating baths (Irenas, 1946). Although copper 

is an essential or beneficial plant nutrient, crops differ widely in 

their tolerances. Low concentrations have been found to be injurious to orange seedlings, flax, sugar beets, tomatoes, and barley, but 

high concentrations had no effect on oats or kale (California State 

Water Pollution Control Board, 1952, p. 230).

D:14a-l DIETHYLDITHIOCARBAMATE SPECTROPHOTOMETRIC METHOD

Principle of determination

Copper reacts with the substituted carbamate, dihydroxyethyldithiocarbamic acid, to give a stable colored complex, which is a 

true solution. With freshly prepared reagents, the color development is instantaneous. If copper is determined infrequently, preparation of fresh carbon disulfide and diethanol-amine solutions is 

recommended for each set of samples. Refrigerated reagents are 

sufficiently stable for about 1 month.

Nickel, cobalt, and bismuth also give colored solutions with the 

reagents, but these metals must be present in concentrations 20 

times that of copper to equal the copper color. At the wavelength 

used, the absorbancy of natural color in the water is appreciable, 

and a color correction is often necessary. Iron reacts with the 

reagent to give a color of less intensity than copper. This inter157
158 ANALYTICAL PROCEDURES D : 14a^l

ference is eliminated by complexing iron with citrate and compensating for the slight yellow color of the iron-citrate complex.

With the listed apparatus, results are usually accurate and reproducible to ±0.002 mg.

Additional information on the principle of the determination is 

given by Sandell (1950, p. 309).

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength: 425 m/j 

Cells : 40-mm optical depth 

Phototube: Blue-sensitive 

Blank : Metal-free water plus reagents 

Initial sensitivity setting: 2 

Slit width : 0.2 mm (approx) 

The following absorbancies have been observed :

mg Cu Absorbancy

0. 0125 0. 33

.0250 . 66

. 0375 . 98

. 0500 1. 31

Copper solution, 1.00 ml=0.005 mg Cu

Nitric acid, 5 percent v/v

Sodium citrate, 10 percent

Ammonium hydroxide, cone (sp gr 0.900)

Dihydroxyethyldithiocarbamicacid

Procedure

Samples for the determination of copper should be treated as directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.05 mg Cu (25.00 ml max) 

into a 50-ml beaker and adjust the volume to 25.0 ml with metal-free water.

2. Prepare a blank of 25.0 ml metal-free water and sufficient standards and 

adjust volumes to 25.0 ml with metal-free water.

3. Add 1.0 ml 5 percent HNO3.

4. Add 1.0 ml 10 percent Na3C6H5Or.

5. Add 1.0 ml cone NH4OH. Stir the solutions vigorously.

6. Add 1.0 ml dihydroxyethyldithiocarbamic acid and mix well.

7. When necessary, make corrections for water color and iron as directed in 

sec. C : la-2, method 2, to obtain true sample absorbancy.

Calculations

1. Determine the mg Cu in sample from a plot of absorbancies of standards 

containing known amounts of constituent.

2. ppm Cu=-; rr~X H 5~Xnig Cu in sample *^ density ml sample

Report copper concentrations of <1.0 ppm to 2 decimal places and 

of >1.0 ppm to 2 significant figures only.
D:14a^l COPPER 159

Preparation of reagents

Copper solution, 1.00 ml=0.005 mg Cu: Dilute 5.00 ml Cu (1.00 ml=1.00 mg Cu)

to 1,000 ml with metal-free water containing 1 or 2 drops cone HNO» (sp

grl.42). 

Copper solution, 1.00 ml=1.00 mg Cu: Dissolve 1.00 g Cu in a minimum of

dilute HNO3 and dilute to 1,000 ml with metal-free water. 

Nitric acid, 5 percent v/v: Mix 5 ml cone HNO3 (sp gr 1.42) with metal-free

water and dilute to approx 100 ml. 

Sodium citrate, 10 percent: Dissolve 10 g NaaCeHsOi in metal-free water and dilute

to approx 100 ml. 

Dihydroxyethyldithiocarbamic acid: Just before use, mix equal volumes of 0.32

percent CSj and 3.6 percent diethanol-amine. 

Carbon disulfide, 0.32 percent v/v: Dissolve 0.32 ml C& in ethyl alcohol and

dilute to 100 ml. Store in refrigerator. 

Diethanol-amine, 3.6 percent v/v: Dissolve 3.6 ml diethanol-amine in ethyl alcohol

and dilute to 100 ml. Store in refrigerafr»r.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Irenas, Z., 1946, The significance of water in the plating room: Water Pollution

Abs. 19 [August]. 

Sandell, E. B., 1950, Colorimetric determination of trace metals: New York,

Interscience Publishers. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public

Health Service Repts., v. 61, no. 11. 

Weckel, K. G., 1942, Watch your water supply it affects food quality: Water

Pollution Abs. 15 [March].

D:15 DENSITY

Density is the mass of any substance per unit volume at a designated standard temperature. Density should not be confused with 

specific gravity, which is a mass-to-mass relation.

The density value has some use in industries that utilize brines and 

whose basic unit of concentration of dissolved material is density. 

Density is used primarily by the chemist in the computation of 

parts per million for highly mineralized waters. In the analysis, 

weight per sample volume is determined, and these values must be 

converted to parts per million. Parts per million by definition is 

the weight of dissolved material per 1 million equal weights of 

solution (milligrams per kilogram). Therefore the determined 

concentration of each constituent is divided by the density to give 

the correct parts per million. This computation is made before 

rounding the results.

A dissolved-materials content of 1,000 ppm increases the density 

of the solution by slightly less than 1 g per kg, or less than 0.1 

percent; the exact amount is dependent on the nature of the solid 

material. Consequently differences in density are not significant 

to the chemical analysis unless the total concentration of dissolved 

solids exceeds about 5,000-10,000 ppm. The Geological Survey has 

arbitrarily selected 7,000 ppm as the dividing line. If the dissolvedsolids content is less than 7,000 ppm it is assumed that exactly 1 ml 

of solution weighs exactly 1 g; if it exceeds 7,000 ppm the density 

is determined and used in the computation of parts per million for 

all constituents. If determination of dissolved solids is not included 

in the scope of the analysis, a specific conductance of 10,000 is then 

taken as the differentiating measurement.

D:15a-l GRAVIMETRIC METHOD

Principle of determination

The density determination is based on the weight of a carefully 

measured volume of solution at a given temperature. Densities are 

determined at 20° C, the same temperature at which volumetric 

glassware is calibrated to deliver a given weight of pure water.

Results are accurate and reproducible to ±0.0005 g.

Apparatus

Weighing bottle: 50-| ml capacity

Volumetric pipet, 50-ml, calibrated. The actual volume delivery of the pipet 

is determined by weighing a delivered volume of dilution water at 20°C. 

The volume is obtained from relative-density tables in handbooks. Alternatively, a 50-ml pycnometer can be used; it must be calibrated also.

161
162 DENSITY

Procedure

1. Adjust the temperature of the sample to 20.0°C.

2. Rapidly withdraw a sample and transfer it to a tared weighing bottle.

3. Stopper the bottle immediately to prevent water loss by evaporation.

4. Weigh the solution to the nearest 0.1 mg.

Calculations

Density=- . g sample 5Li ^-rml sample

If a constant-temperature bath is not available, the determination 

can be made at the sample temperature and a correction applied for 

the departure from 20°C. The temperature is recorded with an 

accurate thermometer, and the relative density for that temperature 

obtained from a table. The density result is then corrected by the 

factor:

Relative density (20° C) 

Relative density (test temp in °C)

Report density to 3 decimal places in terms of grams per milliliter 

at 20°C.
D:16 FLUORIDE

Unlike chlorides, fluorides are only sparingly soluble and are 

present in most natural waters in only small amounts. Calcium 

fluoride (fluorite) is the principal source of fluoride, but there are 

some other complex fluoride-bearing minerals. The element is often 

characteristic of waters from deep strata and is frequently found in 

salt water from oil wells and in water from areas that have been 

subjected to recent vulcanism. Fluorides are used as insecticides, 

disinfectants, preservatives, and for a few other purposes in industry. They are seldom found in large quantities in industrial wastes, 

except as the result of spillage.

Pathological changes in man attributable to fluoride absorption 

are in the nature of ostosis. Large quantities of fluoride are toxic, 

but it has been reported that adults may safely drink 2 gallons 

per day of water containing 10 ppm of fluoride (Smith, 1944, p. 

1293). Daily intakes of about 15-20 mg of fluoride over a period 

of several years are required to induce chronic fluorosis in an adult 

man (California State Water Pollution Control Board, 1952, p. 256). 

Excessive quantities in drinking water during calcification may cause 

discoloration in the teeth of children; adults are not affected. Available evidence indicates that water containing less than 1.0-0.9 ppm 

of fluoride seldom causes mottling of children's teeth, and the literature describing the beneficial effect of 0.88-1.5 ppm in drinking 

water as an aid in the reduction of tooth decay in children is 

abundant. The U. S. Public Health Service (1946) states that the 

concentration of fluoride shall not exceed 1.5 ppm in drinking and 

culinary water on carriers subject to Federal quarantine regulations.

Manufacturers of products for internal consumption generally 

limit the fluoride content of water to about 1.0 ppm. In normal 

concentrations it is not significant in irrigation; it is generally 

concluded that the fluoride content of irrigation water has no consistent influence on the fluoride content of plants. The effect of 

fluoride on livestock is similar to that on humans (California State 

Water Pollution Control Board, 1952, p. 256).

D:16a-l ZIRCONIUM-ERIOCHROME CYANINE R METHOD

The Zirconium Eriochrome Cyanine R method is a much modified version of the procedure of Megregian (1954, p. 1161).

Principle of determination

In acid solution, zirconium reacts with Erichrome Cyanine R to 

form a red complex ion. Fluoride forms a more stable complex 

with zirconium (ZrFG -2 ) and withdraws zirconium from the organic 

complex to produce a bleaching effect. Eriochrome Cyanine R

163
164 ANALYTICAL PROCEDURES

shows a decided specificity to zirconium. Under the experimental 

conditions the dye does not give a color with titanium or beryllium, 

two metals which react with many other zirconium agents. Aluminum reacts to give a positive interference that is easily eliminated. 

This is accomplished by allowing the solution to stand for at least 

2 hr before making color comparison. Up to 10 ppm can be tolerated.

Analytical conditions are not overly critical. The pH is controlled at a highly acid level by the addition of 1.7 ml of concentrated hydrochloric acid to each sample. This assures that high 

concentrations of bicarbonate or other alkaline ions will not affect 

the pH significantly. Sulfate interferes but is removed in the procedure by precipitation as barium sulfate. Overnight standing is 

usually required to assure complete settling of barium sulfate before 

making color comparison. The clarification of the sample can be 

accelerated by centrifuging if the fluoride result is desired immediately. Filtration should not be used.

Eesidual chlorine, chromate, and probably other strong oxidants 

attack the indicator. The susceptibility to attack varies with 

batches of indicator. Stannous chloride is used to eliminate chromate and chlorine interference. Chromium, cadmium, and nickel, 

in concentrations of less than 5 ppm, do not interfere in the lower 

fluoride range. When the fluoride concentration exceeds 1.0 ppm, 

larger quantities of these metals can be tolerated. Ten ppm iron, 

zinc, lead, cyanide and phosphate, cause no appreciable interference if the sample is allowed to stand overnight.

The determination shows "salt effect"; the sensitivity is depressed 

by 5 to 10 percent at a dissolved-solids concentration of 10,000 ppm. 

The effect of the usual type of color is not serious. A color of 70 

on the Hellige scale is equivalent to an absorbancy error of only 

0.005 in the spectrophotometric measurement. Thus, it appears that 

color correction will not often be necessary.

The quality of batches of Eriochrome Cyanine E from different 

sources differs very significantly, and it is necessary to test the 

reagent each time it is prepared. The individual absorbancy curves 

show corresponding differences, and the sensitivity of fluoride 

between reagents may differ by 20 percent.

The method has rather good tolerance for temperature differences. For most purposes, operating at room temperature without 

other precautions is satisfactory.

With listed apparatus, results are accurate and reproducible to 

±0.0005 mg in the lower ranges and approx 0.001 mg in the higher 

fluoride range.
D: 16a-l FLUORIDE 165

Apparatus and reagents 

Spectrophotometer, Beckman Model B:

Wavelength: 540 m/j

Cells: 40-mm optical depth

Phototube: Blue-sensitive

Blank : Dilution water plus reagents

Initial sensitivity setting: 3

Slit width : 0.3 mm (approx)

The following absorbancies have been observed :

ing F-1 Absorbancy

0. 00 1. 50

. 01 1. 15

.02 .82

.03 .56

Sodium fluoride, 1.00 ml=0.010 mg F"1 

Stannous chloride reagent, 2 percent 

Indicator solution

Procedure

If the sample contains an excessive amount of interfering materials, 

the fluoride should be isolated by distillation (see sec. D: 16a-3).

1. Pipet a volume of sample containing less than 0.03 mg F'1 (10.00 ml max) 

into a 50-ml centrifuge tube or test tube.

2. Prepare a blank and sufficient standards, and adjust the volumes to 10.0 ml.

3. If chromate, residual chlorine, or other strong oxidizing agents are present 

in the sample, add 0.1 ml 2 percent SnCl2 and let the solution stand for 

10 min.

4. Add 25.0 ml indicator.

5. Allow the solution to stand overnight for barium sulfate to settle.

6. Decant approx 25 ml pure supernatant solution, taking care not to disturb 

the precipitate.

7. Determine the absorbancy of the test sample and standards against the blank, 

and when necessary make correction for water color as directed in sec. 

C : la-2, method 1.

Calculations

1. Determine mg F"1 in test sample from a plot of absorbancies of standards 

containing known amounts of constituent.

2. ppm F-*= . * X . 1|000 , Xmg F-» in sample *^ density ml sample 6

Keport fluoride concentrations of <1 ppm to 0.1 ppm and of >1 

ppm to 2 significant figures only.

Preparation of reagents

Sodium fluoride, 1.00 ml=:0.010 mg F"1 : Dissolve 0.2210 g NaF in water and 

dilute to 1,000 ml in metal-free water. Dilute 100.0 ml of this stock solution to 1,000ml.

Stannous chloride reagent, 2 percent: Dissolve 1 g SnCl2 -2H2O in 10 ml cone 

HC1 (sp gr 1.19) and dilute to approx 50 ml.
166 ANALYTICAL PROCEDURES D:16a-2

Indicator solution: To about 300 ml metal-free water add 20.0 ml 0.9 percent 

Eriochrome Cyanine B and 10.0 ml 0.2 percent ZrO(NO3 ) 2 -2H2O in acid. 

Add 70 ml cone HC1 (sp gr 1.19) and 4 g BaCk Dissolve and dilute to 1,000 

ml. The solution is consumed rapidly and a larger volume is normally 

prepared.

Eriochrome Cyanine R, 0.9 percent: Dissolve 1.80 g tested Eriochrome Cyanine 

R in water and dilute to 200 ml. The National Aniline product labeled 

"Alizarol Cyanone RC" has been used successfully. With other products, a 

precipitate sometimes forms when the indicator solution is prepared.

Zirconyl nitrate, 0.2 percent in acid: Dissolve 0.40 g ZrO(NOs ) 2 -2H2O in 200 ml 

50 percent HC1 v/v.

D:16a-2 ZIRCONIUM-ALIZARIN METHOD

The zirconium-alizarin method is similar in substance to part ID, 

APHA (1955, p. 105-107) Standard Methods, and method 85, U.S. 

Salinity Laboratory Staff (1954, p. 147) Handbook 60.

Principle of determination

In acid solution, zirconium reacts with alizarin red S to form a 

reddish-violet lake. Upon addition of fluoride ion to the zirconium 

red lake, the more stable fluorozirconate complex ion, ZrF6~2, is formed. 

If the amount of zirconium and alizarin is carefully controlled, the 

amount of zirconium removed by fluoride subtracts from the reddishcomplex color. The reaction simultaneously liberates free alizarin 

sulfonic acid which is yellow in acid solution. With increasing fluoride, the color change is from reddish violet to yellow green. The 

bleaching effect of the fluoride in the sample is compared visually 

with that in standard solutions. As the color approaches yellow green, 

the test is less sensitive; the method is most useful when the sample 

volume contains less than 0.10 nig of fluoride, but the range can be 

extended a little higher. The reaction requires several hours to reach 

equilibrium, and the rate is somewhat dependent on temperature. 

Color comparison can be made at the end of 1 hr if the reagent was 

added to the sample and standards within 2 min and uniform temperature is maintained.

The interference of 20 mg of calcium, 20 mg of magnesium, 100 

mg of sodium, 20 mg of potassium, and 30 mg of nitrate is negligible; even higher concentrations would cause very little trouble. 

Up to 1.5 mg of iron does not interfere seriously, but larger amounts 

cannot be tolerated; 2.5 mg of manganese does not interfere. Free 

residual chlorine interferes but can be removed by adding sodium 

arsenite. Several other ions common to water interfere; some give 

high results and others low. Aluminum gives a negative error because of the formation of aluminum-fluoride complex, which withdraws fluoride from the reaction of zirconium. The magnitude of 

the interference is dependent on the quantity of both alluminum and
D:16a-2 FLUORIDE 167

fluoride, but the percentage of fluoride complexed is approximately 

constant for a given concentration of aluminum. The ratio taken 

from the graph shown below (fig. 16) and multiplied by the apparent fluoride concentration will give the actual fluoride content. For 

fluoride concentrations not exceeding 0.10 mg and aluminum concentrations not exceeding 0.05 mg, the maximum variable errors will 

not exceed 0.01 mg of fluoride.

In the presence of 1.0 mg of aluminum, the ratio of actual to apparent fluoride concentration is 2.8. The interference of 0.05 mg 

of aluminum does not exceed the equivalent of 0.01 mg of fluoride. 

Hexavalent chromium causes a positive error, but 0.05 mg can be 

tolerated if the color comparison is made 1 hr after addition of 

the reagent, but only 0.01 mg of chromium can be tolerated when 

readings are taken between 7 and 18 hr. The interference of orthophosphate and hexametaphosphate is lessened if the reaction is allowed to go 7-18 hr; 0.20 mg can be tolerated. Below pH 4.5, 

hydrogen ion gives positive errors but the effect can be eliminated 

by raising the pH of the sample to 4.5-6.0 with sodium hydroxide.

co

1.0 

0.9 

0.8

0.7 

oc.

3 0.6

1

2 0.5

s

2 0.4

±)

^ 0.3

0.2 

0.1

1.0 1.2 1.4 1.6 1.8 2.0 2.2 2.4 2.6 2.8

RATIO

FIGURE 16. Variation of the ratio of actual to apparent fluoride content with

aluminum content.
168 ANALYTICAL PROCEDURES

Sulfate increases the fluoride reading, while chloride and alkalinity decrease it. These effects are substantially additive, and 

corrections taken from the following table can be applied to the 

apparent fluoride concentration. Alkalinity in excess of 10 mg of 

calcium carbonate can be neutralized with an equivalent quantity of 

nitric acid.

Corrections for fluoride determination 

[mg Actual F=mg apparent F+value indicated]

Concentration of foreign ion

10 _ _ _

20 _________ ___ __

30___ _____________________________

40

50 __________

60

100 _

200

Sulfate

-0. 002

-. 002

-. 003

-. 006

-. 009

-. Oil

-. 018

mg per 100 ml

Chloride

0. 002

. 003

. 004

. 006

.007

. 012

. 018

Alkalinity 

as CaCO3

0. 005

. Oil

. 017

. 023

.028.

Turbidity interferes with color precipitation, as does the natural 

color of the water to a lesser extent. Turbid water should be filtered through the Millipore filter membrane. Most quantitative 

filter papers are washed with hydrofluoric acid and can add fluoride 

to the sample. The effect of color can be partly compensated for 

by placing a tube of distilled water below the sample and an equal 

volume of water treated with acid below the standard and by 

making a color comparison through both tubes.

When color and turbidity do not interfere with color comparison, 

results are generally accurate and reproducible to ±0.01 mg.

When interfering substances are present in significant amounts, 

distillation of the fluoride from the sample as hydrofluosilicic acid 

(see sec. D:16a-3) followed by reaction with zirconium-alizarin 

reagent is preferable to corrections of apparent fluoride concentration.

Additional information on the principle of the determination is 

given by Lamar and Drake (1955, p. 563-572).

Apparatus and reagents

Color comparator, 3-hole, that permits longitudinal viewing of the contents of

the tube.

Nessler tubes, low form, matched 

Sodium an enite, 1.0 mlol.O mg CU 

Acid indicator reagent 

Sodium fluoride, 1.001111=0.010 mg F"1
FLUOEIDE 169

Procedure

If the sample contains an excessive quantity of interfering materials, 

the fluoride should be isolated by the distillation procedure (see sec. 

D:16a-3).

1. Measure a volume of sample containing less than 0.12 mg F"1 (100.0 ml max) 

into Nessler tubes and adjust the volume to 100 ml.

2. Neutralize residual chlorine with NaAsO2 (1.0 mlol.O mg Clz ), adding 0.1 

ml in excess.

3. Prepare a blank and sufficient standards, and adjust the volumes to 100 ml.

4. Add 10.0 ml acid indicator reagent and mix.

5. Allow the solutions to stand 7-18 hr and visually compare the color of the 

sample with that of standards in the comparator.

Calculations

1. Correct the apparent F"1 concentration for sulfate, chloride, and alkalinity 

as required from figure 16 and the table shown under "Principle of determination."

2 ppm F-i=-^ X MPO_ x /corrected mg F~'V aluminum \ 

density ml sample \ in sample /\correction ratio/

Report fluoride concentrations of <1 ppm to 0.1 ppm and of >1 

ppm to 2 significant figures only.

Preparation of reagents

Acid indicator reagent: Add 25.0 ml ZrCKNOsh solution to approx 100 ml water.

With constant stirring, slowly add 25.0 ml alizarin solution and dilute to

approx 450 ml with water and mix well. Add 500 ml 2.1N H2SO4 and dilute

to 1,000 ml with water. Allow the indicator to stand approx 1 hr before

use. Storage in a lightproof bottle increases the stability of the indicator.

If the indicator precipitates on standing, shake it thoroughly before use. 

Zirconyl nitrate (or zirconyl chloride) : Dissolve 1.84 g ZrO(NO3 ) z -2H2O or 2.22

g ZrOCl2 -8H2O in water and dilute to 250 ml. Turbid solutions should not

be filtered. 

Alizarin red S : Dissolve 0.37 g alizarin monosodium sulfonate in water and dilute

to 250 ml. 

Sodium arsenite, 1.0 mlol.O mg C12 : Dissolve 0.916 g NaAsO2 in water and dilute

to 500.0 ml. 

Sodium fluoride, 1.00 ml=0.010 mg F-1 : Dissolve 0.2210 g NaF in water and

dilute to 1,000 ml. Dilute 100.0 ml of this stock solution to 1,000 ml.

D:16a-3 ISOLATION OF FLUORIDE BY DISTILLATION

The isolation of fluoride by distillation is similar in substance to 

part IA, APHA (1955, p. 99-101) Standard Methods. 

Principle of determination

The alkalized sample is concentrated to a small volume to facilitate 

quantitative steam distillation. After acidification, silver sulfate is 

added to prevent distillation of chloride as hydrochloric acid and the 

fluoride is microdistilled as hydrofluosilicic acid.
170 ANALYTICAL PROCEDURES D:16ar-3

Apparatus and reagents

Distillation assembly consisting of steam-generating assembly, fluoride-distillation 

section, and a distillate-receiving flask. (See flg. 17.)

FIGURE 17. Distillation apparatus for the Isolation of fluorlde.

K. Steam-inlet tube, with Hoffman

clamp.

L. Fluoride-distillation flask adapter. 

M. Fluoride-distillation flask, 300 ml. 

N. Heating mantle with asbestos shield. 

0. Thermometer, 250 e C. 

P. Trap.

Q. Condenser adapter. 

R. Condenser. 

8. Distillate-receiving flask, 200 ml.

A. Pressure-control tube.

B. Hot-water reservoir, 1 liter.

C. Hotplate.

D. Hot-water-supply tube clamp.

E. Steam generator flask, 1 liter.

F. Heating mantle.

<?. Pressure-relief tube.

H, Steam-regulation tube with Hoffman

clamp.

/. Glass steam line (asbestos covered). 

J. Steam trap (asbestos covered), with

clamp on relief-water drain.

Phenolphthalein indicator solution 

Sodium hydroxide, pellets 

Sulfuric acid, cone (sp gr 1.84) 

Silver sulfate, powder

Procedure

1. Measure 200.0 ml of sample into the fluoride distillation flask (M).

2. Add a few drops of phenolphthalein indicator solution and make alkaline 

with NaOH. Insert some glass beads to prevent bumping.

3. Concentrate sample by evaporation to a volume of 15-20 nil. If solution tends 

to buinp at this point, add more glass beads.

4. Cool the concentrated sample, and acidify with a few drops of cone Hi-SO-i.

5. Add a slight excess of Ag2 SOi powder.

6. Disconnect steam-inlet tube (K), and in its place insert a long-stemmed 

separatory funnel containing 20 nil cone H2SO4 .
D: 16a-3 FLUORIDE 171

7. Turn condenser on.

8. Slowly introduce 20 ml cone H2SO4 into distillation flask (M). Before removal of funnel, wash tip of funnel in flask.

9. Reconnect steam-inlet tube (K) to flask.

10. Pass steam through assembly by opening valve on steam trap (J) and closing 

valve on steam-regulating tube (H). Adjust clamp so that a pressure of 

20-40 cm is maintained in the steam-generator pressure-relief tube (<?) 

The distilled water in the steam generator should be made slightly alkaline 

with NaOH.

11. Apply even heat to distillation flask (M). Use an asbestos pad, and do 

not allow heat to be applied around sides of flask above solution.

12. Maintain a temperature of 135°-140°C in sample by means of the steam 

flow and heat applied to the flask.

13. Collect a sufficient volume of distillate to fill a 200-ml volumetric flask. The 

distillate should be free of interferences.

14. Continue as directed in sec. D: 16a-l or 2.

Preparation of reagents

Phenolphthalein indicator solution: Dissolve 2.5 g phenolphthalein in approx 

500 ml 50 percent ethyl alcohol. Neutralize with Q.02N NaOH.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and Industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Lamar, W. L-, and Drake, P. G., 1955, Factors affecting the determination of

fluoride in water with zirconium-alizarin: Am. Water Works Assoc. Jour.,

v. 47. 

Megregian, Stephen, 1954, Rapid spectrophotometric determination of fluoride

with Zirconium-Eriochrome Cyanine R Lake: Anal. Chemistry, v. 26. 

Smith, O. M., 1944, The detection of poisons in public water supplies: Water

Works Eng., v. 97. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts., v. 61, no. 11. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline

and alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60.

D:17 HARDNESS

Hardness of water is the property attributable to the presence of 

alkaline earths. Calcium and magnesium are the principal alkaline 

earths in natural waters, whereas strontium and barium are usually 

present only in small quantities.

Hardness of water results from the solution of alkaline-earth 

minerals from the soil and rocks, or the minerals may enter from 

direct pollution by wastes. Calcium and magnesium carbonates 

(limestone and dolomite) are prevalent in the earth's crust but are 

only sparingly soluble in pure water. Water that contains carbon 

dioxide or other acidic constituents readily dissolves carbonate 

minerals; in the presence of carbon dioxide the carbonates are converted to the more soluble bicarbonates.

GaCO3 + CO2 + H2O -» Ca (HGO3 ) 2

Many waters with a hardness of less than 200-300 ppm may derive 

practically all of their alkaline earths from carbonate rocks. Gypsiferous shale and evaporites often contain large quantities of more 

soluble sulf ates and chlorides of calcium and magnesium, and waters 

that traverse these deposits may have a hardness of several hundred 

parts per million or more.

Hard water is not generally believed to have harmful effects on 

man, although the relation to urinary concretions is controversial. 

Hard water decreases the sensitivity of fish to toxic metals (California State Water Pollution Control Board, 1952, 1954, p. 265- 

26Y), but experiments with calves and chicks have indicated that 

those supplied with hard water develop somewhat better than those 

supplied with distilled water.

Hardness in conjunction with other chemical properties, such as 

acidity and other polyvalent cations, is an indication of the soapconsuming power of the water. Soap will not cleanse or lather 

until these constituents have either been neutralized or precipitated 

as insoluble salts of the fatty acids. Hard water is recognized by 

the curd formed with soap. Carbonates and some sulfates of the 

alkaline earths are sparingly soluble and tend to precipitate on 

evaporation; heating converts bicarbonate to carbonate, which precipitates calcium and magnesium carbonate in boilers, pipes, cooking vessels, etc. Calcium sulfate is also relatively insoluble and 

in addition to silica may comprise a portion of many incrustations 

found in the home or in industry.

Hardness limitations for water used in steam generation are 

very exacting and only 2 ppm or less is generally considered per173
174 ANALYTICAL PROCEDURES D:17a

missible in feed water for boilers operating at 400 psi or more. The 

tolerances of process water differ from one industiy to another and 

range from less than 10 ppm to several hundred (California State 

Water Pollution Control Board, 1952, 1954, p. 265-267). Hard 

water is usually superior to soft water for irrigation, and it is often 

stated that hard water makes soft soil and soft water makes hard 

soil. (See significance of calcium and sodium in sees. D:8 and 

D:35, respectively.)

D:17a HARDNESS

D:17a-l COMPLEXOMETEIC METHOD

The complexometric method is similar in substance to part 1C, 

APHA (1955, p. 112-117) Standard Methods; D 1126-53 T, ASTM 

(1954, p. 253-254) Manual on Industrial Water; and method 7, 

U.S. Salinity Laboratory Staff (1954, p. 94-95) Handbook 60.

The procedure is applicable to most natural and treated waters, 

but the method fails conspicuously at times with acid or polluted 

waters that contain excessive amounts of heavy metals.

Principle of determination

Disodium dihydrogen ethylenediamine tetraacetate (Na2EDTA) 

forms a slightly ionized colorless stable complex with alkaline earth 

ions. The indicator Eriochrome Black T is bright blue in the 

absence of alkaline earths, but with them forms a deep red complex 

which has a higher ionization constant than the Na2EDTA complex. Hence, by using Eriochrome Black T as an indicator, the 

alkaline earths can be titrated with Na2EDTA. For example, with 

calcium, the reaction is :

o o

NaO C CH» CHj C ONa

N CHj CHj N +Ca«

HO C CHj CHj C OH

I A

NaO C CHj CHj C ONa

O N CHj CHj N O +2H+*

CH, \ / CH» 

\ Ca /

V</ \-/ I &

All alkaline earths titrate approximately stoichiometrically. The 

titration should proceed immediately upon addition of the indicator, as the color of the solution will fade upon standing. The 

optimum pH of the titration is 10.4 or above.
D: 17a-l HARDNESS 175

The salt Na2EDTA also reacts with iron, manganese, copper, 

lead, cobalt, zinc, and nickel. Heavy-metal interferences can usually be eliminated by complexing the metals with cyanide. In the 

presence of cyanide, the procedure can be used with undiluted samples for analysis of water having metal concentrations as high as 

10 ppm iron, 10 ppm copper, 10 ppm zinc, 10 ppm lead.

The higher oxidation states of manganese above Mn+2 react 

rapidly with the indicator to form discolored oxidation products. 

Hydroxylamine hydrochloride reagent is used to reduce manganese 

to the divalent state. The divalent-manganese interference can be 

removed by addition of 1 or 2 small crystals of potassium ferrocyanide.

In the presence of high aluminum concentrations, a characteristic effect will be observed as the end point is approached. The 

blue color that indicates that the end point has been reached will 

appear and then on short standing will revert to red. This reversion should not be confused with the gradual change that normally 

takes place in the titrated sample several minutes after the titration has been completed. Although it has been stated that the titration cannot be used in the presence of excessive amounts of heavy 

metals, it should be pointed out that most heavy metal interference 

can be alleviated by applying conventional separations described in 

quantitative-analysis texts such as Hillebrand and Lundell (1929, 

p. 487).

Results are accurate and reproducible to ±0.1 mg.

Additional information on the principle of the determination is 

given by Goetz, Loomis, and Diehl (1950) and Botha and Webb 

(1952).

Apparatus and reagents

Titration assembly: Some analysts prefer to use conventional lighting and hand 

stirring. Others have reported better results by using visual-titration 

assembly consisting of a motor-driven stirrer, 25-ml buret, white porcelain 

base buret holder, and shaded incandescent lamp. The sample beaker ia 

placed near the front of the porcelain base and the reaction is viewed 

diagonally downward through the side of the beaker and against the 

white background. Illumination is from behind the beaker and in the 

plane of vision. The photometric titration assembly described in sec. C: 2h 

may also be used.

Buret, 25-ml

Hydroxylamine hydrochloride, 3 percent

Ammonium hydroxide, cone (sp gr 0.900)

Sodium cyanide, 2.5 percent

Potassium ferrocyanide, crystals

Eriochrome Black T indicator solution

Na-EDTA, 1.00*nil=1.00 mg CaCO3
176 ANALYTICAL PROCEDURES Dll7a-2

Procedure

1. Pipet a volume of sample containing less than 25 mg hardness (50.00 ml max) 

into a 150-ml beaker and adjust the volume to approx 50 ml.

2. Insert the beaker in the titration assembly and start the stirrer.

3. Add 1 ml 3 percent NH2OH HC1.

4. Add 1 ml cone NH4OH (if NH4OH is not tightly stoppered it tends to lose 

strength, and 1 ml of weak NH4OH will not buffer the solution to the 

desired pH).

5. Add 2 ml 2.5 percent NaCN (CAUTION deadly poison). NaCN may be 

eliminated if copper, zinc, lead, cobalt, and nickel are entirely absent and 

if the sample contains less than 0.25 mg Fe and 0.025 mg Mn.

6. If manganese is present, add 1 or 2 small crystals K4Fe(CN) 6 -3H2O. Stir 

and wait at least 5 min until the Mn2Fe(CN) 6 precipitates.

7. Add 2.0 ml Eriochrome Black T indicator.

8. Titrate with Na2BDTA (1.00 mlol.OO mg CaCO3 ) until blue or purple swirls 

begin to show. The end point is reached when all traces of red and purple 

have disappeared and the solution is clear blue in color. The detection of 

the end point may be facilitated by comparison of the titration solution 

with a color-free blank prepared of metal-free water and the reagent.

Calculations

ppm Hardness as CaCOa^-; : X r-2 :-X ml titrant density ml sample

Report hardness of < 1,000 ppm to whole numbers and of > 1,000 

ppm to 3 significant figures only.

Preparation of reagents

Hydroxylamine hydrochloride, 3 percent: Dissolve 30 g NH2OH-HC1 in metalfree water and dilute to approx 1 liter.

Sodium cyanide, 2.5 percent (CAUTION NaCN is a deadly poison and the 

reagent solution must be so marked) : Dissolve 2.5 g NaCN in metal-free 

water and dilute to approx 100 ml.

Eriochrome Black T indicator solution: Dissolve 0.40 g Eriochrome Black T 

in 100 ml metal-free water and dilute to approx 1 liter with 95 percent 

ethyl alcohol. This indicator is stable for at least 2 months. The Eastman 

reagent product has been found to be satisfactory.

Na2EDTA, 1.00 mlol.OO mg CaCOs: Dissolve 3.72 g Na2EDTA, which has been 

dried overnight in an H2 SO4 desiccator, in metal-free water and dilute to 

1,000 ml. The reagent is stable for several weeks and a larger volume is 

usually prepared. Check the titer of the reagent by titrating 25.00 ml 

CaCl2 (1.00 mlol.OO mg CaCO3 ), as described in the procedure for sample 

analysis.

Calcium chloride, 1.00 ml = 1.00 mg CaCO3 : Suspend 1.000 g CaCOs, dried at 

180°C for 1.0 hr before weighing, in approx 600 ml metal-free water and 

dissolve cautiously with a minimum of dilute HC1. Dilute to 1,000 ml.

D:17a-2 CALCULATION METHOD

Principle of determination

Hardness is computed from the individual determinations of the 

alkaline earths. The procedure does not require the determination of each constituent because of the interference of strontium
D:17b-l HARDNESS 177

and barium in the calcium determination. Strontium is included in 

the calcium values determined by any of the methods given in sec. 

D: 8. Its inclusion is not strictly stoichoimetric or quantitative, 

but for practical purposes can be assumed to be so. Barium interferes with detection of the complexometric calcium end point. 

Hence, a good end point is a reasonable indication that barium is 

not present in significant quantities. When barium is present in 

significant quantities it should be determined separately by the 

procedure given in sec. D: 5 and included in the calculated hardness. Some barium may be included with calcium in the singleprecipitation permanganimetric calcium determination (sec. D:8a-2). 

Therefore, if much barium is present and this method is used for 

calcium, double precipitation is recommended.

Calculations

ppm Hardness as CaCOa = (2 epm Ca+Mg+Ba) X 50.05

Eeport hardness of less than 1,000 ppm to whole numbers and 

of more than 1,000 ppm to 3 significant figures only.

D:17b CARBONATE AND NONCATCBONATE HARDNESS

The hardness caused by the alkaline-earth equivalent of the alkalinity in a water is called carbonate hardness, and the remainder, 

if any, noncarbonate hardness. These terms approximate the terms 

"temporary hardness" and "permanent hardness," which are based 

on the fact that upon boiling hard water the bicarbonate is decomposed and most of the calcium corresponding to the bicarbonate is 

precipitated as calcium carbonate. The consumption of soap by 

water of a given hardness is normally the same whether the hardness is caused by carbonate or noncarb,onate hardness.

D:17b-l CALCULATION METHOD

Carbonate and noncarbonate hardness are computed from the 

hardness and alkalinity determinations. No negative values are 

reported. However, "negative noncarbonate hardness" will counteract "positive noncarbonate hardness" in a mixture of two or more 

waters. Hence in all calculations of averages concerned with a 

mixture of waters for which two or more analyses are available, 

carbonate and noncarbonate hardness of the mixture must be computed from the average hardness and average alkalinity and not by 

averaging the noncarbonate hardnesses of the individual samples.

Calculations

ppm Noncarbonate hardness as CaCO3=50.05 (epm hardness epm alkalinity)

Computed "negative noncarbonate hardness" is reported as "0." 

Report noncarbonate hardness of < 1,000 ppm to whole numbers 

and of > 1,000 ppm to 3 significant figures only.
178 ANALYTICAL PROCEDURES D:17b-l

REFERENCES

American Public Health Association and others, 1955, Standard methods for 

the examination of water, sewage, and industrial wastes: New York, Am. 

Public Health Assoc., Inc., 10th ed.

American Society for Testing Materials, 1954, Manual on industrial water: Spec. 

Tech. Pub. 148-A.

Botha, C. R., and Webb, M. M., 1952, The versenate method for the determination of calcium and magnesium in mineralized waters containing large 

concentrations of interfering ions: Inst. Water Engineers Jour., v. 6.

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

 1954, Water quality criteria : Pub. no. 3, Addendum no. 1.

Goetz, C. A., Loomis, T. C., and Diehl, Harvey, 1950, Total hardness in water 

the stability of standard disodium dihydrogen ethylenediamine tetraacetate 

solutions: Anal. Chemistry, v. 22.

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis: 

New York, John Wiley and Sons, Inc.

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline and 

alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60.
D:18 IODIDE

Iodide, like bromide, is a minor element in the earth's crust and is 

normally present in natural waters in only minute quantities. Measurable amounts may be found in some streams that receive industrial 

wastes, and some natural brines may contain rather high concentrations.

D:18a-l OXIDATION METHOD

The oxidation method is similar in substance to D 1246-53 T, ASTM 

(1954, p. 260-263) Manual on Industrial Water, and U.S. Salinity 

Laboratory Staff (1954) Handbook 60.

Principle of determination

Iodide in a buffered solution is oxidized with bromine to iodate, 

and the excess bromine is subsequently removed with sodium formate 

(Kolthoff and Sandell, 1952, p. 585-605).

'+GHBr

Iodine equivalent to the iodate is then liberated by addition of potassium iodide to an acid solution.

The liberated iodine is then titrated with standard thiosulf ate, using 

starch as the indicator.

I2+2S2Oa-2-> S4O

Iron, manganese, and organic material interfere with the basic 

reactions of the method, but their interference is removed by preliminary treatment with calcium oxide.

Kesults are accurate and reproducible to ±0.02 mg.

Apparatus and reagents

Iodine flasks 

Buret, 10-ml

Calcium oxide, anhydrous powder 

Methyl red indicator, 0.01 percent 

Sulfuric acid, 20 percent v/v 

Sodium acetate, 16.5 percent 

Acetic acid, 12.5 percent v/v 

Bromine water, saturated 

Sodium formate, 50 percent 

Potassium fluoride, KF-2H2O, crystals 

Potassium iodide, crystals 

Sulfuric acid, 20 percent v/v 

Sodium thiosulfate solution, 0.010# 

Starch indicator, stable

179
180 ANALYTICAL PROCEDURES D:18a-l

Procedure

1. Remove soluble iron, manganese, and organic matter by adding a slight excess 

of CaO to approx 400 ml of sample, shake, let stand about 1 hr, and filter 

through dry paper. Discard the first 75 ml of filtrate.

2. Pipet a volume of the filtrate containing less than 5.0 mg I"1 (100.0 ml max) 

into a 250-ml iodine flask and adjust the volume to approx 100 ml.

3. Prepare a blank of approx 100 ml water and carry it through the procedure 

with the sample.

4. Add 1 drop 0.01 percent methyl red indicator and barely acidify with 20 percent H2 SO4 .

5. Add 15.0 ml 16.5 percent NaCaH3O2 .

6. Add 5.0 ml 12.5 percent HCzHsOa.

7. Add sufficient Br2 water to produce a light-yellow color, mix, and allow to 

stand 5 min.

8. Reduce the excess Br2 by adding 50 percent NaOHO2 until the yellow tinge in 

the sample disappears, then add an excess of 1 ml.

9. Wash down the sides of the flask with a small amount of water and blow out 

Bra vapors with a syringe and a glass tube inserted through the mouth of 

the flask.

10. If any iron precipitates at this point, add 0.5 g KF-2H2O.

11. Add approx 1 g KI.

12. Add 10 ml 20 percent H2 S<X

13. Let stand 5 min in the dark.

14. Titrate the liberated I2 with 0.0102V Na2S2O3 , adding 2-3 ml of starch indicator 

as the end point is approached. Disregard the return of the blue color after 

the end point has been reached.

Calculations

ppm I-i=: * X 1>QOQ X21.15X[(ml titrant - ml blank) X #] 

^ density ml sample

where N= normality of thiosulfate.

Report iodide concentrations of <10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.

Preparation of reagents

Methyl red indicator, 0.01 percent: Dissolve 0.01 g water-soluble methyl red in

approx 100 ml of water. 

Sodium acetate, 16.5 percent: Dissolve 273.5 g NaC2H3O2 -3H2O in water and dilute

to 1,000 ml. 

Acetic acid, 12.5 percent v/v: Mix 125 ml glacial HC2H3O2 (sp gr 1.049) with water

and dilute to 1,000 ml. 

Bromine water, saturated: Add to approx 250 ml water slightly more liquid Br2

than will dissolve when shaken. Store in a glass-stoppered actinic-glass

bottle. 

Sodium formate, 50 percent: Dissolve 50 g NaCHO2 in hot water and dilute to

approx 100 ml. Prepare fresh daily. 

Potassium iodide crystals, I Of1 free: The KI can be tested for lOr1 by

dissolving about 0.1 g in 5 ml water, acidifying with 1 or 2 drops cone H2SO4

(sp gr 1.84), and adding 2-3 ml starch indicator. Immediate appearance of
D:18a-2 IODIDE 181

a blue color indicates the presence of ICV1 ; slow color formation is due to 

atmospheric oxidation.

Sodium thiosulfate solution, O.Ol/W: Dilute 100.0 ml 0.1^ Na2 S2O3 to 950 ml with 

carbon dioxide-free water and standardize against KIO3 as follows: Dry 

approx 0.5 g KIO3 for 2 hr at 180°C. Dissolve 0.3567 g in water and dilute 

to 1,000 ml. Pipet 25.00 ml of the KIO3 into a 250-ml iodine flask, then add 

successively 75 ml water and 0.5 g KI crystals. After solution is complete, 

add 10 ml 20 percent H2 SO4 . Allow the stoppered flask to stand 5 min in 

the dark, then titrate with Na2 S2O3 using 2 ml starch indicator as the end 

point is approached (light straw color).

0 25 Normality of Na2S2O3=-ml Na2S2 O3

Sodium thiosulfate, Q.1N: Dissolve 25 g Na2 S2O3 -5H2O in carbon dioxide-free 

water. Add 1 g Na 2CO3 and dilute to approx 1,000 ml.

D:18a-2 STARCH METHOD

The starch method is extremely sensitive and is useful for the determination of trace quantities of iodide. It is also useful as a screening 

test or semiquantitative test in higher concentrations to select the 

proper sample volume for method D: 18a-l.

Principle of determination

Iodide is oxidized to iodate with bromine water (ASTM, 1954, p. 93).

r'-f 3Br2 +3H2O-> ICV'+GHBr

The excess bromine is then neutralized with phenol, and the iodate is 

then reacted with added iodide to produce 6 atoms of iodine for each 

atom of iodide originally present:

IOr1+5I-1 +6H+1-> 3I2 +3H2 0

The blue color developed by the iodine with starch is compared visually 

by that developed in standard solutions of iodide.

Bromide does not interfere, but any agent that will oxidize iodide 

to iodine will interfere.

The test as usually run in test tubes is sensitive to ±0.0002 mg, but 

the use of Nessler tubes or other means to increase the optical viewing 

depth can increase the sensitivity proportionally.

Apparatus and reagents

Test tubes, matched, Pyrex

Boiling-water bath

Potassium iodide, 1.00 ml=0.10 mg I'1

Sulf uric acid, 10 percent v/v

Bromine water

Phenol, solid

Potassium iodide, 5 percent

Starch indicator, stable
182 ANALYTICAL PROCEDURES D:18a-2

Procedure

1. Pipet a volume of sample containing less than 0.02 mg I'1 into a test tube and 

adjust the volume to 10.0 ml.

2. Prepare a blank of dilution water and standards containing known amounts of 

I"1 (0.02 mg max) in a 10.0-ml volume.

3. Add 1 drop 10 percent H2SO4.

4. Add bromine water dropwise until a slight excess has been added, as indicated 

by a yellowish color.

5. Heat gently over a flame or immerse in a boiling-water bath for 5 min.

6. Add a few crystals of phenol and shake until the yellow Br2 color is entirely 

discharged. A white precipitate of tribromophenol may form if too much 

Br2 water was added in step 4.

7. Add 1.0 ml 5 percent KI.

8. Add 1.0 ml starch indicator solution.

9. Compare the color of the sample with that of the standards.

Calculations

ppm 1-1= 1 » 000 x 0.1 Xml standard 

vv ml sample

Preparation of reagents

Potassium iodide, 1.00 ml=0.10 mg I"1 : Dissolve 0.131 g KI, dried overnight in

a sulfuric acid desiccator, in water and dilute to 1,000 ml. 

Bromine water: Add sufficient Brz to approx 100 ml dilution water to produce

about 1 ml excess. Store in a dark bottle protected from light. Commercial

preparations of bromine water are satisfactory. 

Sulfuric acid, 10 percent v/v : Mix 10 ml cone HaSOu (sp gr 1.84) with water and

dilute to approx 100 ml. 

Potassium iodide, 5 percent: Dissolve 5 g KI in water and dilute to approx 100 ml.

REFERENCES

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

Kolthoff, I. M. and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline and

alkali soils: U.S. Dept. Agriculture, Agriculture Handb. 60.
D:19 IRON

Iron is one of the most abundant minerals in the earth's crust. It 

occurs in the dark-colored silicate minerals of igneous rocks and as 

sulfides and oxides. In sandstone, iron oxide and iron hydroxide are 

often present as cementing materials. Iron is also present as oxides 

and sulfides in shale. Some of the common humic complexes causing 

color in water may also be iron bearing.

Because iron is readily precipitated as the hydroxide, it seldom is 

one of the major constituents in water. The metal occurs in water in 

both the ferrous and ferric state. Ferrous iron is unstable in solution 

in the presence of oxygen.

2Fe+2+4HC(V1 +H2O+y2 O2-»- 2Fe (OH) 3+4CO2

If the sample is subjected to a strong reducing environment the reaction is reversed, and the solution may contain large quantities of 

ferrous iron. In the 6-8 pH range the amount of ferric iron in true 

solution is theoretically limited to the solubility of ferric hydroxide, 

about 4X10~10 to 5X10~6 ppm. At lower pH values, the solubility is 

greatly increased. Most of ferric iron in apparent solution in alkaline 

water is believed to be principally in the colloidal form (Kolthoff and 

Sandell,1952,p.6l7).

Concentrations greater than 1 ppm of iron are rare in alkaline 

surface water. However, higher concentrations are common in ground 

water and acid surface water.

Iron in concentrations much greater than 0.20 ppm is objectional 

in waters for public supply. The U.S. Public Health Service (1946) 

recommends that the sum of iron and manganese in drinking and 

culinary water on carriers subject to Federal quarantine regulations 

riot exceed 0.3 ppm. These limits are not based on toxicity but on 

esthetic and taste considerations. Iron tends to stain laundry and 

porcelain, and also it can be tasted in concentrations higher than about 

0.5-1.0 ppm. Livestock is sensitive to the taste of iron and may not 

drink water with a high iron content. Ninety-five percent of the 

waters that support good fish fauna in the United States have 0.7 ppm 

or less of iron (Ellis, 1937).

Industries' tolerance for iron varies, but concentrations exceeding 

1-2 ppm are generally not satisfactory (California State Water Pollution Control Board, 1952, p. 277). It has been established that iron 

in irrigation water is of no practical significance to plant growth or 

soil texture.

D:19a IRON

The category "iron" includes ferric and ferrous iron in the ionized 

form and in suspension as the colloidal hydroxide. In all probability

183
184 ANALYTICAL PROCEDURES D:19a-l

the iron bound in the common humic complexes is also determined. 

Ferrous iron in solution at time of analysis can be determined by the 

procedure modification given in sec. D: 19b-l, and ferric .iron can be 

computed as the difference.

D:19b-l BIPYRIDINE METHOD

Principle of determination

The bipyridine determination utilizes the reaction between ferrous 

iron and 2,2'-bipyridine that yields a red complex. Hydroxylamine 

hydrochloride reduces ferric iron to ferrous. The color develops 

immediately and is stable for several hours. The color intensity is 

independent of pH in the range 3 to 10.

Copper reacts with bipyridine at pH values greater than 6, but the 

method embodies pH adjustment to about 5.5 where copper does not 

show appreciable interference. Silver and bismuth must be absent. 

The procedure can be used with undiluted samples for analysis of 

water having metal and anion concentrations as high as:,

Sb ___ __

Cd _ _

Cr _ _

Co

Cii _ __ _

Interfering

CN _ _

CrO4 _ _

Ppm

on

50

on

10

10

Hg

Ni

Sn

W

7n

anions must be limited to

Ppm

_ _ _ 10

20

P04

Ppm

1

2

20

10

10

the following concentrations:

Ppm

20

20

At 520 m/*, the wavelength used in the determination, the natural color 

of the water contributes a measurable absorbancy and must be corrected for if the color exceeds about 50 units. This applies only to 

the "normal" type of yellow color found in water.

With the listed apparatus, results are accurate and reproducible 

to ± 0.005 mg.

Additional information on the principle of the determination is 

given by Moss and Mellon (1942, p. 862-865).

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength: 520 m/u 

Cells : 40-mm optical depth 

Phototube: Blue-sensitive 

Blank : Metal-free water plus reagents 

Initial sensitivity setting: 1 

Slit width : 0.1 mm (approx)
D:19a-l IRON 185

The following absorbancies have been observed: 

mff\ Fe Absorbancy

0. 0125 0. 265

. 0250 . 540

. 0500 1. 070

. 1000 2.140 

Iron chloride, 1.00 ml=r0.004 mg Fe 

Bipyridine, 0.2 percent 

Hydroxylamine-hydrochloric acid reagent 

Sodium acetate, 35 percent

Procedure

Samples for the determination of iron should be treated as directed 

in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.10 mg Fe (25.00 ml max) 

into a 50-ml beaker and adjust the volume to 25.0 ml with metal-free water.

2. Prepare a blank of metal-free water and sufficient standards, and adjust the 

volumes to 25.0 ml.

3. Add 1.0 inl 0.2 percent bipyridine solution. Color development at this stage 

indicates the presence of ferrous iron.

4. Add 2.0 ml hydroxylamine-hydrochloric acid reagent, mix, and allow the 

solution to stand 0.5 hr.

5. Add 2.0 ml 35 percent NaC2H3O2 and mix.

6. Determine the absorbancies of the test sample and standards against the blank 

and, when necessary, make correction for water color as directed in sec. 

C : la-2, method 1, to obtain true sample absorbancy.

Calculations

1. Determine mg Fe in sample from a plot of absorbancies of standards containing 

known amounts of constituent.

2. ppmFe = j X r-* r~ X mg Fe in sample density ml sample

Keport iron concentrations of < 1.0 ppm to 2 decimal places and of 

> 1.0 ppm to 2 significant figures only.

Preparation of reagents

Iron chloride, 1.00 ml=0.004 mg Fe: Dilute 10.00 ml FeCl3 (1.00 ml=0.400 mg Fe) 

to 1,000 ml with metal-free water containing 1 or 2 drops of cone HC1 

(spgrl.19).

Iron chloride, 1.00 ml=0.400 mg Fe: Weigh out 0.400 g analytical-grade iron 

wire which has been cleaned in dilute HC1, rinsed, and dried. Dissolve in a 

minimum of dilute HC1 and dilute to 1,000 ml.

Bipyridine, 0.2 percent: Dissolve 1.0 g a,a'-dipyridyl in metal-free water and 

dilute to approx 500 ml.

Hydroxylamine-hydrochloric acid reagent: Dissolve 100 g NH2OH HC1 in metalfree water. Add 40 ml cone HC1 (sp gr 1.19). Add 1 g BeSO4-2H2O. 

Dilute to approx 1 liter with metal-free water.

Sodium acetate, 35 percent: Dissolve 350 g NaC2H3O2 in metal-free water and 

dilute to approx 1 liter.
186 ANALYTICAL PROCEDURES D:19b 

I>:191t> FERROUS AND FERRIC IRON

The determined values of ferrous and ferric iron may not be truly 

representative of the sample at the time of collection because the oxidation states of iron are not completely stabilized by the preservative 

treatment prescribed in sec. A: 4d.

D:19b-l BIPYRIDINE METHOD

The method distinguishes between ferrous and ferric iron in solution 

at the time of analysis but does not distinguish between the ionized 

and nonionized ferric iron.

Principle of determination

Ferrous bicarbonate is relatively soluble in water, as is ferrous hydroxide. Therefore, the ferrous iron remains in solution in many 

natural waters and can be determined by direct reaction with bipyridine in the pH range 3-10 without addition of reductant.

The chemistry of the procedure, interferences, and reagents, are 

identical to those described in sec. D: 19a-l.

Apparatus and reagents 

Spectrophotometer, Beckman Model B :

Wavelength: 520 m/*

Cells: 40-mm optical depth

Phototube: Blue-sensitive

Blank: Metal-free water plus reagents

Initial sensitivity setting: 1

Slit width : 0.1 mm (approx)

The following absorbancies have been observed :

my Fe Ab&orlxincy

0. 0125 0. 265

. 0250 . 540

. 0500 1. 070

. 1000 2.140

Iron chloride, 1.00 ml=0.004 mg Fe 

Bipyridine, 0.2 percent 

Hydroxylamine-hydrochloric acid reagent 

Sodium acetate, 35 percent

Procedure

Samples for the determination of ferrous iron should be treated as 

directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.1 mg Fe+2 (25.00 ml max) 

into a 50-ml beaker and adjust the volume to 29.0 nil with metal-free water.

2. Prepare a blank of metal-free water and sufficient standards, and adjust the 

volumes to 25.0 ml.

3. Add 1.0 ml bipyridine solution.
D:19b-l IRON 187

4. Treat the blank and standards only as follows:

(a) Add 2.0 ml hydroxylamine-hydrochloric acid reagent, mix, and 

allow the solution to stand 0.5 hr.

(b) Add 2.0 ml 35 percent NaCzHsOis and mix.

5. Determine the absorbancies of the test sample and standards against the 

blank, and, when necessary, make correction for water color as directed in 

sec. C: la-2, method 1, to obtain true sample absorbancy.

Calculations

1. Determine mg Fe+2 in sample from a plot of standards containing known 

amounts of the constituent.

1 1 000 2. ppm Fe+2=-. X ^ rXmg Fe in sample ^ density ml sample ^

3. ppm Fe+3=ppm Fe [sec. D: 19a-l] ppm Fe+2

Report iron concentrations of <1.0 ppm to 2 decimal places and 

of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Iron chloride, 1.00 ml=0.004 mg Fe: Dilute 10.00 ml FeCU (1.00 ml=0.400 mg 

Fe) to 1,000 ml with metal-free water, containing 1 or 2 drops cone HC1 

(spgrl.19).

Iron chloride, 1.00 ml=0.400 mg Fe: Weigh out 0.400 g analytical-grade iron 

wire which has been cleaned in dilute HC1, rinsed, and dried. Dissolve in 

a minimum of dilute HC1 and dilute to 1,000 ml.

Bipyridine, 0.2 percent: Dissolve 1.0 g a,a'-dipyridyl in metal-free water and 

dilute to approx 500 ml.

Hydroxylamine-hydrochloric acid reagent: Dissolve 100 g NH2OH HC1 in metalfree water. Add 40 ml cone HC1 (sp gr 1.19) and dilute to approx 1 liter.

Sodium acetate, 35 percent: Dissolve 350 g NaC2H3Oj in metal-free water and 

dilute to approx 1 liter.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Ellis, M. M., 1937, Detection and measurement of stream pollution: U. S. Dept.

of Commerce, Bur. Fisheries Bull. 22. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

Moss, M. L., and Mellon, M. G., 1942, Colorimetric determination of iron with

2,2'-bipyridyl and with 2,2',2"-terpyridyl: Indus, and Eng. Chemistry, AnaL

Ed., v. 14. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public Health

Service Repts., v. 61, no. 11.

D:20 LEAD

Lead is only a minor element in most natural waters, but industrial or mine and smelter effluents may contain relatively large 

amounts of lead. Many of the commonly used lead salts are water 

soluble. Lead acetate is used in printing and dyeing operations; 

lead chloride and sulfate are used in the manufacture of some 

paints; lead nitrate is used in photography, dyeing, engraving, and 

in the manufacture of certain explosives.

Lead is a cumulative poison to humans and animals, but the 

individual sensitivity differs. Lead does not have to be in solution 

to be toxic. Reports on human tolerance for lead vary widely, but 

the U.S. Public Health Service (1946, p. 12) states that lead shall 

not exceed 0.1 ppm in drinking and culinary water on carriers subject to Federal quarantine regulations. The maximum safe concentration for animal watering has been reported to be 0.5 ppm 

(Pierse, 1938, p. 145). The toxicity of lead to fish results from 

a film of coagulated mucus which forms over the gills and causes 

suffocation. The toxicity to fish is decreased by water hardness; 

50 ppm of calcium has destroyed the toxic effect of 1.0 ppm lead 

(Ohio River Valley Water Sanitation Commission, 1950).

D:20a-l DITHIZONE METHOD 

Principle of determination

Dithizone forms an insoluble red chela ted complex with lead in 

basic solution. The complex is insoluble in water but soluble in 

chloroform or carbon tetrachloride. The color of the complex in 

organic solvent changes from green to red, depending on the quantity of lead present.

The procedure is extremely sensitive, so much so that measurable 

amounts of !ead are often picked up from glassware and reagents. 

All glassware must be soaked and rinsed with dilute nitric acid. If 

the analysis is made frequently it is advisable to reserve certain 

cleansed glassware for lead determinations. Reagents used in large 

quantities must be repeatedly extracted with a carbon tetrachloride 

solution of dithizone until all lead is removed. Some batches of 

dithizone may require purification; when necessary the reagent can 

be purified by the ammonia-extraction method (Sandell, 1944). A 

reagent blank is carried through the procedure to compensate for 

the lead picked up from other reagents and the glassware.

189
190 ANALYTICAL PROCEDURES D:203r-l

Double extraction at a rather carefully controlled pH is used to 

isolate lead from other metals which also react with dithizone. 

Copper and zinc are often present with lead and are particularly 

troublesome. In the first extraction, ammonium citrate and cyanide are used to complex many heavy metals; hydroxylamine inhibits the oxidation of dithizone by ferricyanide formed by iron. 

Copper also oxidizes dithizone in a basic medium, and the color 

developed with dithizone is bleached if the combined concentration 

of copper and zinc exceeds 0.02 mg. The color of the first dithizone 

extract is seldom, if ever, suitable for color comparison. The lead 

dithizone complex from the first extraction is then decomposed with 

dilute nitric acid, and the lead is taken up in the aqueous phase. 

This procedure separates lead from iron and copper whose dithizonate is more appreciably decomposed by acid. The pH of the 

aqueous phase is then adjusted, and the lead complexed with 

dithizone.

With listed apparatus and reagents, results are usually accurate and 

reproducible to about ± 0.002 mg.

Additional information on the principle of the determination is 

given by Sandell (1944, p. 283-284) and Clifford (1943, p. 26-53).

Apparatus and reagents

Separatory funnels, 100- to 125-ml capacity, washed free of lead with dilute HNO3. 

Spectrophotometer, Beckman Model B:

Wavelength: 510 m/t

Cells: 10-mm optical depth; plastic cells are unsatisfactory because they 

are attacked by the organic solvent.

Phototube: Blue-sensitive

Blank: Metal-free water

Initial sensitivity setting: 1

Slit width: 0.11 mm (approx)

The following absorbancies have been observed :

mg Pb Absorbancy

0. 00 0.11

.01 .71

.02 1.30 

Lead nitrate, 1.00 ml=10.0 mg Pb 

Lead nitrate, 1.00 ml=0.002 mg Pb 

Ammonium citrate, 50 percent 

Hydroxylamine hydrochloride, 20 percent 

Thymol blue indicator solution, 0.04 percent 

Ammonium hydroxide, cone (sp gr 0.900) 

Potassium cyanide, 10 percent 

Dithizone, 0.005 percent in CCU 

Ammonia-cyanide reagent, 10 percent 

Nitric acid, 1 percent
D:20a-l LEAD 191

Procedure

Samples for determination of lead should be collected and treated 

as directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.020 mg Pb (50.0 ml max) 

into a separatory funnel and adjust the volume to 50 ml with metal-free 

water.

2. Prepare a blank of 50.0 ml metal-free water and sufficient standards, and 

adjust volumes to 50 ml.

3. Add 10 ml 60 percent (NH4 ) 3C8H5O7.

4. Add 2 ml 20 percent NH2OH HC1.

5. Add 10 drops 0.04 percent thymol blue indicator solution and make alkaline 

with cone NH4OH.

6. Add 4 ml 10 percent KCN ( CA UTION use hood).

7. Carefully adjust the pH of the solution to 8.5-9 (green color) with 1 percent 

HNO3 (CAUTION use hood).

8. Immediately extract the solution with 5'inl portions of the dithizone solution, 

draining off each extract into another separatory funnel until the lead is 

completely removed and the dithizone retains its original green color.

9. Remove the lead from the combined CC14 extracts by shaking the mixture with 

20 ml 1 percent HNOa. Discard the CCU layer.

10. Dilute the samples and standards to 50.0 ml with 1 percent HNO3.

11. Add 4.0 ml 10 percent ammonia-cyanide reagent.

12. Add 5.0 ml dithizone and immediately shake for 1 min.

13. Transfer part of the CCU layer to the absorption cell and determine the 

absorbancy of samples and standards against the blank. Water or air 

bubbles in the CCU can cause serious errors.

Calculations

1. Determine mg Pb in sample from a plot of absorbancies of standards containing known amounts of constituent.

2. ppm Pb== * X 1 1'000 1 xmg Pb in sample ** density ml sample 6

Report lead concentrations of < 1.0 ppm to 2 decimal places and of 

> 1.0 ppm to 2 significant figures only.

Preparation of reagents

Lead nitrate, 1.00 ml=0.002 mg Pb: The reagent should be prepared fresh immediately before use as follows: Dilute 5.00 ml Pb(NO3 ) 2 (1.00 ml=10.0 mg 

Pb) to 1,000 ml with metal-free water. Dilute 10.00 ml of this solution to 

250.0 ml with metal-free water.

Lead nitrate, 1.00 ml=10.0 mg Pb: Dissolve 15.986 g dry Pb(NO3 ) 2 in metal-free 

water and 1 ml lead-free concHNO, (sp gr 1.42) and dilute to 1,000 ml.

Ammonium citrate, 50 percent: Dissolve 50 g (NH*) 3C«H5OT in approx 100 ml 

metal-free water. Make the solution ammoniacal (pH 8.5-9) and shake 

with successive portions of 0.005 percent dithizone until all the lead has 

been removed (green color in the final dithizone extract). Remove any 

excess dithizone by extracting with CCU.

Hydroxylamine hydrochloride, 20 percent: Dissolve 20 g NHjOH-HCl in approx 

65 ml metal-free water and make the solution alkaline to thymol blue with
192 ANALYTICAL PROCEDURES D:20a-l

cone NH4OH (sp gr 0.900). Add 5 ml 4 percent sodium diethyldithiocarbamate. After a few minutes extract with successive portions of CCU until 

no color is developed in the CCU layer when the extract is shaken with a 

dilute solution of a copper salt. Make the NH2OH-Hd solution just acid 

with HC1 and dilute to approx 100 ml.

Thymol blue indicator solution: Grind 0.1 g thymol blue with 21.5 ml O.Oltf 

NaOH. When solution is complete, dissolve to approx 250 ml with metalfree water.

Potassium cyanide, 10 percent: Dissolve 50 g KCN in metal-free water and 

dilute to approx 500 ml. Extract the solution with successive small portions 

of dithizone until all lead has been removed. Extract the dithizone from 

the aqueous layer with CHCU.

Dithizone, 0.005 percent: Dissolve 50 mg of diphenylthiocarbazone in approx 1 

liter CCU. Keep the solution tightly stoppered and store in the refrigerator.

Nitric acid, 1 percent: Dilute 10 ml cone HNO3 (sp gr 1.42) to approx 1 liter 

with metal-free water.

Ammonia-cyanide reagent, 10 percent: Dissolve 10 g KCN in approx 500 ml cone 

NH4OH (sp gr 0.900). Add 10 g citric acid and dilute to approx 1,000 ml 

with metal-free water.

REFERENCES

Clifford, P. A., 1943, Report on lead-dithizone method and its interferences : Assoc. 

Official Agr. Chemists Jour., v. 26.

Ohio River Valley Water Sanitation Commission, 1950, Metal Finishing Industries Action Comm., Subcomm. on Toxicities, Rept. 3.

Pierse, R. R., 1938, Some cases of interest met with in practice chronic lead 

poisoning [in cows] : Veterinary Rec., v. 50.

Sandell, E. B., 1944, Colorimetric determination of traces of metals: New York, 

Interscience Publishers Inc.

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health 

Service Repts., v. 61, no. 11.
D:21 LITHIUM

Lithium is present in some minerals but is not abundant in nature. 

From available information, most fresh waters rarely contain lithium concentrations in excess of 10 ppm, but larger quantities may 

be contained in brines and thermal waters. Lithium is used in 

metallurgy, medicinal water, and some types of glass and storage 

battei-ies. Waste from such industries may contain lithium.

D:21a-l FLAME-PHOTOMETRIC METHOD

The exact procedure used in the flame-photometric method is governed principally by the design and performance of the particular 

flame photometer used; hence, no specific directions can be given. 

See sees. C: Ib and C: 2e for a discussion of flame photometry and 

flame photometers. Many helpful suggestions are provided by the 

manufacturers of the individual instruments.

Keport lithium concentrations of <10 ppm to 1 decimal place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm 

to 3 significant figures only.

193

D:22 LOSS ON IGNITION

The loss of weight on ignition of residue on evaporation has frequently been taken as a measure of the organic matter in the water. 

By subtracting loss on ignition from the residue on evaporation a 

result for the so-called fixed residue is obtained. For some waters, 

the loss on ignition may be chiefly water that was not removed by 

drying at 180°C. In the majority of samples, the only value of 

the determination is to afford a check on the residue on evaporation 

and provide another reference point for comparison of the calculated dissolved solids. This latter value should lie somewhere between residue on evaporation and residue after ignition (Collins, 

1928, p. 247). (See sec. C: 8b, this report.)

D:22a-l GRAVIMETRIC METHOD

Apparatus

Radiator, consisting of a nickel dish and cover. A nichrome triangle is placed 

in the radiator to keep the evaporating dish containing the solids from 

coming in contact with the sides of the radiatoi.

Sulfuric acid desiccator.

Procedure

1. Determine the dissolved solids as directed in sec. D : 36a-l.

2. Preheat the radiator with a burner until the center of the bottom is 

cherry red.

3. Ignite the residue in the covered radiator for 3.0 min.

4. Cool in the sulfuric acid desiccator and weigh.

Calculations

ppm Loss on ignition=-7 X H rX[mg(R on E m.g(R on /)] & density ml sample &v &

where

R on JB=weight of residue on evaporation, and 

R on /=weight of residue after ignition.

Report loss on ignition of < 1,000 ppm to whole numbers and of 

> 1,000 ppm to 3 significant figures only.

REFERENCE

Collins, W. D., 1928, Notes on practical water analysis: U. S. Geol. Survey 

Water-Supply Paper 596-H.

195

D:23 MAGNESIUM

Magnesium is relatively abundant in the earth's crust. Ferromagmasian minerals in igneous rocks and magnesium carbonate in 

carbonate rocks are generally considered to be the principal sources 

of magnesium in natural waters. Carbon dioxide plays an important role in the solution of magnesium from both silicate and carbonate minerals. Water associated with granite and silicious sand 

may contain less than 5 ppm of magnesium, but water from dolomite or limestone rich in magnesium may contain 10-50 ppm. 

Sulfates and chlorides of magnesium are very soluble, and water 

in contact with such deposits may contain several hundred parts 

per million. Magnesium is widely used in industry, and industrial 

wastes may add appreciable quantities to streams and ground water.

It is not believed that water can contain toxic quantities of magnesium and remain potable. Salts of magnesium act as cathartics 

and diuretics; the U.S. Public Health Service (1946) recommends 

that magnesium should not exceed 125 ppm in drinking and 

culinary water on carriers subject to Federal quarantine regulations. High concentrations of magnesium may cause scouring diseases among livestock, but less than 5,000 ppm is reported to be 

harmless to cattle that have become accustomed to it (Frens, 1948).

Magnesium imparts the property of hardness to water and is 

therefore of concern to industrial users of water. Magnesium acts 

similarly to calcium in that it flocculates soil colloids and tends 

to maintain good soil structure and permeability, hence it complements calcium in consideration of sodium ratios in irrigation water.

D:23a-l CALCULATION METHOD

The calculation method is applicable to all waters with which 

complexometric calcium and hardness determinations can be made 

satisfactorily.

Principle of determination

Equivalents per million of hardness is calculated from the hardness determination (sec. D: I7a-l). The equivalents per million of 

calcium is subtracted and the remainder multiplied by the equivalent 

weight of magnesium. Strontium does not interfere because strontium is included almost quantitatively in either the complexometric 

or perrnanganimetric calcium determination. Barium interferes 

with both the complexometric and single-precipitation permanganimetric calcium determination but in different ways. (See sec. 

D: 8.) In the presence of barium, magnesium should be determined 

by the pyropliosphate method (see sec. D: 23a-2).

197
198 ANALYTICAL PROCEDURES D:23a-2

Nonrounded values of hardness and calcium should be used in 

the calculation.

With the listed apparatus and reagents, results are accurate and 

reproducible to about ±0.002 mg.

Calculations

1. epm Hardness =. 1 .. X rr r^^ r Xml titrantXU.01998 density ml hardness sample

2. epm Ca+2= J---X- . ^A?00 r X (mltitrant- ml blank) X 0.5X0.0499 density ml Ca+2 sample

3. ppm Mg+2=12.l6 X (epm hardness epm Ca+2 )

Report magnesium concentrations of <10 ppm to 1 decimal place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm to 

3 significant figures only.

D:23a-2 PYROPHOSPHATE METHOD

The pyrophosphate method is normally used in the analysis of 

brines, some acid waters, and concentrated industrial wastes in which 

calcium and hardness cannot be determined satisfactorily by complexometric methods.

Principle of determination

Diammonium hydrogen phosphate precipitates magnesium in ammoniacal solutions; the precipitate is ignited to magnesium pyrophosphate and weighed.

Mg+2+NH4+1 +P<V3+6H2O-> MgNEUPO* 6H2O 

2MgNEUPO4£+ Mg2PaO7+2NH3+H2O

MgNH4PO4 is one of the most soluble precipitants in gravimetric 

analysis. A large excess of ammonium hydroxide serves to quantitatively precipitate the magnesium as the ammonium salt. Magnesium 

also forms slightly soluble MgHPO4 and Mg3 (POi) 2 , which do not 

yield pyrophosphate on ignition.

Calcium, strontium, and manganese form sparingly soluble phosphates and must be absent in the precipitation media. Iron and 

aluminum if present may precipitate as hydroxides, and coprecipitation of silica is a potential source of error. These interferences 

can be eliminated or sufficiently minimized if the magnesium determination is made on the filtrate of the gravimetric calcium determination. Some magnesium may be lost by precipitation as the 

phosphate during precipitation of the iron and aluminum in the 

calcium determination, if the concentration of phosphate and arseiiate exceeds that of iron and aluminum. If the original sample
D:23a-2 MAGNESIUM 199

contains an excess of phosphate or arsenate, sufficient iron may be 

added before precipitation as described in section D: 8a-3 to bring 

down all phosphate and arsenic and leave all magnesium in solution. Coprecipitation of various salts of magnesium, ammonium, 

different phosphates, chlorides, and oxalates in the presence of a 

large excess of (NH4 ) 2HPO4 and large amounts of ammonium salts 

is so prevalent that double precipitation is usually required for accurate results. The concentration of ammonium salts is decreased by 

ignition of the salts in the filtrate from the calcium determination. 

In the first precipitation, quantitative removal of magnesium is 

about all that can be achieved. The precipitate is then redissolved 

with a minimum of acid and the magnesium reprecipitated under 

conditions more favorable to a precipitate of the theoretical composition uncontaminated by coprecipitants.

Ignition temperature and time are important. If the precipitate 

is of proper composition it need not be ignited at temperatures above 

1,000°-1,050°C. However, the precipitate usually contains a little 

Mg(NH4 ) 4 (PO4 ) 2 which decomposes on ignition to Mg(PO3 ) 2 ; 5-10 

min final heating at 1,150°-1,200°C will complete the conversion to 

pyrophosphate. However, pyrophosphate slowly liberates phosphorus pentoxide at 1,200°C, hence heating at the elevated temperature 

should not be prolonged. Ignition for 30 min at 1,100° C is a practical 

compromise and is recommended for most water analyses.

Either porcelain or platinum crucibles may be used with proper 

precautions. Porcelain tends to soften somewhat at 1,100°-1,200°C, 

and repeated heating at this temperature embrittles the crucible. If 

porcelain is used it is advisable to allow the crucibles to cool a little 

before handling them with tongs. Platinum may be used satisfactorily if the precipitate is of normal composition and if the filter paper 

is adequately ashed before ignition at the high temperature. 

MgNH4PO4 can be ignited safely in platinum, but salts that decompose 

to Mg(PO3 ) 2 may cause corrosion of the platinum. The filter paper 

must be ashed completely in an oxidizing atmosphere before ignition 

to decomposition temperature. Adequate ashing is essential regardless of the crucible composition. Impurities in the precipitate lower 

the melting point of the pyrophosphate appreciably and may cause a 

fused mass at temperatures even lower than 1,100°C.

By careful work, results may be accurate and reproducible to ±0.05. 

Single precipitation is sometimes used with an attendant decrease in 

accuracy and reproducibility.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 352-363) and Hillebrand and 

Lundell (1929, p. 506-516).
200 ANALYTICAL PROCEDURES D:23a-2

Apparatus and reagents

Steam bath

Ice-water bath

Crucibles, platinum or porcelain

Muffle furnace

Hydrochloric acid, cone (sp gr 1.19)

Ammonium phosphate, 10 percent

Ammonium hydroxide, cone (spgr 0.900)

Methyl orange indicator solution

Hydrochloric acid, 1 percent v/v

Ammonium hydroxide, 8 percent v/v

Procedure

1. Dilute or concentrate the filtrate from the gravimetric Ca+2 determination 

(see sec. D: 8a-3, procedure 13) to a convenient volume. Take an aliquot 

containing less than 60 mg Mg+2 for the determination.

2. Evaporate the solution to dryness on a steam bath and decompose the ammonium salts by gentle ignition.

3. Take up the residue in approx 100 ml water containing several drops of 

cone HC1.

4. Place the beakers in the ice bath.

5. Add 10 ml 10 percent (NEU) 2HPO4.

6. Add cone NHiOH slowly and with vigorous stirring until the solution is alkaline to methyl orange and until a precipitate begins to form.

7. Stir occasionally for % hr.

8. Add 15 ml cone NH4OH and stir vigorously.

9. Allow the mixture to stand overnight.

10. Filter the supernatant solution through Whatman No. 42 filter paper, bm 

retain as much precipitate as possible in the beaker.

11. Quantitatively return the precipitate to the beaker by washing the filter paper 

3 times with 10-ml portions of warm 1 percent HC1. Adjust the volume to 

approx 100 ml with water.

12. Place the beaker in the ice bath.

13. Add 2-3 ml (NHOiHPO*

14. Add cone NH4OH slowly and with vigorous stirring until the solution is alkaline to methyl orange and the precipitate begins to form.

15. Add 5 ml cone NHiOH and stir vigorously.

16. Allow the mixture to stand for at least 4 hr.

17. Quantitatively collect the precipitate on Whatman No. 42 filter paper.

18. Rinse the beaker and precipitate 3-5 times with 5- to 10-ml volumes of 8 

percent NH4OH.

19. Transfer the precipitate and filter to the crucible.

20. Completely ash the filter paper over an oxidizing flame. Heat the crucible 

gently at first and gradually increase the heat as the ashing proceeds.

21. Ignite the crucible and contents at 1,100°C for 0.5 hr.

22. Cool in a desiccator and weigh. Record the weight to the nearest 0.0001 g.

Calculations

ppm Mg+2= X ' X mg Mg2P2O7 X0.2185 density ml sample
D:23a-3 MAGNESIUM 201

Report magnesium concentrations of <10 ppm to 1 decimal place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm to 

3 significant figures only.

Preparation of reagents

Ammonium phosphate, 10 percent: Dissolve 100 g (NH4 )2HPO4 in water and

dilute to approx 1 liter. 

Hydrochloric acid, 1 percent v/v: Mix 10 ml cone HC1 (sp gr 1.19) with water

and dilute to approx 1 liter. 

Ammonium hydroxide, 8 percent v/v: Mix 80 ml cone NH4OH (sp gr 0.900) with

water and dilute to approx 1 liter. 

Methyl orange indicator solution: Dissolve 0.05 g methyl orange in water and

dilute to approx 100 ml.

D:23a-3 ERIOCHROME BLACK T METHOD

The Eriochrome Black T method is not applicable to waters containing more than 100 ppm magnesium because of the small sample 

volumes required.

Principle of determination

Eriochrome Black T reacts with calcium, magnesium, strontium and 

other metals, especially zinc, copper, manganese, and aluminum, to 

give a red color. The sensitivity for a particular metal varies greatly 

with the pH. When the alkaline-earth metals alone are considered, 

the indicator can be made specific for magnesium with relative ease. 

The color development with magnesium is good in the pH range from 

9.5 to 12. Calcium and strontium do not give appreciable color until 

the pH rises to about 11. Calcium gives a much more intensive color 

than strontium; barium does not react at any pH value. Zinc, copper, 

manganese, and possibly other metals react over a broader pH range 

than magnesium. The reaction is specific for magnesium only in the 

absence of these metals.

For optimum reaction conditions the pH is adjusted to 10.8, at 

which point calcium interferes appreciably. This interference is 

eliminated by the use of a silicate-oxalate buffer. The oxalate complexes calcium and the silicate complexes aluminum. Permissible 

concentrations of other interfering ions are as follows:

TVT P Tl ff P Tl CkG f*

Aluminum

Ppm 

1

1

10

2

20

Residual chlorine

Beryllium 

rpm 

100

20

10

10

5
202 ANALYTICAL PROCEDURES D:233r-3

With listed apparatus, results are accurate and reproducible to 

± 0.005 mg.

Apparatus and reagents

Spectrophotometer. Beckman Model B: 

Wavelength: 520mp 

Cells: 10-mm corex 

Phototube: Blue-sensitive 

Blank : Distilled water plus reagents 

Initial sensitivity setting: 2 

Slit width: 0.5 mm (approx) 

The following absorbancies have been observed: 

mg Mg Absorbancy

0. 025 0. 54

. 050 1. 02

. 075 1. 50

. 100 1. 88

. 125 2. 00

Test tubes, glass-stoppered, graduated at 50 ml 

Magnesium sulfate, 1.00 ml=1.00 mg Mg+2 

Ammonium hydroxide, 2 percent v/v 

Briochrome Black T reagent 

Silicate-oxalate buffer

Procedure

1. Pipet a volume of sample containing less than 0.100 mg Mg+2 (5.00 ml max) 

into a glass-stoppered test tube and adjust the volume to 5.0 ml.

2. Prepare a blank of dilution water and sufficient standards, and adjust the 

volumes to 5.0 ml.

3. Add 1.0 ml 2 percent NH4OH.

4. Add 20 ml Briochrome Black T reagent. Run the solution down the side of 

the tube so that the reagent forms a layer on top of the sample.

5. Dilute to 50 ml with silicate-oxalate buffer.

6. Invert the stoppered test tubes several times until complete homogeneity, as 

evidenced by the absence of striations in the alcohol-water mixture, is 

achieved. Allow to stand at least 15 min but no longer than 30 min.

7. Determine the absorbancy of the test sample and standards against the blank, 

and when necessary make correction for water color as directed in sec. C : la2, method 1, to obtain the true sample absorbancy.

Calculations

1. Determine the quantity of Mg*2 in the sample from the plot of absorbancy of 

standards containing known amounts of the constituent.

2. ppra Mg+2== 1 .. X , 1>00() , Xmg Mg+2 in sample *^ B density ml sample B B

Report magnesium concentrations of <10 ppm to 1 decimal place 

and of > 10 ppm to 2 significant figures only.
D:23a-3 MAGNESIUM 203

Preparation of reagents

Magnesium sulfate, 1.00 ml=1.00 mg Mg+2 : Ignite MgSCVHzO in a muffle furnace 

at 300° C for 2 hr. Oool in a stoppered bottle in a sulfuric acid desiccator. 

Rapidly weigh a little more than 4.95 g into a tared weighing bottle. Return 

the bottle to the muffle furnace at low heat for y2 hr. Cool and adjust the 

salt weight to 4.950 g. Quantitatively transfer the dry salt to a 250-ml beaker. 

Cautiously add water and immediately cover; the hydration of the salt is 

rather violent. Add 1 or 2 drops cone HC1 (sp gr 1.19), a few drops of 

CHC1,, and dilute to 1,000 ml.

Ammonium hydroxide, 2 percent: Dilute 2 ml cone NH4OH (sp gr 0.900) to 

approx 100 ml.

Briochrome Black T reagent: Mix 0.40 g Eriochrome Black T with 100 ml dilution 

water and dilute to approx 1 liter with 95 percent ethyl alcohol.

Silicate-oxalate buffer: Dissolve 1.0 g Na2SiOs -9H2O in 500 ml water. Dissolve 

1.5 g Na2C2O4 in. approx 200 ml water. Mix the 2 solutions and dilute to 1,000 

ml.

REFERENCES

Frens, A. M., 1948, Salt drinking water for cows: Water Pollution Abs. 2i. 

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis: New

York, John Wiley and Sons, Inc. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public

Health Service Rept? , v. 61, no. 11.

D:24 MANGANESE

Manganese is much less abundant in nature than iron and is usually 

present in lower concentrations. Manganese may exist in natural 

waters in the bivalent or quadrivalent state. Manganous bicarbonate is relatively soluble but oxides readily to the quadrivalent 

form. Manganese dioxide and tetrahydroxide are nearly insoluble, 

and manganese in water under aerobic environment is probably carried primarily as a colloidal suspension. The reducing environment in bottom muds of lakes and reservoirs may promote solution 

of manganese from sediments. Manganese from natural sources 

seldom exceeds 1 ppm in water.

Waters containing less than 0.1 ppm seldom prove troublesome 

as public supplies, but those containing more than 0.5 ppm are 

reported to form objectionable deposits on cooked food, laundry, 

and plumbing fixtures. In concentrations that do not cause unpleasant taste, manganese is not regarded to be toxicologically significant. The U.S. Public Health Service (1946, p. 13) recommends 

that the concentration of iron and manganese combined not exceed 

0.3 ppm in drinking and culinary water on carriers subject to Federal 

quarantine regulations.

Tolerances for manganese in industrial water supplies are generally very low (less than 0.2 ppm), particularly for textile dyeing, 

food processing, distilling and brewing, paper, plastics, and photography. Manganese stimulates the growth of CrenotJirix and similar organisms that are troublesome in wells and water-supply systems. Crop tolerances to manganese differ widely; concentrations 

ranging from 0.5 to 500 ppm have been reported to be harmful 

(California State Water Pollution Control Board, 1952, 1954, p. 

291-292).

D:24a-l PERMANGANATE METHOD

The Permanganate method is quite sensitive and specific for manganese but does not distinguish the forms in which manganese was 

originally present in the water.

Principle of determination

Manganese may be oxidized by several oxidizing agents such as 

sodium bismuthate, ammonium persulfate with silver nitrate, and 

potassium periodate. Preference is usually given potassium periodate. In hot-acid solution, periodate oxidizes manganese as 

follows:

5I04-1 + 3H2O -»2MnO4 -1 + 5IQ,-1 + 6H+1

205
206 ANALYTICAL PROCEDURES D:24a-l

Phosphoric acid is needed in solution to prevent the possible precipitation of the periodates or iodates of manganese and to decolorize ferric iron (Kolthoff and Sandell, 1952, p. 681).

Few interfering ions affect the accuracy. No interference is offered 

by 500 ppm of chloride, 5 ppm of organic material as sugar, 436 ppm 

of methyl orange alkalinity, 100 ppm of ammonium salts as nitrogen, 

or 5 ppm of iron. Higher concentrations of chloride may reduce the 

permanganate. High concentrations of organic matter not oxidized 

by the nitric acid treatment may be attacked by the permanganate and 

the color thereby bleached. Charred organic material imparts a 

brownish tint to the solution. Interference of high concentrations of 

chloride and organic matter can be minimized by first evaporating the 

sample almost to dryness in the presence of hydrogen peroxide and a 

few drops of sulfuric acid. The concentrate is then diluted in a little 

water before step 3 of the procedure. The interference of organic 

matter can also be minimized by increasing the digestion time with 

nitric acid, or by the addition of an excess of potassium periodate. 

The evaporation is best carried out on a hotplate as the acid mixture 

tends to bump and spatter when concentrated over a flame.

Accuracy and reproducibility of results depend on the other constituents in solution. With the listed apparatus results may be reproducible and accurate with normal water to ±0.003 mg in the lower 

concentration range. With high permanganate color the final dilution 

error may amount to 2-3 percent.

Additional information on the principle of the determination is 

given by Sandell (1950, p. 312-314).

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength: 525 m/* 

Cells: 40-mm optical depth 

Phototube: Blue-sensitive 

Blank: Dilution water plus reagents 

Initial sensitivity setting: 1 

Slit width: 0.2 mm (approx) 

The following absorbancies have been observed: 

mg Mn Abaorbancy

0. 05 0. 31

.10 .63

. 25 1. 59 

Graduated cylinder, 25-ml 

Manganous sulfate, 1.00 inl=0.010 mg Mn 

Nitric acid, cone (sp gr 1.42) 

Phosphoric acid, 85 percent 

Potassium periodate, crystals
D:24a-2 MANGANESE 207

Procedure

Samples for the determination of manganese should be collected and 

treated as directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.25 mg Mn (100.0 ml max) 

into a 250-ml Erlenmeyer flask.

2. Prepare a blank of metal-free water and sufficient standards, and adjust the 

volumes to approx 100 ml.

3. Add 5 ml cone HNO3 .

4. Concentrate to approx 10 ml by boiling.

5. Cool and add 5 ml cone HNO3.

6. Add 5 ml 85 percent H3PO4 .

7. Boil for 3-5 min but do not allow any part of the bottom of the flask to become 

dry. If there is any evidence of charred organic materials, continue the 

digestion until the solution clears.

8. Add 10 ml dilution water.

9. Add approx 0.3 g KIO4 and boil until maximum MnO-T1 color is produced; 3 

min is usually sufficient. Add a little more KICK to check completeness of 

the color development.

10. Cool the solution and decant into a 25-nil graduated cylinder. Wash the 

contents of the flask with several small portions of water and decant as 

before, taking care not to transfer any crystals of KIO4 into the graduate. 

Adjust the volume to 25 ml.

11. Determine the absorbancy of the test sample and standards against the blank, 

and when necessary make correction for water color as directed in sec. 

C : la-2, method 3, to obtain true sample absorbancy.

Calculations

1. Determine mg Mn in sample from a plot of absorbancies of standards containing 

known amounts of the constituent.

2. ppm Mn=-j r X H r Xmg Mn in sample density ml sample

Report manganese concentrations of <1 ppm to 1 decimal place 

and of > 1 ppm to 2 significant figures only.

Preparation of reagents

Manganous sulfate, 1.00 ml=0.010 mg Mn: Heat 0.5 g MnSO4 -H2O for 1 hr at 

180°C. Dissolve 0.275 g in metal-free water containing 1 ml cone H2SO4 (sp 

gr 1.84) and dilute to 1,000 ml. Dilute 100.0 ml of this intermediate solution 

to 1,000 ml with metal-free water.

D:24a-2 TETRABASE METHOD

The tetrabase method is rapid, but it is limited for quantitative 

purposes to very low manganese levels, 0.02 ppm and below. Hence, 

the method is essentially a screening test to identitf y those samples that 

require quantitative determination by the method described in sec. 

D:24a-l. It can, however, be used to estimate up to 0.1 ppm 

manganese.
208 ANALYTICAL PROCEDURES D:24a-2

Principle of determination

Permanganate and manganese dioxide react with tetramethyldiaminodiphenylmethane (Arnold's base, tetrabase) to give an intensely 

blue oxidation product. With manganese concentrations above about 

0.02 ppm, the oxidation continues to a higher stage where a somewhat 

more stable yellow color is developed, but the stability of this color 

is not satisfactory for quantitative purposes. The method requires 

careful adjustment of the pH which is provided by the buffer. A 

short standing period eliminates negative tests.

Strong oxidizing agents interfere; persulfate, chromate, peroxide, 

and vanadate must be absent. However, false positive tests due to 

unsuspected oxidizing agents are rarely obtained. The method will 

tolerate residual chlorine up to 4 ppm. Chloride in concentrations up 

to 1,000 ppm does not interfere.

Additional information on the principle of the determination is 

given by Gates and Ellis (1947, p. 537) and Harry (1931, p. 434T).

Apparatus and reagents

Manganous sulfate, 1.00 nil=0.010 mg Mn 

Manganous sulfate, 1.00 ml=0.00002 mg Mn 

Test tubes, approx 1-cin diameter 

Ammonium hydroxide, 0.5 percent 

Potassium periodate, saturated solution 

Sodium acetate buffer 

Tetrabase solution

Procedure

Samples for the determination of manganese should be treated as 

directed in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.00002 mg Mn (1.00 ml max) 

into a test tube.

2. Prepare a blank of metal-free water, and sufficient standards, and adjust the 

volumes to 1.0 ml.

3. Add 1 drop 0.5 percent NH4OH.

4. Add 2 drops saturated KIO4 solution, and mix.

.">. Add 1 drop sodium acetate buffer. Mix and let stand for at. least 30 sec.

6. Add 1 drop tetrabase solution and mix.

7. Immediately compare the color of the Sample with that of the standards. 

The color is stable for only 2 or 3 min.

Calculations

ppm Mn= ~ r Xmg Mn 

^^ ml sample

Record the manganese concentration of the sample to 2 decimal 

places.
D:24a-2 MANGANESE 209

Preparation of reagents

Manganous sulfate, 1.00 ml=0.00002 mg Mn: Dilute 2.00 ml Mn (1.00 ml=0.010 

mgMn) to 1,000 ml.

Manganous sulfate, 1.00 ml=0.010 ing Mn: Heat 0.5 g MnSCVEUO for 1 hr at 

180°C. Dissolve 0.275 g in metal-free water containing 1 ml cone HaSO* 

(sp gr 1.84) and dilute to 1,000 ml. Dilute 100.0 ml of this intermediate 

solution to 1,000 ml with metal-free water.

Ammonium hydroxide, 0.5 percent: Mix 1.00 ml cone NH4OH (sp gr 0.900) with 

water and dilute to approx 200 ml.

Sodium acetate buffer solution: Dissolve 10 g NaC2H3O2 in approx 75 ml metalfree water. Add 3.0 ml glacial HC2H3O2 (sp gr 1.049) and dilute to approx 

100 ml. The pH of this solution is 4.8.

Tetrabase solution: Dissolve approx 10 mg tetramethyldiaminodiphenylmethane 

in approx 100 ml acetone. The solution is stable for about 2 weeks.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

 1954, Water quality criteria: Pub. no. 3, Addendum no. 1.

Gates, B. M., and Ellis, G. H., 1947, Jour. Biol. Chem., v. 168.

Harry, R. G., 1931, The determination of manganese: Soc. Chem. Industry 

Jour., v. 50.

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic analysis : New York, Macmillan Co., 3d ed.

Sandell, E. B., 1950, Colorimetric determination of traces of metals: New York, 

Inter science Publishers, Inc., 2d ed.

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public 

Health Service Repts., v. 61, no. 11.

D :25 NITROGEN

Nitrogen makes up about 78 percent by volume of normal air at 

20°C, and therefore water in equilibrium with air will contain 

about 15.6 ppm of dissolved nitrogen gas. Gaseous nitrogen is inert 

and not important in water chemistry, but the combined forms of 

nitrogen are. Although ammonia, nitrite, and nitrate are highly 

soluble, natural waters contain only small quantities, usually less 

than 10 ppm. Legumes take nitrogen from the atmosphere and fix 

it in the soil as nitrate. Most plant, animal, and bacterial life 

require nitrogen, and this nitrogen is constantly being converted 

from one form to another by tissue synthesis, metabolism, and decomposition. Bacteria are very important in converting the forms of 

nitrogen because some obtain their energy for life from these chemical reactions.

Both organic and inorganic nitrogen in water may result from 

the leaching of soils and rocks, from fertilizers, normal decomposition of plants and animals, sewage, industrial effluents, and living 

organisms.

D:25a AMMONIA NITROGEN

 Ammonia nitrogen includes nitrogen in the form of NH3 and 

NH4+1. As a component of the nitrogen cycle, it is often present 

in water, from natural causes, but usually in only small amounts. 

Ammonia is used in some water-treatment processes. More than 

0.1 ppm usually indicates organic pollution (Rudolph, 1931).

There is no evidence that ammonia nitrogen in water is physiologically significant to man or livestock. Fish, however, cannot 

tolerate large quantities. The toxicity to fish is directly related to 

the amount of undissociated ammonia in solution; hence, the toxicity 

is dependent on the pH of the water. Ammonia decreases the ability 

of hemoglobin to combine with oxygen, and the fish suffocate. Although the tolerances of fish differ, 2.5 ppm of ammonia nitrogen 

is considered harmful in the 7.4-8.5 pH range (Ellis, Westfall, and 

Ellis, 1946).

The low concentrations of ammonia in natural waters are of little 

industrial significance, except that ammonium salts are destructive to concrete made from portland cement.

D:25a-l NESSLERIZATION METHOD

The Nesslerization method is similar in substance to part IB, 

APHA (1955, p. 142-147) Standard Methods.

The method is the most sensitive of those given in this manual 

and is recommended for determination of low concentrations of less 

than 2 ppm.

211
212 ANALYTICAL PROCEDURES

Principle of determination

Ammonia is distilled from a buffered solution, and an aliquot of 

the distillate is then nesslerized. Essentially, nesslerization is the 

reaction between potassium mercuric iodide and ammonia which 

forms a red-brown complex of mercuric ammonobasic iodide:

HgI2 -2KI+NH3H>Hg(NH2 )I+2KI+HI

High concentrations of calcium (>250 ppm) interfere in the distillation by combining with the phosphate of the buffer and thereby 

lowering the pH. More than the prescribed amount of buffer can 

be added to maintain a pH of 7.4 ±0.2, or the buffer can be added first 

and then the pH adjusted with sodium hydroxide. Calcium, magnesium, iron, and sulfide interfere with the nesslerization, but the 

interference of the metals is eliminated by the distillation. Sulfide 

can be precipitated in the distillation flask with a little lead carbonate. 

Several amines, organic chloramine, acetone, aldehydes, and other 

undefined organic compounds are reported to form off-colors or a 

turbidity with Nessler reagent (APHA, 1955, p. 142-147). These 

compounds would normally be expected only in wastes or in grossly 

polluted waters, but when present in troublesome quantities, the procedure described in sec. D: 25a-2 is recommended for the determination of ammonia nitrogen.

All glassware should be rinsed with ammonia-free water.

With the listed apparatus, results are accurate and reproducible 

to ± 0.01 mg nitrogen.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 633-635).

Apparatus and reagents

Kjeldahl distillation apparatus, 1,000-ml flask 

Graduated cylinders, 50-ml 

Spectrophotometer, Beckman Model B :

Wavelength: 425 m/i

Cells : 40-nim optical depth

Phototube: Blue-sensitive

Blank : Annnonia-f ree water plus reagent

Initial sensitivity setting : 1

Slit width : 0.3 mm (approx)

The following absorbancies have been observed : 

mg N Abaorbancy

0. 02 0. 24

.04 .47

.06 .70

. 10 1. 16 

Phosphate buffer, O.oi/ 

Ammonium chloride, 1.00 ml=1.00 mg N 

Ammonium chloride, 1.00 ml=0.010 mg N 

Nessler reagent
D:25a-l NITROGEN 213

Procedure

Samples for the determination of ammonia nitrogen should be 

treated as directed in sec. A: 4d.

1. Free the distillation apparatus of ammonia by boiling ammonia-free water 

until the distillate shows no trace with Nessler reagent (CAUTION 

deadly poison).

2. Measure a volume of sample containing less than 1.0 mg ammonia nitrogen 

(500 ml max) into the distillation flask and adjust the volume to approx 

500 ml with ammonia-free water.

3. Add 10 ml phosphate buffer.

4. Immediately distill at a rate of not more than 10 ml nor less than 6 ml per 

minute; catch the distillate in a 200-ml volumetric flask.

5. Collect approx 190 ml of distillate, dilute to 200.0 ml with ammonia-free 

water, and mix.

6. Pipet an aliquot of distillate containing less than 0.1 mg ammonia nitrogen 

(50.00 ml max) into a graduate and adjust the volume to 50.0 ml with 

ammonia-free water.

7. Prepare a blank of ammonia-free water, sufficient standards, and adjust 

the volumes to 50.0 ml.

8. Add 1.0 ml Nessler reagent (CAUTION deadly poison) and mix.

9. Allow the solutions to stand at least 10 min but not over 30 min.

10. Determine absorbancy of test sample and standards against the blank.

Calculations

1. Determine quantity of NH3 (as N) in aliquot from a'plot of absorbancies of 

standards containing known amounts of constituent.

2. ppm Ammonia nitrogen as N=-: : density

1,000 ml distillate /mg N in\ 

' ml sample ml aliquot \ aliquot /

ppm Ammonia nitrogen as NH4+1=ppm as N X 1.288

Report ammonia nitrogen concentrations of <1 ppm to 2 decimal 

places and of between 1 and 10 ppm to 2 significant figures only.

Preparation of reagents

Phosphate buffer, 0.5M: Dissolve 14.3 g KH2PO4 and 90.1 g RHPOvSHzO in ammonia-free water and dilute to approx 3 liter.

Ammonium chloride, 1.00 ml=0.010 mg N : Dilute 10.00 ml NH4C1 (1.00 ml=a.OO 

mg N) to 1,000 ml with ammonia-free water. Prepare fresh.

Ammonium chloride, 1.00 ml=1.00 mg N : Dissolve 3.819 g NH4C1, dried overnight 

over sulfuric acid, in ammonia-free water and dilute to 1,000 ml.

Nessler reagent (CAUTION HgCla is a deadly poison and the reagent must be 

so marked) : Dissolve 61.75 g KI in approx 250 ml ammonia-free water. 

Cautiously add a sufficient quantity of a cold solution of HgCU, which has 

been saturated by boiling with an excess of the salt, to make the color a 

permanent bright red (about 400 ml) ; avoid an excess of HgCU. Dissolve 

the red precipitate by adding 0.75 g KI. Add 150 g KOH dissolved in 250 ml 

ammonia-free water and dilute to 1,000 ml. Allow the precipitate to settle 

and use the supernatant liquid.
214 ANALYTICAL PROCEDURES D:25a-2

D:25a-2 TITRATION METHOD

The titration method is recommended for the determination of 

ammonia nitrogen concentrations of above 2 ppm and for water containing substances that interfere with the Nessler reaction.

Principle of determination

Ammonia is distilled from a strongly alkalized sample and collected 

in boric acid. The ammonia is then titrated with standard acid. 

Cyanides may interfere by hydrolyzing to ammonia in presence of hot 

alkali:

ON-1+2H2O-» flCOO-1+NHS

This interference can, however, be greatly reduced if mercury salts are 

added before the alkali (Feigle, 1947, p. 183). Undissociated mercuric 

cyanide is stable to alkali.

Results are accurate and reproducible to ±0.1 mg nitrogen.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 633-635).

Apparatus and reagents

Kjeldahl distillation apparatus, 500-ml flasks

Buret, 25-ml

Boric acid, 2 percent

Sodium hydroxide, ION

Bromcresol green-methyl red indicator solution

Sulfuric acid, 0,0357#, 1.00 mloO.fiO mg N

Procedure

Samples for the determination of ammonia nitrogen should be 

treated as directed in sec. A: 4d.

1. Free the distillation apparatus of ammonia by boiling ammonia-free water 

until the distillate shows no trace with Nessler reagent ( CA UTION deadly 

poison).

2. Prepare a reagent blank of approx 250 nil ammonia-free water and carry it 

through the distillation procedure.

3. Pour approx 50 ml 2 percent H3BO3 into a 250-uil receiving flask. The tip 

of the delivery tube must be below the surface of the acid.

4. Measure a volume of sample containing less than 10 mg N (250.0 ml max) 

into the distillation flask, and adjust the volume to approx 250 ml.

5. Add 5 ml lOtf NaOH.

6. Digest the solution over low heat for at least 15 min.

7. Gradually increase the heat and collect approx 100 ml distillate. Check the 

last drops of the distillate with litmus paper to be sure that no more NH3 

is coming over.

8. Titrate blank and sample with 0.0357# H2 SO4 to identical colors, using 1.0 

ml mixed indicator.

Calculations

ppm Ammonia nitrogen as N==^ - X H r X0.5X (ml titrant ml blank) 

** & density ml sample

ppm Ammonia nitrogen as NH4+1=ppm as NX 1.288
D:25a-3 NITROGEN 215

Report ammonia nitrogen concentrations of between 1 and 99 ppm 

to 2 significant figures and of > 100 ppm to 3 significant figures only.

Preparation of reagents

Boric acid, 2 percent: Dissolve 20 g H3BO3 in ammonia-free water and dilute 

to approx 1 liter.

Sodium hydroxide, WN: Dissolve 400 g NaOH in ammonia-free water and dilute 

to approx 1 liter.

Bromcresol green-methyl red indicator solution: Dissolve 0.3 g bromcresol 

green and 0.1 g methyl red in approx 500 ml methyl alcohol.

Sulfuric acid, 0.0357^, 1.00 mloO.50 mg N: Mix 1.3 ml cone H2 SO4 (sp gr 1.84) 

with water and dilute to approx 950 nil before standardization. Standardize by titrating 25.00 ml 0.0357A7 Na*CO« to pH 4.5.

Sodium carbonate, 0.03572V: Dissolve 1.892 g primary standard Na2CO3 in carbon 

dioxide-free water and dilute to 1,000 ml.

D:25a-3 TEST-PAPER METHOD

The test-paper method is a qualitative screening test which is 

useful in detecting relatively large quantities of ammonia nitrogen.

Principle of determination

Ammonia is liberated from the sample by treatment with strong 

alkali and heat, and is detected by litmus paper. The sensitivity 

of the test is dependent primarily on the quality and age of the 

test paper. As little as 0.1 ppm has been detected with fresh litmus 

paper, but as much as 1 ppm has not been detected by other batches 

of paper. The sensitivity of the paper should be checked with 

standard ammonium solutions before the test is used. Cyanides interfere as with procedure D: 25a-2.

Additional information on the principle of the determination is 

given by Feigle (1947, p. 183).

Apparatus and reagents

Test tubes, 25-rnm 

Sodium hydroxide, WN 

Litmus paper, red

Procedure

1. Carefully measure approximately 25-ml sample into the test tube so as to 

leave the neck of the tube dry.

2. Add 2 ml WN XaOH well below the neck of the tube.

3. Immediately loop a piece of moist litmus paper below a stopper and insert 

stopper in the tube.

4. Mix the contents of the tube by swirling.

o. Place the tube in a water bath at 70°-80°C for 5 min.

6. A positive test is indicated by color change of the litmus.

Preparation of reagents

Sodium hydroxide, 102V: Dissolve 40 g NaOH in water and dilute to approx 

100 ml.
216 ANALYTICAL PROCEDURES D:25b 

D:25b NITRATE NITROGEN

Nitrate is usually the most prevalent form of nitrogen in water 

because it is the end product of the aerobic decomposition of organic 

nitrogen. Nitrate from natural sources is attributed to the oxidation of nitrogen of the air by bacteria and to the decomposition of 

organic material in the soil. Fertilizers may add nitrate directly 

to water resources. Nitrate concentrations range from a few tenths 

to several hundred parts per million, but in unpolluted water seldom exceed 10 ppm. Nitrate and chloride are major components 

of human and animal wastes, and the occurrence of abnormally high 

concentrations of both constituents suggests possible pollution of the 

water resources.

Cyanosis due to methemoglobinemia may occur in infants whose 

drinking or formula water contains a high concentration of nitrates. 

The nitrates, when ingested, are converted to nitrites in the digestive 

system of some infants. The nitrite ion oxidizes hemoglobin to 

methemoglobin and thereby causes cyanosis. It is widely recommended that water containing more than 10-20 ppm of nitrate, 

expressed as nitrogen, should not be used in infant feeding (Comly, 

1945).

Nitrates in large amounts are injurious to the dyeing of wool and 

silk and are undesirable in fermentation processes (California State 

Water Pollution Control Board, 1952, p. 301). At least 2 ppm of 

nitrate prevents intercrystalline cracking of steel in steam boilers.

D:25b-l PHENOLDISTTLFONIC ACID METHOD

The phenoldisulfonic acid method is similar in substance to part 

IB, APHA (1955, p. 149-151) Standard Methods.

Because of its sensitivity, the method is applicable for the determination of low nitrate concentrations, generally less than 30 ppm. 

It is not recommended for brines and other waters that contain 

more than 5,000 ppm of chloride or 0.2 ppm of nitrite.

Principle of determination

The water is evaporated to dryness and the residue mixed with 

1,2,4-phenoldisulfonic acid in the cold. Ammonium hydroxide is 

added until the nitrated sulfonic acid ring is converted to the yellow 

ammonium salt (see next page).

All acid samples must be neutralized or made slightly basic 

before evaporation to prevent loss as volatile nitric acid. Some 

samples lose nitrate rapidly if the heating is prolonged. Hence, 

the evaporation is carried almost to dryness on a steam bath; then 

the dish is removed from the bath and the remaining moisture 

allowed to evaporate.
D:25b-l

i. OH

-SOsH

O 2NNITROGEN

ONO2-'

217

OH

SO3H I SO3NH4

Chloride interferes by forming oxides of nitrogen in a strong 

acid solution, but reports of the permissible quantities differ widely; 

as little as 10 ppm is reported by one authority to interfere. Chloride can be precipitated with silver sulf ate, but the use of this salt 

presents its own problems. If an excess of silver sulfate is used, 

the unprecipitated silver ion produces an off color and (or) turbidity. 

As a practical compromise between potential chloride error and 

silver error, chloride is not removed unless the quantity in the volume 

of sample exceeds about 1 mg. The precision with which the silver 

sulfate can be measured is limited; therefore, it is not advisable 

to attempt to remove less than 0.1-0.2 mg of chloride from the 

sample volume. The inherent manipulative difficulties of voluminous 

precipitate and large volumes of nitrate decrease the accuracy of 

the determination, and for this reason the maximum volume of 

silver sulfate is generally limited to about 25ml. Nitrite in excess 

of 0.2 ppm erratically increases the apparent nitrate concentrations. 

Colored materials that physically modify the color system should 

be absent; the sample can be decolorized with alumina suspension.

Time is an important factor in the determination. Once the procedure is started it should be carried through to the development of 

the color with a minimum of delay between steps. The color is stable 

for several hours.

With the listed apparatus and reagents and in absence of interfering substances, results are generally accurate and reproducible to 

±0.005 mg of NCV1 (0.001 mg N). For some water the error may 

be appreciably greater.

Additional information on the principle of the determination is 

given by Chamot and Pratt (1909, 1911) and Taras (1950, p. 1020).

Apparatus and reagents

Evaporating dishes, porcelain

Steam bath

Graduated cylinders, 50-ml
218 ANALYTICAL PROCEDURES D:251>-1

Spectrophotometer, Beckman Model B : 

Wavelength: 410 m/tt 

Cells: 40-mm optical depth 

Phototube: Blue-sensitive

Blank: Dilution water carried through the procedure 

Initial sensitivity setting: 2 

Slit width : 0.2 mm (approx) 

The following absorbancies have been observed: 

mg NO3 Abaorbancy 

0. 08 0. 66 

. 16 1. 31

Silver sulfate, 1.00 mlol.OO mg Cl'1 

Aluminum hydroxide suspension 

Potassium nitrate, 1.00 ml=0.016 mg NOa"1 

Phenoldisulf onic acid 

Ammonium hydroxide, cone (spgr 0.900)

Procedure

If distribution of nitrogen in the nitrogen cycle is to be determined, 

the sample should be treated as directed in sec. A: 4d. Highly colored 

water may be decolorized with aluminum hydroxide suspension.

1. Pipet a volume of sample containing less than 0.16 mg NCV1 (10.00 ml max) 

into a small evaporating dish.

2. If the Or1 in the test sample exceeds about 1.0 mg, precipitate all but 0.1-0.3 

mg Cr1 with appropriate volume of Ag2 SO4 (1.00 mlol.OO mg Cr1 ). Cover 

the dish and set aside in a dark place overnight to obtain a filterable mass 

of precipitate, or the precipitate can be immediately flocculated with aluminum hydroxide suspension. Filter the suspension through Whatman No. 

42 filter paper and thoroughly Avash the precipitate with hot water, combining the filtrate and washings.

3. Prepare a blank and sufficient standards, and adjust the volumes to 10 ml.

4. Evaporate almost to dryness on a steam bath.

5. Within the next 5 min add 1.0 ml phenoldisulfonic acid to the cooled dry 

residue.

6. Immediately grind the residue and acid.

7. Cautiously dilute to approx 20 ml and mix.

8. Add dropwise and with constant stirring cone NH4OH until full color is 

developed, then add approx 1 ml excess.

9. Transfer the solution to a graduate, dilute to 50.0 ml, and mix. 

10. Determine the absorbancies of test sample and standards against the blank.

Calculations

1. Determine quantity of NOs"1 in test sample from a plot of absorbancies containing known amounts of the constituent.

2. ppm NOa-1 ^. 1 .. X 1>000 Xmg NOr1 in sample ^ density ml sample & ^

ppm Nitrate as N=ppm NO3'1 X0.226

Report nitrate (NOs"1 ) concentrations of <10 ppm to 1 decimal 

place and of > 10 ppm to 2 significant figures only.
D:25b-2 NITROGEN 219

Preparation of reagents

Silver sulfate, 1.00 mlol.OO mg Cl"1 : Dissolve 4.3972 g Ag2SO4, dried overnight 

over H2SO4, in water and dilute to 1,000 ml. Check the titer by titrating 

25.00 ml NaCl (1.00 ml=1.00 mg cr1 ) using K2CrO4 indicator (see sec. 

D : 10a-l). Store in lightproof bottle.

Sodium chloride, 1.00 ml=1.00 mg Cl"1 : Dissolve 1.6484 g NaCl, which has been 

dried at 900° C for V2 hour, in water and dilute to 1,000 ml.

Aluminum hydroxide suspension: Dissolve 125 g KA1(SO4 ) 2 -12H2O or NH4A1 

(SO4 ) 2 -12H2O in approx 1 liter water. Precipitate A1(OH) 3 by cautiously 

adding cone NH4OH (sp gr 0.900). Wash the precipitate by successive 

additions of water and decantation until the washes contain less than 0.04 

ppm NOr1 . Keep the solution basic during the washing.

Potassium nitrate, 1.00 ml=0.016 mg NOr1 : Dissolve 2.6100 g KNO$, dried overnight over H2SO4, in water and dilute to 1,000 ml. Dilute 10.00 ml of this 

intermediate solution to 1,000 ml.

1,2,4-phenoldisulfonic acid: (CAUTION use hood). Dissolve 113.5 g phenol 

(%-lb bottle) by placing the bottle in warm water. Pour the dissolved 

phenol slowly and with constant stirring into 766 ml cone H2SO4 (sp gr 

1.84). Slowly add 255 ml fuming H2SO4 (20 percent SO*) and stir well. 

Heat the mixture 2 hr on a steam bath. Care must be taken in the preparation of phenoldisulfonic acid to be sure that only 1,2,4-acid and not ortho 

or para mono acids are produced. The ortho and para acids give many 

side reactions with nitrate.

D:25b-2 REDUCTION METHOD

The reduction method is recommended for water that contains more 

than about 30 ppm of nitrate or more than 5,000 ppm of chloride or 

0.2 ppm of nitrite.

Principle of determination

Nitrate is reduced to ammonia by aluminum and zinc in alkaline 

solutions. The ammonia is distilled, collected in boric acid, and 

titrated with standard acid.

The sample must be free of ammonia or ammonium salts. The 

residue from the ammonia nitrogen determinations (see sec D:25a-l) 

can be used, or the ammonia can be volatilized from a fresh sample 

buffered to pH 7.4. Nitrite also is reduced to ammonia during the 

digestion, and the results must be corrected for significant quantities 

of nitrite. Nitrogenous organic matter is not readily decomposed in 

hot alkali and adds little, if any, nitrogen to a measured quantity of 

nitrate.

The reduction is not instantaneous, and recovery may not be quantitative if the digestion time is insufficient.

Results are accurate and reproducible to ±0.1 mg N (0.5 mg NCV1 ) 

Additional information on the principle of the determination is 

given by the Association of Official Agricultural Chemists (1945).

Apparatus and reagents

Kjeldahl distillation apparatus, 500-ml flask 

Buret, 25-ml
200 ANALYTICAL PROCEDURES D:25b-2

Phosphate buffer (see sec. D : 25a-l)

Boric acid, 2 percent

Devarda alloy

Sodium hydroxide, WN

Sulfuric acid, 0.0357N, 1.00 mlc=0.50 mg N

Bromcresol green-methyl red indicator solution

Procedure

If the distribution of nitrogen in the nitrogen cycle is to be determined, samples for the determination of nitrate should be treated as 

directed in sec. A : 4d.

1. Free the distillation apparatus of ammonia by boiling ammonia-free water 

until the distillate shows no trace of ammonia when tested with Nessler 

reagent (CAUTION deadly poison).

2. Dilute the residue of the ammonia nitrogen distillation (see sec. D:25a-l) 

to a convenient volume and transfer an aliquot containing less than 50 mg 

nitrate and nitrite combined (250 ml max) to a Kjeldahl flask and adjust 

the volume to approx 250 ml, or buffer a suitable volume of sample to pH 

7.4 with phosphate buffer and evaporate to approx 20 percent of the 

original volume to drive off ammonia nitrogen ; cool and adjust the volume 

to approx 250 ml with ammonia-free water.

3. Pour approx 50 ml 2 percent H3BO3 into a 250-ml receiving flash. The tip 

of the delivery tube must be below the surface of the acid.

4. Add approx 3 g Devarda alloy.

_ 5. Add 5 ml ION NaOH dawn the side of the flask and immediately connect the 

flask to the distillation apparatus.

6. Mix the contents of the flask by swirling.

7. Digest the solution over low heat for about 15 min.

8. Gradually increase the heat and collect approx 100 ml distillate.

9. Prepare a reagent blank of approx 250 ml ammonia-free water and carry

it through the distillation procedure.

10. Titrate the blank and 1 the sample with 0.0357N H2SO4 to identical colors, 

using 1.0 ml mixed indicator.

Calculations

XT'4. * AT f '

ppm Nitrate as N= -

-3 rr--^ -T~ 

Ldensity ml sample

X0.5 (ml titrant-ml blank) lppm' NO3 -1 =ppm as NX4.427

Eeport nitrate (NOs-1 ) concentrations of <999 ppm to whole 

numbers and of >999 ppm to 3 significant figures only.

Preparation of reagents

.Boric acid, 2 percent: Dissolve 20 g HsBO3 in ammonia-free water and dilute

to approx 1 liter. 

Bromcresol green-methyl red indicator solution : Dissolve 0.3 g bromcresol

green and 0.1 g methyl red in approx 500 ml methyl alcohol. 

Sodium hydroxide, ION : Dissolve 400 g NaOH in ammonia-free water and

dilute to approx 1 liter.
D:25c-l NITROGEN 221

Sulfuric acid, 0.0357N, 1.00 mloO.50 mg N: Mix 1.3 ml cone H2SO4 (sp gr 1.84) 

with water and dilute to approx 950 ml before standardization. Standardize by titrating 25.00 ml 0.0357tf Na2CO3 to pH 4.5.

Sodium carbonate, 0.0357N: Dissolve 1.892 g primary standard Na2COs in 

carbon dioxide-free water and dilute to 1,000 ml.

D:25c NITRITE NITROGEN

Nitrite is unstable in the presence of oxygen and is, therefore, 

absent or present in only minute quantities in most natural waters 

under aerobic conditions. The presence of nitrite in water is sometimes an indication of organic pollution.

Recommended tolerances of nitrite in domestic water supplies 

differ widely. A generally accepted limit is 2 ppm, but as little as 

0.1 ppm has been proposed (California State Water Pollution Control Board, 1952, p. 303). Nitrite is undesirable in water used in 

dyeing wool and silk and in brewing.

D:25c-l DIAZOTIZATION METHOD

Principle of determination

The diazotization method is based on a diazotization reaction 

with nitrite and sulfanilic acid followed by coupling with alpha 

napthylamne.

r N=N n+i

+ HONO HC1

S03H . SChH.Cl-'

r N=N -

_. SChH_ClF N=N NH 

I I

-N=N

+ HC1

303H NHj 

, Red azo dye

The development of the azo dye is affected by several variables 

including temperature, pH, standing time, and the degree of purity 

of the reagents. In addition, the azo dye is photosensitive, and, 

therefore, the solutions must be shielded from direct sunlight during 

the development of the color. Ordinary artificial light usually has 

no effect, although close fluorescent lamps may be detrimental. The 

effect of pH is probably the most significant. The diazotization 

step requires strong acid, while the coupling step proceeds best at 

moderate acidities. Therefore, at the completion of the diazotization 

stage, sodium acetate is introduced to raise the pH. At room 

temperature, the diazotization requires about 3 min for practical
222 ANALYTICAL PROCEDURES D:25C-1

completion. After 10 min, the diazotization product shows significant decomposition.

The common anions, with the exception of iodide, are not known 

to interfere. Iodide should not exceed 0.1 mg. Lead, zinc, and 

manganese are considered not to interfere. Copper interferes, and 

the quantity in the test sample should not exceed 0.05 mg. Disodium 

ethylenediamine tetraacetate effectively complexes iron to prevent 

its interference.

With the listed apparatus, results are accurate and reproducible 

to ± 0.0002 mg of NCV1 .

Additional information on the principle of the determination is 

given by Karrer (1950, p. 481-497) and Mellan (1941).

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength: 520 mji 

Cells: 10-mm optical depth 

Phototube: Blue-sensitive 

Blank : Distilled water plus reagents 

Initial sensitivity setting: 1 

Slit width : 0.2 mm (approx) 

The following absorbancies have been observed :

mg NO.,-1 Absorbancy

0. 01 0. 73 

. 02 1. 40 

Na2EDTA, 0.5 percent 

Sulfanilic acid solution, 0.6 percent 

Sodium nitrite, 1.00 ml=0.001 mg N 

Alpha napthylamine hydrochloride, 0.6 percent 

Sodium acetate buffer, 27 percent

Procedure

Samples for the determination of nitrite should be treated as directed 

in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.02 mg NCV1 (10.00 ml max) into 

a small beaker or test tube and adjust the volume to 10.0 ml.

2. Prepare a blank and sufficient standards, and adjust the volumes to 10.0 ml.

3. Add 0.2 ml 0.5 percent Na2EDTA solution

4. Add 0.2 nil sulfanilic acid and mix.

5. Let stand for 5.0 min.

6. Add 0.2 ml 0.6 percent alpha napthylamine hydrochloride.

7. Add 0.2 ml 27 percent sodium acetate buffer and mix.

8. Allow the liquids to stand at least 10 min and no longer than 20 min.

9. Determine the absorbancies of the test sample and standards against the 

blank, and, when necessary, make correction for water color as directed in 

sec. C : la-2, method 1.

Calculations

1. Determine quantity of NOT1 in test sample from a plot of absorbancies containing known amounts of the constituent.
D:25d-l NITROGEN 223

2. ppm NOr*=. * X 1 '000 xmg NCV1 in sample ^ density ml sample & ^

ppm Nitrite as N=ppm NO^X 0.304

Report nitrite (NCV1 ) concentrations of <1 ppm to 2 decimal 

places and of > 1 ppm to 2 significant figures only.

Preparation of reagents

Na2EDTA, 0.5 percent: Dissolve 5 g disodium salt in water and dilute to approx 

1,000 ml.

Sulfanilic acid (4-aminobenzenesulfonic acid), 6 percent: Dissolve 6 g sulfanilic 

acid in about 700 ml of hot water. Cool the solution. Add 200 ml cone HC1 

(sp gr 1.19). Dilute to approx 1,000 ml.

Sodium nitrite, 1.00 ml=0.100 nig N: Dissolve 0.4926 g NaNO2 in water and 

dilute to 1,000 ml.

Alpha-napthylamine hydrochloride: The commercially available reagent is of 

variable composition. The reagent should be purchased in small bottles, 

stored in a desiccator, and kept in the refrigerator. The pure reagent is 

white and comparatively odorless. If the reagent turns dark, it should 

be purified by recrystallization before use. To recrystallize 4 or 5 g., dissolve 

the material in 100 nil boiling water plus 5 ml cone HC1 (sp gr 1.19). Add 

2 g activated carbon, stir, and filter rapidly through a Buchner funnel fitted 

with filter paper. Cool in ice and collect the purified crystals. To prepare 

the solution, dissolve 0.6 g of the amine salt in water, add 1 ml HC1, and dilute 

to 100 ml.

Sodium nitrite, 1.00 ml=0.001 mg N: Add 1 nil CHC13 to 10.00 ml NaNO2 (1.00 

ml=0.100 mg N) and dilute to 1,000 ml with water. The standard should be 

used for the preparation of the calibration curve without delay.

Sodium acetate buffer solution, 27 percent: Dissolve 270 g NaC2HsO2 -3H2O in 

water and dilute to approx 1,000 ml.

D:25d ORGANIC NITROGEN

Organic nitrogen includes all nitrogenous organic compounds, 

such as amino acids, polypeptides, and proteins. It is present 

naturally in all surface waters as the result of inflow of nitrogenous 

products from the watershed and the normal biological life of the 

stream. Effluents of sewage and waste from slaughter houses and 

chemical plants often contain nitrogen in varying combinations. 

Organic nitrogen in unpolluted ground water is usually very low.

Organic nitrogen is not pathologically significant but is sometimes 

an indication of pollution. Organic nitrogen is important to considerations involving aquatic biology.

D:25d-l KJELDAHL METHOD

The Kjeldahl method is similar in substance to part I, APHA 

(1955, p. 155-156) Standard Methods.

Principle of determination

Organic nitrogen is degradated to the ammonium ion by digestion 

with sulfuric acid in the presence of copper sulfate, which acts as a 

catalyst. The solution is made alkaline with sodium hydroxide, and
224 ANALYTICAL PROCEDURES D:25d-l

the free ammonia is distilled off and Nesslerized. The color developed is proportional to the organic-nitrogen content

Nitrate and nitrite do not interfere. The effect of ammonium ions 

and ammonia is strictly additive. Therefore, the organic nitrogen is 

normally made on the residue of the ammonia nitrogen determination 

(see sec. D: 25a-l). If the residue from the method described in sec. 

D: 25a-2 is used, additional sulfuric acid should be added for the 

digestion.

Calcium, magnesium, iron, and sulfide interfere with the Nesslerization, but the interference of the metals is eliminated by the distillation. 

Sulfides can be precipitated in the distillation flask with a little lead 

carbonate before addition of sodium hydroxide. Several amines, 

organic chloramines, acetone, aldehydes, and other undefined organic 

compounds are reported to form off-color or turbidity with Nessler 

reagent (APHA, 1955, p. 155-156). These compounds would normally be expected only in wastes and polluted waters, but, when present in troublesome quantities, the ammonia can be collected in boric 

acid and titrated with standard acid rather than nesslerized (see sec. 

D: 25a-2). Cyanides are liberated during acid digestion and do not 

interfere.

All glassware should be rinsed with ammonia-free water. With 

the listed apparatus results for organic nitrogen in solution are accurate and reproducible to ±0.02 mg of nitrogen. The adequacy of 

results for total organic nitrogen in solution and suspension will depend on the completeness of the acid degradation, but will usually 

approach that obtained with soluble nitrogenous compounds.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell, 1952, p. 537-538).

Apparatus and reagents

Kjeldahl distillation apparatus, 1,000-ml flask 

Graduated cylinders, 50-ml 

Spectrophotometer, Beckman Model B :

Wavelength: 425 m/j.

Cells : 40-mm optical depth

Phototube: Blue-sensitive

Blank: Ammonia-free water plus reagents

Initial sensitivity setting: 1

Slit width : 0.3 mm (approx)

The following absorbancies have been observed : 

mg N Absorbancy

0. 02 0. 24

.04 .47

.06 .70

.10 1.16
D:25d-l NITROGEN 225

Phosphate buffer, 0.5M

Sulfuric acid, cone (sp gr 1.84)

Copper sulfate, 10 percent

Sodium hydroxide, ION

Ammonium chloride, 1.00 ml=:0.010 mg N

Nessler reagent

Procedure

Samples for the determination of organic nitrogen should be 

treated as directed in sec. A: 4d.

1. Free the distillation apparatus of ammonia by boiling ammonio-free water 

until the distillate shows no trace with Nessler reagent (CAUTION 

deadly poison).

2. Take the residue of the ammonia nitrogen determination for this determination. Or, buffer a volume of sample containing less than 1.0 mg N (500.0 

ml max) to pH 7.4 with phosphate buffer and evaporate to approx 20 percent of original volume to drive off ammonia nitrogen.

3. Oool and add 10 ml cone H2 SO4.

4. Add 1 ml 10 percent CuSO4 .

5. Digest under a hood until copious fumes are given off and the liquid becomes 

colorless or pale yellow.

6. Cool and dilute to approx 300 ml with ammonia-free water.

7. Add 50 ml 102V NaOH cautiously down the side of the flask.

8. Immediately connect the flask to the distillation apparatus and cautiously

mix the contents by swirling gently. 

!>. Distill at a rate of no more than 10 ml nor less than 6 ml per min; catch

the distillate in a 200-ml volumetric flask.

10. Collect approx 190 ml distillate, dilute to 200.0 ml with ammonia-free water, 

and mix.

11. Pipet an aliquot of distillate containing less than 0.1 mg ammonia nitrogen 

(50.0 ml max) into a graduate and adjust the volume to 50.0 ml with 

ammonia-free water.

2. Prepare a collorimetric blank of ammonia-free water and sufficient standards, and adjust volumes to 50.0 ml.

3. Add 1.0 ml Nessler reagent (CAUTION deadly poison) and mix.

4. Allow the solutions to stand at least 10 min but not more than 30 min.

5. Determine the absorbancies of the aliquot and standards against the colorimetric blank.

Calculations

1. Determine a reagent blank for each new batch of H2SO4 by taking 300 ml 

ammonia-free water through the entire procedure.

mg Reagent blank = mg N per 10 ml H2SO4 X2. Determine the quantity of nitrogen in the aliquot from a plot of absorbancies 

of standards containing known amounts of the constituent.

3. ppm Organic nitrogen as N=-5 rr-X i 2 r- Kl^ & & density ml sample

ml distillate P/mg N in\_/mg reagent\~| 

ml aliquot |_\ aliquot / V blank / J
22,6 ANALYTICAL PROCEDURES D:25d-l

Report organic nitrogen concentrations of <1 ppm to 2 decimal 

places, of. between 1 and 10 ppm to 1 decimal place, and of > 10 ppm 

to 2 significant figures only.

Preparation of reagents

Phosphate buffer, 0.5M: Dissolve 14.3 g KH2PO4 and 90.1 g K2HPO4-3H2O in 

ammonia-free water and dilute to approx 1 liter.

Copper sulfate, 10 percent: Dissolve 10 g CuSOi-SHcO in ammonia-free water 

and dilute to approy 100 ml.

Sodium hydroxide, 102V: Dissolve 400 g XaOH in water and dilute to approx 

1 liter.

Ammonium chloride, 1.00 ml=0.010 nig N : Dilute 10.00 ml NH4C1 (1.00 inl=1.00 

mg X) to 1,000 ml with ammonia-free water. Prepare fresh.

Ammonium chloride, 1.00 ml=1.00 mg X: Dissolve 3.819 g XH4C1, dried overnight over sulfuric acid, in ammonia-free water and cilute to 1,000 ml.

Xessler reagent (CAUTION HgCU is a deadly poison and the reagent must be 

so marked) : Dissolve 61.75 g KI in approx 250 ml of ammonia-free water. 

Cautiously add a sufficient quantity of a cold solution of HgCU, which has 

been saturated by boiling with an excess of the salt, to make the color a 

permanent bright red (about 400 nil) : avoid an excess of HgCl2 . Dissolve 

the red precipitate by adding 0.75 g KI. Add 150 g KOH dissolved in 250 

ml ammonia-free water and dilute to 1,000 ml. Allow the precipitate to 

settle and use the supernatant liquid.

REFERENCES

American Public Health Association and others, 1955, Standard methods for 

the examination of water, sewage, and industrial wastes: New York, Am. 

Public Health Assoc., Inc., 10th ed.

Association of Official Agricultural Chemists, 1945, Official and tentative methods of analysis of the Association of Official Agricultural Chemists: Washington, D. C., 6th ed.

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

Chamot, E. M., and Pratt, D. S., 1909, 1911, A study of phenoldisulfonic acid 

method for the determination of nitrate in water: Am. Chem. Soc. Jour., 

v. 31, 33.

Comly, H. H., 1945, Cyanosis in infants caused by nitrates in well water: Am. 

Med. Assoc. Jour., v. 129.

Ellis, M. M., Westfall, B. A., and Ellis, M. D., 1946, Determination of water 

quality: U. S. Fish and Wildlife Service Reserve Kept. 9.

Feigle, Fritz, 1947, Qualitative analysis by spot tests: Xew York, Elsevier Publishing Co. Inc.

Karrer. Paul, 1950, Organic chemistry: Xew York, Elsevier Publishing Co., 

Inc., 4th ed.

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganicanalysis : Xew York, Macmillan Co., 3d ed.

Mellan, I., 1941, Organic reagents in inorganic analysis: Philadelphia, Pa., 

The Blakistron Co.

Rudolph, Z., 1931, Principles of the determination of the physical and chemical 

standards of water for drinking, industrial, and domestic purposes: Water 

Pollution Abs. 4 [March].

Taras, M. J., 1950, Phenoldisxilfonic acid method of determining nitrate in 

water-photometric study: Anal. Chemistry, v. 22.
D:26 OILS AND WAXES

Potential contributors of oil to water are engaged in production, transportation, handling, and use of oil and petroleum products, such as ships, oil wells, oil-loading points, refineries, railroads, 

civic dumps, salvage dumps, garages, and many industries. Some 

oil in natural waters may be derived from the decomposition of 

plankton or higher forms of aquatic life (Galtsoff, 1936, p. 550- 

555). Heavy oils, greases, and waxes are somewhat insoluble in 

water, but they may be emulsified or saponified by detergents, alkalis, or other chemicals. Oil films interfere with gas exchange, coat 

bodies of birds and fish, exert a direct toxic action on some organisms 

as a result of water-soluble components, and interfere with fishfood organisms and the natural food cycle. Oils from surface films 

become adsorbed on clay particles, settle to the bottom, and remain 

a source of pollution because they may be stirred up and float again 

or release toxic principles (Galtsoff, 1936, p. 550-555).

The principal objection to oils in water for domestic use is the 

resulting disagreeable taste and odor. Aging causes petroleum 

odors to become musty, probably as a result of biochemical decomposition (California State Water Pollution Control Board, 1952, 

p. 30T).

Oils are detrimental to any industrial process water that is used 

in food and beverage preparation. The Portland Cement Association 

(1948) recommends that water used for mixing concrete be free 

of oil.

Oil pollution is harmful to marine life in at least four ways: 

through the formation of surface films, by emulsification, by sedimentation and coating of benthal organisms, and by toxic action of 

water-soluble principles (Gage, 1924, p. 237). Also, oil may affect 

the taste of fish.

D:26a-l WET-EXTRACTION METHOD

The wet-extraction method is applicable to waters containing 

small amounts of oil, grease, wax, and other diethyl ether-soluble 

substances in solution, suspension, or attached to sediment.

The wet-extraction method is similar in substance to part I 18B, 

APHA (1946, p. 42) Standard Methods.

Principle of determination

Dissolved or emulsified oils, greases, waxes, and other extractable 

substances are extracted from water by diethyl ether. The diethyl 

ether is evaporated and the residue dried and weighed. Volatile 

oils may be lost in the process.

227
228 ANALYTICAL PROCEDURES D:268r-l

Apparatus and reagents

Steam bath

Separatory funnel, "No-lub" glassware is recommended (one source of supply

is the Scientific Glass Co.) 

Diethyl ether

Procedure

Samples for the determination of oils and waxes must be treated 

in accordance with directions given in sec. A: 4d.

1. Mark the sample bottle at the water-plus-oil line and pour the treated 

sample into the separatory funnel.

2. Thoroughly wash the sample bottle with ether and add washings to sample 

in separatory funnel. Extract the sample with a sufficient number (minimum of 2) of 25-ml portions of diethyl ether.

3. The ether extractions are collected in a previously dried and tared evaporating dish or beaker.

4. Evaporate just to dryness on a steam bath.

5. Dry in a desiccator for 1 hr and weigh.

6. Determine the blank correction by weighing the residue of a portion of ether 

equal to that used for extracting the sample.

Calculations

ppm Oils and waxes= T-2 =-X (me sample residue me blank) ml sample

Report oil and wax concentrations of < 1,000 ppm to whole numbers 

;md of > 1,000 ppm to 3 significant figures only.

REFERENCES

American Public Health Association and American Water Works Association, 

1946, Standard methods for the examination of water and sewage: New 

York, American Public Health Assoc., Inc., 9th ed.

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

Gage, S. DeM., 1924, The control of oil pollution in Rhode Island: Boston Society Civil Engineers Jour. 11.

Galtsoff, P. S., 1936, Oil pollution in coastal waters: North Am. Wildlife 

Conf. Proc., v. 1.

Portland Cement Association, 1948, Design and control of concrete mixtures: 

Chicago, 111., 9th ed.
D:27 OXYGEN CONSUMED

The oxygen-consumed determination is a measure of the readily 

oxidizable material in the water and furnishes an approximation of 

the minimum amount of organic and reducing material present. 

In reality, the term "oxygen consumed" is defined by the method 

used for its determination. In the test given below it is defined as 

the amount of oxygen used by the sample when digested 30 minutes 

in a boiling-water bath with an excess of potassium permanganate 

in an acid solution. Oxygen consumed may be correlated with 

natural-water color or with some carbonaceous organic pollution 

from sewage or industrial wastes.

Tolerances for oxygen consumed in feed water for low- and highpressure boilers are 15 and 3 ppm, respectively (Northeast Water 

Works Association, 1940). Wash water containing more than 8 ppm 

has been reported to impart a bad odor to textiles; concentrations for 

water used in beverages and brewing range from 0.5 to 5.0 ppm (California State Water Pollution Control Board, 1952, 1954).

D:27fr-l PEBMANGANIMETBIC METHOD

Principle of determination

The sample is acidified and digested with potassium permanganate 

for 30 minutes in a boiling-water bath, during which time reducing 

substances are oxidized along with part of the carbonaceous material. The remaining permanganate is reacted with a volume of 

oxalate equivalent to the permanganate originally added, and the 

excess oxalate is determined by back titration with permanganate. 

The permanganate required in the back titration is therefore equivalent to the oxygen consumed. Nitrogen is not oxidized by the permanganate. The carbon in nitrogenous organic materials such as 

proteins is not as easily oxidized as that in simple carbonaceous 

materials, and the results may not therefore include the effect of all 

carbon. Similarly, some organic matter is fairly stable and may not 

react completely under the conditions of the test. Chloride will 

partially react with permanganate, and waters containing more than 

1,000 ppm of chloride will give significantly higher oxygen-consumed 

values.

In hot acid solution, permanganate slowly decomposes with the 

evolution of oxygen, and the blank must be run with the sample to 

evaluate the decomposition of permanganate; the blank also compensates for any oxidizable material in the acid.

To obtain reproducible results, it is essential that the directions be 

followed explicitly.

229
230 ANALYTICAL PROCEDURES D:27a-l

Additional information on the principle of the determination is 

given by APHA (1946, p. 122-124) and Treadwell and Hall (1935, 

p. 543).

Apparatus and reagents

Buret, 25-ml

Potassium permanganate, 1.00 mlc=0.100 nag Oz

Sulf uric acid, 25 percent v/v

Sodium oxalate, 1.00 mlol.OO ml KMnO4 (1.00 mloO.100 mg O2 )

Procedure

1. Pipet a volume of sample containing less oxygen-consuming material than 

can be oxidized by 7 ml KMnO4 (100.0 ml max) into a 250-ml Brlenmeyer 

flask and adjust the volume to approx 100 ml.

2. Prepare a blank of approx 100 ml dilution water.

3. Add 10.00 ml KMnO* (1.00 mloO.100 mg O3 ).

4. Add 10 ml 25 percent H2SO*, and mix.

5. Place the flasks in a boiling-water bath and digest for 30 min. The flasks 

must be at such depth that the surface of the sample is below the surface of 

the boiling water throughout the digestion period. This step is critical and 

determines the reproducibility of the determination. It is essential to maintain constant boiling conditions if determinations between different sets of 

samples are to be comparable. That is, the water bath should be at full boil 

before the samples are introduced. The water bath should be large enough 

to accommodate all the samples plus the blank. Porcelain rings can be 

placed around the necks of the flasks to weigh them down.

6. Remove the flask from the bath and cool to about 70°C.

7. Immediately add 10.00 ml Na2C2O4 (1.00 mlol.OO ml KMnO4 ), and mix.

8. Immediately back titrate with KMnO4 (1.00 mloO.100 mg O2 ) to the first pink 

tint that persists for 30 sec. The temperature during the titration should 

be between 55° and 60°C.

Calculations

ppm Oxvgen eonsumed = -: : X j-2 r-XO.l (ml titration ml blank) 

yv - density ml sample

Report oxygen consumed of < 10 ppm to whole numbers and of > 10 

ppm to 2 significant figures only.

Preparation of reagents

Potassium permanganate, 1.00 mloO.100 mg O2 : Dissolve 0.40 g KMnO4 in water 

and dilute to approx 1 liter. Store the solution in the dark for at least 1 

week to permit precipitation of uianganous salts. Heat the solution to 

boiling and filter through asbestos fiber or fritted glass. Cool and standardize against primary standard Na2C2O4 as follows: Dry about 0.1 g of the 

salt at 105°C for 1 hr. Dissolve approx 25 mg accurately weighed to 0.1 mg 

in 50 ml water and 20 ml 25 percent H2 SOi.

i T^IV/T ^ j Na2C2 O4 ml KMnO4 required=mg

U.OO/U

Add from the buret 90 percent of the required KMnO4 to the oxalate solution 

and heat to 55°-60°C. Continue the titration slowly and with constant
D:27ar-l OXYGEN CONSUMED 231

stirring until a faint pink color persists for 30 sec. For pfactical work, 

no blank correction is required. Store the standardized KMnO4 solution in 

a lightproof bottle. The titer of the KMnO4 solution should be checked 

frequently.

Sulfuric acid, 25 percent v/v : Mix 250 ml cone HZ SO< (sp gr 1.84) with water and 

dilute to approx 1 liter.

Sodium oxalate, 1.00 mlol.OO ml KMnO*: Dissolve 0.8376 g Na2C2O4, dried at 

105° C for 1 hr in water. Add 1 ml 25 percent H2 SO4 and dilute to 1,000 ml.

REFERENCES

American Public Health Association and American Water Works Association, 

1946, Standard methods for the examination of water and sewage: New 

York, Am. Public Health Assoc., Inc., 9th ed.

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

Northeast Water Works Association, 1940, Progress report, Committee on Quality Tolerances of Water for Industrial Uses: Northeast Water Works 

Assoc. Jour., v. 54.

Tread well, F. P., and Hall, W. T., 1935, Analytical chemistry: New York, John 

Wiley and Sons, Inc., 8th ed., v. 2.

D:28 OXYGEN DISSOLVED

Oxygen dissolved in water is derived from the air and from the 

oxygen grven off in the proce&s of photosynthesis by aquatic plants. 

The solubility of oxygen in water is dependent upon the partial 

pressure of oxygen in air, the temperature of the water, and the 

mineral content of the water.

Dissolved oxygen in water has no adverse physiological effect and 

actually increases the palatability of the water (California State 

Water Pollution Control Board, 1952, p. 241). In general, no 

minimum concentration of dissolved oxygen required to support fish 

life has been listed because the oxygen requirements of fish vary 

with the species and age, with temperature, and concentration of 

other substances in the water. Ellis believes that under average 

stream conditions, 3.0 ppm of dissolved oxygen, or less, should be 

regarded as hazardous or lethal, and that to maintain a varied fish 

fauna in good condition the dissolved-oxygen concentration should 

be at least 5.0 ppm (Ellis, 1937). Water supersaturated with dissolved oxygen has been reported as detrimental to fish (American 

Water Works Association Journal, 1938, p. 1420).

Dissolved oxygen is responsible for many of the corrosion problems in industry. For many industrial uses of water, zero dissolved 

oxygen would be desirable as a means of inhibiting corrosion.

D:28a-l ALSTERBEBG (AZIDE) METHOD

The Alsterberg (azide) method is similar in substance to part IIB, 

APHA (1955, p. 160, 255) Standard Methods, and D 888-49 T 

Non-Referee Method A, ASTM (1954, p. 243-245) Manual on Industrial Water.

This method is applicable to natural waters and most other waters 

which are not heavily polluted. It is not recommended for samples 

that contain more than 1 mg per liter of ferrous iron or appreciable 

quantities of sulfite, thiosulfate, polythionate, hypochlorite, or free 

chlorine.

Principle of determination

This method is a modification of the Winkler process for the 

determination of dissolved oxygen, and the principle is the same. 

The method depends on the formation of a precipitate of manganous 

hydroxide. The oxygen dissolved in the water is rapidly absorbed

233
234 ANALYTICAL PROCEDURES D:28a-l

by manganous hydroxide, forming a higher oxide, which may be in 

the following form:

MnSO4+2KOH-» Mn ( OH ) 3+K2SO4 

2Mn ( OH ) 2+ O2-> 2MnO ( OH ) 2

Mn(OH) 2 floe acts as a "gathering" agent for oxygen; therefore, the 

floe is passed twice through the solution to insure quantitative reaction. 

Upon acidification in the presence of iodide, iodine is released in 

a quantity equivalent to the dissolved oxygen present.

H+1

MnO(OH)2+2KI+H2O-> Mn(OH) 2+I2 +2KOH

The liberated iodine is then titrated with a standard sodium thiosulf ate solution using starch as the indicator.

Oxidizable organic matter consumes iodine and thereby causes 

low results. This effect of organic matter can be minimized if the 

liberation of iodine is followed immediately by the titration with 

thiosulfate. Any readily oxidizable or reducible constituents interfere by reaction with either the iodine or thiosulfate. Oxidizable 

substances cause low results and reducible substances high results. 

The magnitude of the net effect is, of course, proportionate to the 

absolute and relative concentrations of the different interferences.

In the listed procedure, sodium azide eliminates the interference 

of nitrite, and potassium fluoride overcomes the effect of ferric salts, 

provided the ferric iron concentration does not exceed 200 mg per 

liter and there is no delay in titration. The ferrous iron concentration should not exceed 1 mg per liter. High concentrations of 

suspended solids, which interfere, are removed by alum flocculation 

as directed in sec. A : 4d.

When the method given is used on water which is not heavily 

polluted the results are generally reproducible to ±0.01 mg, except 

that at low concentrations (below 0.1 nig) the error may be ± 50 

percent.

Additional information on the principle of the determination is 

given by Adams, Barnett, and Keller (1943).

Apparatus and reagents

Buret, 25-ml

Alum solution, 6 percent

Potassium fluoride solution, 25 percent

ManganoU^ suliate solution, 32 percent

Alkaline-iodide sodium azide solution

Sulfuric acid, cone (sp gr 1.84)

Starch indicator, stable

Sodium thiosulfate, 0.025tf
D:28a-l OXYGEN DISSOLVED 235

Procedure

Samples for the determination of dissolved oxygen must be collected and treated as directed in sees. A : 3c and A : 4d.

1. Pipet an aliquot of I2 solution from the treated sample (see sec. A:4d) 

that represents less than 2.0 mg O2 (200.0 ml max).

2. Titrate the liberated I2 with 0.025N Naa&Os to a pale straw color.

3. Add 1-2 ml starch indicator and continue the titration to the first disappearance of the blue color. Subsequent recoloration should be disregarded.

Calculations

1. If the sample has been treated with alum as in sec. A : 4d, correct the sample 

volume reacted with MnSO4 as follows :

, . _ 300 X ml taken for Speculation

ml taken for flocculation-f ml A1K(SO4)2

2. ppm02= -j-J X 1 1 '0°° 1 X (ml ^at<f fyot + 7) XmiNa.8AX0.2 

density ml sample ml treated aliquot

Report dissolved-oxygen concentration to 1 decimal place.

Preparation of reagents

Alum solution, 6 percent: Dissolve 10 g AlK(SO*)a'12H2O in water and dilute 

to approx 100 ml.

Potassium fluoride solution, 25 percent: Dissolve 40 g KF-2H2O in water and 

dilute to approx 100 ml.

Manganous sulfate solution, 32 percent: Dissolve 48 g MnS(V4H2O or 40 g 

MnSO4 -2HaO or 36 g MnS(VH,O in water, filter, and dilute to approx 100 ml.

Alkaline-iodide sodium azide solution : Dissolve 1 g NaN3 in 4 ml water. Add 

this solution with constant stirring to 95 ml alkaline iodide reagent. The 

reagent should not give a color with starch indicator when diluted and 

acidified.

Alkaline-iodide reagent: Dissolve 50 g NaOH or 70 g KOH and 13.5 g Nal or 

15 g KI in water and dilute to 100 ml.

Sodium thiosulfate solution, G.025N: Dissolve 6.205 g Na2S2O3 -5H2O in carbon 

dioxide-free water, add 1 g Na2CO3, and dilute to 1,000 ml. Store the thiosulfate in a glass-stoppered bottle which has been cleaned with dichromatesulfuric acid cleaning solution and rinsed with hot water. Standardize the 

Na2S3O2 against KIO3 as follows : Dry approx 1 g KIO3 for 2 hr at 180°C. 

Dissolve 0.8918 g in water and dilute to 1,000 ml. Pipet 25.00 ml of the 

KIOs into a 250-ml iodine flask, then add successively 75 ml water and 

2 g KI. After solution is complete, add 10 ml 20 percent HaSO*. Allow the 

stoppered flask to stand 5 min in the dark. Titrate with Na:S2O3 using 2 ml 

starch indicator as end point is approached.

Normality of Na2S2O3= A X>/.025 ml
236 ANALYTICAL PROCEDURES D:28ar-l

REFERENCES

Adams, R. C., Barnett, R. D., and Keller, D. E., Jri, 1943, Field and laboratory

determination of dissolved oxygen: Am. Soc. for Testing Materials Proc. 43. 

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

American Water Works Association Journal, 1938, Water bureau aids poor

fish: v. 30. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Ellis, M. M., 1937, Detection and measurement of stream pollution: U. S. Dept.

Commerce, Bur. Fisheries Bull. 22.
D:29 pH VALUE

The pH of a solution is a measure of the effective hydrogen-ion concentration, or more specifically, the hydrogen-ion activity («H+0-

pH=log - 

aH+i

In water solutions, deviations in pH from 7 are primarily the 

result of the hydrolysis of salts of strong bases and weak acids, 

or vice versa. Dissolved gases such as carbon dioxide, hydrogen 

sulfide and ammonia also affect the pH appreciably. The effect of 

dissolved salts of strong acids and strong bases is slight.

Because carbonates are prevalent in nature, most natural waters 

of the United States are slightly basic. The overall normal pH 

range is between 6.0 and 8.0. Industrial wastes on the other hand 

may be strongly acidic or basic, and the effect on the pH of water 

resources is largely dependent on the normal buffering capacity of 

the water.

The pH of potable water is not pathologically significant. Acid 

water with a pH of less than about 4 usually has a sour taste. The 

U.S. Public Health Service (1946) states that the maximum pH of 

treated drinking and culinary water on carriers subject to Federal 

quarantine regulations should be about 10.6. Optimum pH for fish 

is between 7.8 and 8.5 while acid water with pH below 4.4 and 

strong basic waters with pH greater than 8.8 generally causes gill 

irritation and death (Ellis, 1944). pH in conjunction with other 

factors affects the corrosion potential of water on metals. For 

detailed discussion the chemist is referred to the work of Langelier 

(1946, p. 169). Low pH also increases the corrosive action of water 

on concrete (Antill,'1937, p. 1803). Extremes in pH cannot usually 

be tolerated by industry. The optimum pH for irrigation water 

depends on the type of crops to be grown and on the physical and 

chemical properties of the soil. Acid soils will tolerate the more 

alkaline water and some alkaline soils will tolerate the more acid 

water.

D:29a-l INSTRUMENT METHOD

The instrument method is similar in substance to part I, APHA 

(1955) Standard Methods; D E70-52 T ASTM (1954) Manual on 

Industrial Water; and method 21a, U.S. Salinity Laboratory Staff 

(1954) Handbook 60.

Principle of determination

See sec. C: 2c for the principles of pH-meter operation.

The pH obtained in the laboratory may not be the same as that 

of the water at the time it was collected owing to reactions with 

sediment, hydrolysis, and oxidation taking place within the sample

237
238 ANALYTICAL PROCEDURES D:29a^2

bottle. Also the pH may change appreciably through loss of dissolved gases, the absorption of fumes in the laboratory, and from 

the deposition of calcium carbonate or other salts. Therefore, a 

value more representative of the pH at the time of collection is 

obtained if the determination is made as soon as the sample bottle 

is opened.

With the apparatus listed below, results are reproducible to 

±0.05 pH.

Apparatus and reagents

Beckman pH meter with glass or saturated calomel electrodes. 

Buffer solution, pH 7.00±0.02 at 25°C.

Procedure

1. After an appropriate warmup period, standardize the instrument with the 

buffer solution at the approximate temperature of the sample.

2. With a minimum of aeration and agitation, measure the pH in accordance 

with the instrument-manufacturer's instructions.

Calculations

For water having an abnormally high sodium content, a correction 

may be necessary. This correction will differ with the electrodes, 

hence the analyst is referred to manufacturer's instructions for the 

computations necessary.

Eeport pH of <4.5 to the nearest 0.05 and of >4.5 to 1 decimal 

place.

D:29a-2 FIELD INSTRUMENT METHOD

Beckman Model N and N-l and Coleman Model 20 pH meters 

have been used successfully.

D:29a-3 PAPER METHOD

The method is only semiquantitative and is useful only for fieldwork when high precision is not required. If fresh paper is used, 

accuracy of about ±0.5 pH can be obtained with the Fisher product.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

Antill, J. M., 1937, Resistance of concrete to chemical attack: Am. Water

Works Assoc. Jour., v. 29. 

Bllis, M. M., 1944, Water purity standards for fresh water fishes: U. S. Fish

and Wildlife 'Service Spec. Rept. 2. 

Langelier, W. F., 1946, Chemical equations in water treatment: Am. Water

Works Assoc. Jour., v. 38. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public

Health Service Repts., v. 61, no. 11. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline

and alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60.
D:30 PHENOLIC MATERIAL

Phenolic material in water resources is invariably the result of 

pollution. Phenols are widely used as disinfectants and in the synthesis of many organic compounds. Waste products from oil refineries, coke areas, and chemical plants may contain high concentrations. Fortunately, phenols decompose in the presence of oxygen 

and organic material, and their persistence downstream from point 

of entry is relatively short lived. The rate of decomposition is 

dependent on the environment.

Very low concentrations impart such a disagreeable taste to water 

that it is highly improbable that harmful amounts could be consumed 

unknowingly. Keported thresholds of detection of taste and odor 

range from 0.01 to 0.1 ppm. Chlorination of water containing 

phenols produce chlorophenols, which have much more disagreeable 

tastes and odor than the parent material. The California State 

Water Pollution Control Board (1952, p. 320) reports that thresholds of taste and odor for chlorophenols range from 0.001 to 0.02 

ppm. The U.S. Public Health Service (1946, p. 361) recommends 

that phenolic material should not exceed 0.001 ppm, reported as 

phenol, in drinking and culinary water on carriers subject to Federal 

quarantine regulations. Concentrations up to 1,000 ppm are not 

believed to be toxic to animals (Heller and Pursell, 1938, p. 99). 

Lethal concentrations for fish are related somewhat to the species, 

time of contact, temperature, and other conditions. Literature on 

the subject is voluminous. On the basis of compiled data, 5.0 ppm 

could be expected to be harmful to most fish, whereas 1.0 ppm or 

less is probably safe (California State Water Pollution Control 

Board, 1952, p. 320).

Phenols are undesirable in water supplies for food and beverage 

industries because of the attendant taste and odor problems.

D:30a-l GIBBS REACTION 

Principle of determination

The analytical method given here is not specific for phenol but is 

a collective determination of some of the simpler monohydroxy 

aromatic compounds. Because of the higher molecular weights and 

differences in light-absorbance characteristics of the various other 

materials determined, the reported value represents the minimum 

concentration by weight present. Phenolic materials are distilled 

from the acidified water sample, reacted with 2,6-dibromoquinone 

chloroimide (Gibbs reagent) in a buffered alkaline solution, and the 

resulting indigo-type dye extracted with n-butyl alcohol. Temperature and pH play an important role in the Gibbs reaction. A

239
240 ANALYTICAL PROCEDURES D:30a-l

borate buffer is employed to keep the pH at approximately 9.6. It 

is important that both samples and standards are at the same 

temperature during the reaction.

The extraction of the dye with n-butyl alcohol is not strictly 

quantitative, but is sufficiently linear if the volumes of the aqueous 

and alcoholic phases are uniform and if the extractions are carried 

to equilibrium. Although serial extraction quantitatively removes 

the dye, the color intensity of the combined extracts is always less 

than that of the first extract; hence serial extraction decreases the 

sensitivity of the test.

Different phenolic materials produce colors including green, blue 

green, blue, and purple. Therefore, arbitrary spectrophotometric 

comparison must be used as it would be difficult to prepare visual 

standards if the sample contained a complex mixture of phenols, 

cresols, chlorophenols, etc. The butyl alcohol extracts of the dye 

formed in Gibbs reaction for different phenolic materials have 

similar wavelengths of maximum absorbance. Accordingly, a wavelength of 670 m/* is used, and the results reported on the basis of 

phenol.

The sensitivity of the reaction is dependent on the batch of dye 

used. Absorbancies of about 0.4 and 0.8 read in 40-mm cells for 

0.01 mg of phenol have been observed with different batches of dye.

With the listed apparatus, results with pure phenol are accurate 

and reproducible to ±0-.0004 mg. Although good color is developed 

with Gibbs reagent at somewhat higher concentrations, an extension 

of the range above 0.030 mg for the determination with mixtures is 

not recommended.

Additional information on the principle of the determination is 

given by Ettinger and Buchhoft (1948) and Ettinger and Kroner 

(1949).

Apparatus and reagents

Spectrophotometer, Beckman Model B:

Wavelength: 670 m^

Absorption cells: 23-mm, 25-mm, or 40-mm optical depth

Phototube: Red-sensitive

Blank: Dilution water plus reagents

Initial sensitivity setting: 3

Slit width: 1 mm (approx) 

Distillation apparatus, 500-ml, with Graham condenser and ground-glass joint

and stopper (Corning No. 3360) 

Volumetric flask, 250-ml, glass-stoppered 

Incubator, 37°C (optional)

Multiple electrical stirrer or other agitation apparatus 

Separatory funnel, 500-ml 

Phosphoric acid, 10 percent
D:30a-l PHENOLIC MATERIAL 241

Methyl orange indicator solution

Pumice

Phenol, 1.00 ml=1.00 mg phenol

Borate buffer solution

Copper sulfate, 0.005 percent

2,6-Dibromoquinone chloroimide solution

n-Butyl alcohol

Sodium sulfate, anhydrous

Procedure

Samples for the determination of phenolic material must be preserved as directed in sec. A: 4d, and analysis should proceed as soon 

as possible after collection.

1. Measure a 250-ml volume of sample into the distillation apparatus.

2. Add 0.7 ml-10 percent H3PO4 and 2 drops methyl orange indicator solution. 

If the sample is still alkaline, add more EUPO* until it is just acid to 

methyl orange.

3. Add a small quantity of pumice to the flash to inhibit bumping.

4. Collect approx 230 ml distillate in a 250-ml glass-stoppered volumetric flask.

5. Allow the apparatus to cool.

6. Add approx 20 ml dilution water to the distillation flask and distill to the 

250.0-ml mark on the receiving flask.

7. Mix the distillate thoroughly, and pipet an aliquot containing less than 

0.025 mg phenolic material (200.0 ml max) into a 300-ml Brlenmeyer flask 

and adjust the volume to approx 200 ml.

8. Prepare a blank of dilution water and sufficient standards, and adjust 

volumes to approx 200 ml.

9. Bring aliquot, blank, and standards to the same temperature.

10. Add 10 ml borate buffer solution.

11. Add 2 ml 0.005 percent CuSO«.

12. Add 3.0 ml 2,6-dibromoquinone chloroimide solution.

13. Mix well and cover the flasks with inverted beakers.

14. Allow the solutions to stand overnight at room temperature or incubate 

for3-5hrat37°C.

15. Extract the solutions uniformly with 50.0 ml n-butyl alcohol. Vigorous 

stirring with an automatic stirrer for 5 min or shaking for 2 min has 

been found to be satisfactory.

16. Separate the aqueous and solvent phases in a separatory funnel.

17. Clarify the alcoholic extract by shaking it with about 0.5 g anhydrous 

Na2SO« and centrifuging.

18. Decant the clear dye extract into the absorption cells and determine the 

absorbancies of the extract of sample and standard against that of the 

blank.

Calculations

1. Determine the mg phenolic material in the aliquot from a plot of absorbancies 

of standards containing known amounts of phenol.

« , .. . . , 1 1,000 , ml distillate .. mg phenolic 2. ppm Phenolic materials- : X r-1 r~X ; n r-X A . ,

density ml sample ml aliquot material

Report phenolic material concentrations of <0.1 ppm to 3 decimal 

places and of >0.1 ppm to 2 significant figures only.
242 ANALYTICAL PROCEDURES D:30a-l

Preparation of reagents

Phosphoric acid, 10 percent : Dilute 118 ml of 85 percent H3P(>4 to approx 1 liter.

Methyl orange indicator solution: Dissolve 0.50 g methyl orange in water and 

dilute to approx 1 liter.

Phenol, 1.00 ml=0.001 mg phenol: Dilute 10.00 ml phenol (1.00 ml=1.00 mg 

phenol) to 1,000 ml. Dilute 10.00 ml of this intermediate solution to 100.0 

ml. This dilute solution is unstable and should be prepared immediately 

before use.

Phenol, 1.00 ml=1.00 mg phenol : Dissolve 1.000 g phenol in water and dilute to 

1,000 ml. Store in the refrigerator. The strength can be checked as follows : 

Pipet 50.0 ml of the phenol solution into a 250-ml glass-stoppered bottle containing 25 ml water. Add 50.0 ml 0.1N Br2 and 5 ml cone HC1 (sp gr 1.19). 

Allow to stand in cool water for 15 min with occasional shaking. Add about 

1 g KI crystals and shake vigorously. Titrate with 0.100JV Na2S2O3 until the 

I2 color disappears. The end point may be checked by adding a few drops of 

starch solution. The following reactions occur :

+3HBr 

Br2 (excess) +2 r'-J-I^Br-1

I2 +2 S20,-2-> S4CV+2 1'1 

Similarly treat a blank of dilution water.

mg Phenol per ml=0.03134 (ml blank ml sample)

If retained for long periods of time, the solution should be standardized once 

a month.

Bromine, Q.1N : Dissolve 2.784 g KBrO3 and about 10 g KBr in water and dilute to 

1,000 ml.

Sodium thiosulfate, 0.100JV: Dissolve 24.820 g Na2S2O3 -5 H2O and 1.5 g (NH4) 2COS 

in water, add 5 ml CHC13 as a preservative, and dilute to 1,000 ml. Standardize as follows : To 100 nil water in an Erlenmeyer flask, add 5 g KI, 10 ml 

10 percent H2SO4, and 25.00 ml 0.100N KH(IO3 ) 2. Titrate with the Na2S2O3 

using starch near the end point. The following reactions occur :

5 r'+ior'+eir1-^ i2+3 H=O

I2+2 S2O3-2-» S4O

OK A

Normality Na2S2O3= . * a ~ X 0.100 ml

Potassium biniodate, 0.100JW Dissolve 3.250 g KH(IO3 ) 2 in water and dilute to

1,000 ml. 

Borate buffer solution : Dissolve 6.200 g H3BO3 and 7.450 g KC1 in approx 1,500

ml water. Adjust with 0.2N NaOH (about 450 ml) to such a strength that

5 ml, when added to 100 nil of water, will give a pH of about 9.6. The pH

of the buffer will be about 10. 

Copper sulfate, 0.005 percent: Dissolve 0.05 g anhydrous CuSO4 in water and

dilute to approx 1 liter. 

2,6-Dibromoquinone chloroimide solution : Dissolve 0.01 g 2,6-dibromoquinone

chloroimide in 5 ml 98 percent ethyl alcohol and dilute to 50 ml with water.

This solution is stable for about 30 min.
D:30a-l PHENOLIC MATERIAL 243

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria: 

Pub. no. 3.

Ettinger, M. B., and Kroner, B. C., 1949, The determination of phenolic materials in industrial wastes: Industrial Waste Conf., 5th, Purdue Univ., 

Nov. 30.

Ettinger, M. B., and Buchhoft, C. C., 1948, Determination of phenol and structurally related compounds by the Gibbs method: Anal. Chemistry, v. 20.

Heller, V. G., and Pur sell, Lee, 1938, Phenol-containing waters and their physiological action: Jour. Pharmacology and Experimental Therapeutics, v. 63.

U. S. Public Health Service, 1946, Drinking water standards: U. S. Public 

Health Service Bepts., v. 61, no. 11.

D:31 PHOSPHORUS

Phosphorus is prevalent in nature both in the organic and inorganic 

form. Because of its role in animal and vegetable metabolism, phosphorus is a cyclic element, similar to nitrogen, in that the combined 

form is continually changing by decomposition and synthesis. As 

phosphate, it is present in several common minerals but the concentration in water is limited by the relative insolubility of phosphates 

of the alkaline earths. Both organic and inorganic phosphorus in 

water may result from leaching of soil and rocks, and from fertilizer, 

normal decomposition of plants and animals, sewage, and industrial 

effluents. Waters that have undergone phosphate treatment for removal of hardness can also add phosphorus to water resources.

In concentrations found in water, phosphorus is not reported to 

be toxic to man, animals, or fish. However, the element does stimulate 

the growth of algae, which may cause serious odor problems in public 

water supplies.

D:31a ORTHOPHOSPHATES

Orthophosphate is the most common ionized form of phosphorus in 

water and the only one derived from natural sources. Orthophosphates include the three ionization products of phosphoric acid, 

H-jPCV1, HPOr2 , and PO4~3, whose relative concentrations in water 

are a function of the pH. PO4~3 does not exist in solutions below a 

pH of about 9.4; at a pH of approx 6.8, HPO4^ and HaPCV1 are 

present in equal quantities, while at a pH of about 4.4 most of the 

orthophosphate is present as HaPOr1. See Reiman, Neuss, and 

Naiman (1942, p. 323), Kolthoff and Sandell (1952, p. 440-441), and 

sec. C : 8a for ionization of orthophosphate at different pHs.

The various ionization products are not differentiated by the method 

of analysis, nor are they chemically or physiologically significant to 

water uses. Hence, the combined orthophosphates are reported in 

terms of parts per million PO4. The computation of equivalents per 

million (epm), however, necessitates consideration of the ionization 

products ; instructions for estimating equivalents per million are given 

in sec. C : 8a.

D:31a-l PHOSPHOMOLYBDATE METHOD

As far as is known, the phosphomolybdate method is specific for 

the orthophosphate form of phosphorus. Weak tests are reported 

with pyrophosphate and polyphosphates, but these positive tests may 

well result from orthophosphate contamination of the material.

Principle of determination

Orthophosphate is converted to phosphomolybdate by acidified 

ammonium molybdate reagent.

H3PO4+12 (NH4 ) 2MoO4+21H+1-» (NH4 ) 3P

245
246 ANALYTICAL PROCEDURES D:31a-l

When phosphomolybdate is reduced with stannous chloride, a strong 

blue color is developed.

Barium, lead, mercury, and silver interfere by forming a precipitate. Silica gives a pale-blue color that is additive to the phosphate 

color, and a silica correction is made when necessary. The effect of 

silica is somewhat dependent on the reagents; therefore, appropriate 

silica correction should be determined with each batch of reagents. 

Nitrite interferes but can be oxidized to nitrate with hydrogen peroxide before the determination. Residual chlorine should be removed 

by boiling the sample. The reduction is not instantaneous, nor is the 

developed blue color stable. The full color develops in 6-10 min and 

fades gradually thereafter. After addition of stannous chloride the 

color can be measured at the exact predetermined time of maximum 

color development; or within about 30 min thereafter when the time 

effect is negligible. The amount of stannous chloride added affects 

the stability of the color. Using less reductant increases the stability 

but decreases the range of the test by incomplete reduction.

The stannous chloride solution probably causes more trouble than 

any other single factor. The reagent must be fresh. The salt hydrolyzes readily and is unstable even in the solid form. Only fresh salts 

with well defined crystals should be used. The acid solution can be 

kept in the reduced state for several days by adding a small amount 

of granulated tin or by using hydroxylamine hydrochloride.

With listed apparatus, results are accurate and reproducible to 

± 0.0005 mg.

Additional information on the principle of the determination is 

given by Ellis, Westfall, and Ellis (1946) and Zinzadze (1935, p. 227).

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength: 700 m^u 

Filter: Red

Cells : 40-mm optical depth 

Phototube: Red-sensitive 

Blank : Dilution water plus reagents 

Initial sensitivity setting: 1 

Slit width : 0.4 mm (approx)

The following absorbancies have been observed : 

mg PO4 Absorbancy

0. 010 0. 38

. 025 . 95

. 050 1. 91

Potassium phosphate, 1.00 ml=0.005 mg PO* 

Molybdate reagent 

Stannous chloride, 0.25 percent
D:31b-l PHOSPHORUS 247

Procedure

Samples for the determination of orthophosphate should be treated 

as directed in sec. A: 4d.

1. Pi pet a volume of sample containing less than 0.05 mg PO* (25.00 ml max) into 

a 50-ml beaker and adjust the volume to 25.0 ml.

2. Prepare a blank and sufficient standards, and adjust the volumes to 25.0 ml.

3. Add 1.0 ml molybdate reagent and mix.

4. After 5 and before 10 min add 1.0 ml 0.25 percent SnCl2.

5. After about 30-45 min, determine the absorbancies of test samples and standards against the blank, and when necessary, correct for water color as 

directed in sec. C : la-2, method 1.

Calculations

1. Determine mg PO4 from a plot of absorbancies of standards containing known 

amounts of constituent.

1 1 000 2. ppm PO4=-: rr-X :-* r-Xmg PO< in samples ^ density ml sample v

ppm Orthophosphate as P=ppm PCUX 0.326

Report concentrations of orthophosphates of < 1.0 ppm to 2 decimal 

places and of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Potassium phosphate, 1.00 ml=0.00f> mg PO4 : Dilute 50.00 ml potassium phosphate

(1.00 ml=0.100 mg PO*) to 1,000 ml. 

Potassium phosphate, 1.00 ml 0.100 mg P(V. Dissolve 0.1433 g KH2PO4 , dried

overnight over H2 SO4 , in water. Add 2-3 drops chloroform and dilute to

1,000 ml. 

Molybdate reagent: Mix equal volumes of 5 percent ammonium molybdate and

2QN H2SO4 . The reagent is unstable and should be prepared immediately

before use. 

Ammonium molybdate, 5 percent: Dissolve 10 g (NH4)eMo7O24 -4H2O in water and

dilute to approx 200 ml. 

Sulfuric acid, 2QN: Mix 112 ml cone H2 SC>4 (sp gr 1.84) with water and dilute

to approx 200 ml. 

Stannous chloride, 0.25 percent: Dilute 20 ml cone HC1 (sp gr 1.19) to approx

1 liter. Dissolve 2.5 g SnCl2 -2H2O and 10 g NH2OH-HO1 in the dilute acid.

1>:31I> PHOSPHORUS

The category "phosphorus" includes all forms of phosphorus, both 

inorganic and organic, that are in solution.

D:31b-l FHOSPHOMOLYBDATE METHOD

The phosphomolybdate method permits determination of phosphorus in all forms that decompose to yield orthophosphate under 

the drastic conditions of the test. Methane phosphonic acid is not 

determined, nor in all probability are other synthetic organic compounds of similar structure.
248 ANALYTICAL PROCEDURES D:31b-l

Principle of determination

Organic material is decomposed by acid digestion, and the phosphorus from this source as well as other phosphates such as metaphosphate, pyrophosphate and polyphosphates are oxidized to orthophosphate. The orthophosphate is then determined by development 

of the phosphomolybdate blue color as discussed in sec. D: 31a-l. 

Acid retards the development of the blue color; hence, the digested 

sample is neutralized before ammonium molybdate is added.

Additional information on the principle of the determination is 

given by Taylor (193T) and Hansen and Kobinson (1953, p. 55).

Apparatus and reagents

Steam bath

Graduated cylinders, 100-ml

Spectrophotometer, Beckman Model B :

Wavelength: 700 m^

Filter: Red

Cells: 40-mm optical depth

Phototube: Red-sensitive

Initial sensitivity setting: 1

Blank : Dilution water plus reagents

Slit width : 0.4 mm (approx)

The following absorbancies have been observed :

mg PO4 Abaorltancy 

0. 010 0. 38 

. 025 . 95 

. 050 1. 91 

Hydrochloric acid, cone (sp gr 1.19) 

Nitric acid, cone (sp gr 1.42) 

Sulf uric acid, 3.6N 

Phenolphthalein indicator solution 

Sodium hydroxide, IN 

Sulfuric acid, 0.052^

Potassium phosphate, 1.00 ml=0.005 mg PO< 

Molybdate reagent 

Stannous chloride, 0.25 percent

Procedure

1. Pipet 100.0-ml sample into a 250-ml Erlenmeyer flask.

2. Add 3.0 ml cone HC1.

3. Add 0.5 ml cone HNO3 .

4. Digest on steam bath for 2-3 hr.

5. Evaporate to approx 50 ml over a small flame.

6. Add 4.0 ml 3.62V H2SO4
D:31b-l PHOSPHORUS 249

7. Evaporate to approx 3 ml. The acid fumes at this point. Do not let portions 

of the bottom of the flask become dry.

8. Cool and dilute with approx 20 ml water.

9. Add a drop of phenolphthalein and titrate with IN NaOH to a pale-pink color.

10. Add 0.052V H2SO4 carefully until the pink color just disappears.

11. Quantitatively transfer the solution to a Nessler tube and dilute to original 

sample volume.

12. Pipet an aliquot containing less than 0.05 mg PO* (25.00 ml max) into 

a 50-ml beaker and adjust the volume to 25.0 ml.

13. Prepare a blank, sufficient standards and adjust the volumes to 25.0 ml.

14. Add 1.0 ml molybdate reagent and mix.

15. After 5 and before 10 min add 1.0 ml 0.25 percent SnCU

16. After about 30-45 min determine the absorbancies of test sample and 

standards against the blank.

Calculations

1. Determine quantity of PO4 in test sample from a plot of absorbancies of 

standards containing known amounts of constituent.

2. ppm Phosphorus as P04=-: rr- X . '.. - X mg P04 in sample density ml aliquot

ppm Phosphorus (P)=ppm P as PCXX0.326

Report concentrations of phosphorus as PO4 of <1.0 ppm to 2 

decimal places and of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Sulfuric acid, 3.62V: Mix 10 ml cone H2SO4 (sp gr 1.84) with water and dilute

to 100.0 ml. 

Phenolphthalein indicator solution: Dissolve 2.5 g phenolphthalein in approx

500 ml 50 percent ethyl alcohol. Neutralize with 0.022V NaOH. 

Sodium hydroxide, IN: Dissolve 40 g NaOH pellets in approx 1 liter of water. 

Sulfuric acid, 0.052V: Dilute 1.4 ml 3.62V H2 SO4 to 100 ml. 

Potassium phosphate, 1.00 ml=0.005 mg PO4 : Dilute 50.00 ml KH2PO« (1.00

ml=0.100 mg PO4 ) to 1,000 ml. 

Potassium phosphate, 1.00 ml=0.100 mg PO4 : Dissolve 0.1433 g KH2POi, dried

overnight over HaSO4, in water. Add 2-3 drops chloroform and dilute to

1,000 ml. 

Molybdate reagent: Mix equal volumes of 5 percent ammonium molybdate and

202V H2SO4. The reagent is unstable and should be prepared immediately

before use. 

Ammonium molybdate, 5 percent: Dissolve 10 g (NH4)eMo7O24-4H2O in water and

dilute to approx 200 ml. 

Sulfuric acid, 20IV: Mix 112 ml cone H2SO* (sp gr 1.84) with water and dilute to

approx 200 ml. 

Stannous chloride, 0.25 percent: Dilute 20 ml cone HC1 (sp gr 1.19) to approx

1 liter. Dissolve 2.5 g SnCV2H2O and 10 g NH2OH-HC1 in the dilute acid.
250 ANALYTICAL PROCEDURES Dl31b-l

REFERENCES

Ellis, M. M., Westfall, B. A., and Ellis, M. D., 1946, Determination of water 

quality: U. S. Fish and Wildlife Service Reserve Rept. 9.

Hansen, A. L., and Robinson, R. J., 1953, The determination of organic phosphorus in sea water with perchloric acid oxidations: Jour. Marine Research, Repr. 1.

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic 

analysis: New York, Macmillan Co., 3d ed.

Reiman, W., Neuss, J. D., and Naiman, B., 1942, Quantitative analysis: New 

York, McGraw-Hill Book Co., 2d ed.

Taylor, D. C., 1937, Modification of Deniges technique for the determination 

of phosphorus: Am. Water Works Assoc. Jour., v. 29, no. 12.

Zinzadze, C., 1935, Colorimetric methods for the determination of phosphorus: 

Indus, and Eng. Chemistry, Anal. Ed., v. 7.
D :32 POTASSIUM

Although potassium is relatively abundant in the earth's crust, 

the potassium content in natural waters is usually small. Potassium 

occurs in rocks in a form that is not easily brought into solution. 

Also, several geochemical processes tend to remove potassium selectively and return it to the solid phase. Most waters contain less 

than 20 ppm potassium, although several hundred parts per million 

are occasionally found.

Potassium is essential to animal nutrition, but a concentration of 

1,000-2,000 ppm in drinking water is regarded as the extreme limit 

permissible (Moore, 1950). Potassium in water causes foaming, as 

does sodium, but apparently it is not otherwise significant in industrial water supplies. Potassium stimulates plankton growth (Lackey 

and Sawyer, 1946, p. 573) and is reported to be somewhat more 

toxic to fish and shellfish than is calcium, magnesium, or sodium 

(Brandt, 1948, p. 9015).

D:32a-l FLAME-PHOTOMETRIC METHOD

The exact procedure used in the flame-photometric method is 

governed principally by the design and performance of the particular flame photometer used; hence, no specific directions can be 

given. See sees. C: Ib and C: 2e for a discussion of flame photometry and flame photometers. Many helpful suggestions are provided by the manufacturers of the individual instruments.

Report potassium concentrations of <10 ppm to 1 decimal place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm to 

3 significant figures only.

REFERENCES

Brandt, H. J., 1948, Intensified injurious effects on fish, especially the increased toxic effect produced by a combination of sewage poisons: Chem. 

Abs. 42.

Lackey, J. B., and Sawyer, C. N., 1946, Plankton productivity of certain southeastern Wisconsin lakes as related to fertilization: Sewage Works Jour., 

v. 17.

Moore, B. W., 1950, The desalting of saline waters, a review of the present 

status: Natl. Research Council Comm. on Sanitary Bng. and Environment, 

Kept, to Subcomm. on Water Supply.

251

D:33 SELENIUM

Selenium exists in many forms and is found in many localities. 

The element is generally more prevalent in the Great Plains and 

Great Basin than in other parts of the United States. Selenium is 

normally nonexistent or present only in insignificant quantities in 

natural waters but does occur occasionally under certain geologic 

environments. Highly significant is the fact that selenium from 

plants, sedimentary rocks, soils, and animals may find its way into 

water resources. Selenium concentrations in water in excess of 

0.5 ppm are extremely rare and limited primarily to seepage from 

seleniferous soils.

It is generally agreed that selenium is a cumulative poison to 

man and animals; it enters all of the body tissues and is present 

in all secretions and excretions. Because of its chemical similarity 

to sulfur, it is taken up nonpreferentially in tissue assimilations 

and is substituted for sulfur in the proteins. Investigations of 

selenium pathology to man have been somewhat complicated by 

difficulty in diagnosis, and it is understandable that recommendations for maximum limits in drinking and culinary waters are conflicting. The U.S. Public Health Service (1946) states that the 

selenium concentration shall not exceed 0.05 ppm in water on carriers 

subject to Federal quarantine regulations. Chronic selenium poisoning in livestock is known as "alkali disease," and the acute poisoning 

as "blind staggers" (Moxon, 1937). Water with 0.4 to 0.5 ppm is 

not believed to be toxic to livestock; such water may contribute to 

poisoning, but the selenium content of the feed is a more critical 

factor (California State Water Pollution Control Board, 1952, 1954, 

p. 350-351). It is generally agreed that when the selenium content 

of the air-dried diet exceeds about 5 ppm, symptoms of alkali 

disease occr -.

High concentrations of selenium are toxic to some plants, but 

most plants can absorb relatively large amounts of selenium from 

irrigation water or the soil without apparent injury to themselves.

Additional information on the occurrence and significance of 

selenium is given by Trelease and Beath (1949).

D:33a-l NORRIS-FAY TITRATION METHOD 

Principle of determination

For all practical purposes in water analysis, the method is specific 

for selenium but does not distinguish between the forms of the 

element. The sample is concentrated in the presence of sodium

253
254 ANALYTICAL PROCEDURES D:33a-l

peroxide to remove the interference of organic matter and to provide 

a workable volume of solution. Volumes throughout the procedure 

are more critical than for many other determinations in water 

analysis. Selenium is then distilled from sulfuric acid as the 

volatile bromide, precipitated as elemental selenium, collected, and 

oxidized to selenious acid. The selenious acid is titrated with thiosulfate and iodine using the Norris-Fay titration.

6Hrf+SeO,-*+4S,(V*-»S«SeO.-> +S«0.-1 +3H,0

Sulfuric acid is used in the distillation to raise the boiling point 

of the solution and insure quantitative distillation of selenium bromide. The elemental selenium precipitate is extremely fine, and the 

steam and ice-bath treatments are designed to coagulate the precipitate to a filterable size; nevertheless, attention must be given to 

the porosity of the filter. A fine asbestos mat is required with some 

fritted-glass filters. The Norris-Fay titration works best below 

25°C; hence, the solution is cooled in an ice bath before titrating.

With careful work, results are accurate and reproducible to 

±0.002 mg. The possibilities of manipulative errors are many, and 

it is recommended that duplicate samples, one containing a known 

added amount of selenium, be run.

Additional information on the principle of the determination is 

given by the Association of Official Agricultural Chemists (1945), 

Klein (1943), and Norris and Fay (1896).

Apparatus and reagents

Distillation apparatus, 250-ml, all glass, no offsets at joints, thermometer,

dipping adapter 

Steam bath 

Microburet, 5-ml 

Ice bath

Fritted-glass suction filter, fine, 22- X 100-mm 

Pipet-stirrers, 170-mm long (approx), with medicine-dropper bulbs 

Test tubes 

Suction flash 

Sodium peroxide 

Sulfuric acid, cone (sp gr 1.84) 

Hydrooromic acid, 40 percent or 48 percent 

Hydrobromic acid-bromine solution, cone 

Hydrobfomic acid-bromine solution, dilute 

Hydroxylamine hydrochloride, 10 percent 

Phenol, 5 percent 

Starch indicator, 0.5 percent 

Sodium thiosulfate, 0.001# 

Iodine, 0.001#

Selenium bromide, 1.00 ml=0.010 mg Se 

Sulfur dioxide (gas supplied in commercial cylinders is free of selenium)
D:33a-l SELENIUM 255

Procedure

Samples for the determination of selenium probably should be 

collected and treated in accordance with directions given in sec. 

A: 4d.

1. Pipet a volume of sample containing less than 0.05 mg Se (100.0 ml max) 

into the distillation flask.

2. Add sufficient Na2O2, usually approx 1 g, to oxidize any organic matter 

and to make the solution definitely alkaline.

3. Evaporate to about 25 ml on an open flame.

4. Cool and cautiously add 50 ml cone H2SO4 and mix.

5. To the cooled mixture, add 30 ml cone HBr-Br2 solution if the solution was 

prepared with 40 percent HBr, or 25 ml if 48 percent HBr was used.

6. Distill, slowly at first ,into 10-15 ml 48 percent or 40 percent HBr in a 

125-ml Brlenmeyer flask surrounded with cold water. Keep the tip of 

the adapter below the surface of the HBr until all free Br2 is distilled. 

If free Br2 does not distill first, add more cone HBr-Br2 solution to the 

distillation flask. Continue distillation until the temperature of the vapor 

reaches 200° C. At this point there should be no bromine coloration in 

the distillation flask or condenser.

7. Rinse the adapter into the receiving flask with water, keeping the volume 

of the distillate and washings below 75 ml.

8. Pass SO2 with swirling through the distillate until the Br2 color disappears 

and for an additional 30 sec to saturate the solution. (This step should 

be performed under the hood.)

9. Immediately add 1 ml 10 percent NH2OH HC1.

10. Place on a steam bath for 30 min.

11. Place in an ice bath for 30 min.

12. Collect the precipitated Se« in a fritted-glass filter overlaid with a flue 

asbestos mat.

13. Wash the flask and filter with 3-5 ml water.

14. Remove the last traces of SO2 from the flask with natural gas or compressed air.

15. Attach a test tube to the filter and dissolve the Se« in the flask and filter 

with 3 successive 1-ml portions of dilute HBr-Brs solution, sucking the 

filter dry after each addition. After each addition of HBr-Br2 solution, 

start the suction to pull the solution into the filter mat. Allow the solution to remain in the mat for about 1 min before continuing suction.

16. Rinse the flask and filter with three 1.5-ml portions of water using same 

precautions as with the HBr-Br* solution.

17. Detach the test tube and add 3 drops 5 percent phenol to neutralize the fir*

18. Using the pipet-stirrer, rinse down the sides of the tube with the solution.

19. Immerse in a hot-water bath for 5 min to complete the neutralization of 

Br2 and then in an ice bath for 5 min.

20. Add at least 50 percent excess O.OOltf Na2S2O3 and mix well.

21. Add 3 drops starch and sufficient I2 to give a strong blue color. If less than 

1 ml I2 is used, add sufficient O.OOltf Na2S2O3 to require 1 ml I2 . Record I2 

added to nearest 0.01 ml.

22. Back titrate with O.OOltf Na2 S2O3 to disappearance of blue, adding the last 

drops in 0.01-ml portions. Record total Na2S2Os used to nearest 0.01 ml.
256 ANALYTICAL PROCEDURES D:33a-l

Calculations

1. The thiosulfate equivalent of the iodine should be checked each time a 

selenium determination is made. Determine the Na2S2O3 equivalent as follows : 

To 4y2 ml water and 3 ml dilute HBr-Br2 solution in a test tube, add 3 drops 5 

percent phenol and heat in a hot-water bath for 5 min. Place in an ice bath for 

5 min. Add 3.00 ml I2 and titrate with 0.001N Na2 S2O3 , using starch indicator 

near the end point.

ml 0.001 N Na2S2O3 , ... _, . . . , . ,. . ,. ... 

 PJ = /= thiosulfate equivalent of the iodine solution

1.00 ml 0.001 N Na2S2O3 is equivalent to 0.0198 mg Se.

-(ml I2 X/)]X0.0198

Report selenium concentrations of < 1.0 ppm to 2 decimal places and 

of > 1.0 ppm to 2 significant figures only.

Preparation of reagents

Selenium bromide solution, 1.00ml=0.010 mg Se : Dilute 10.0 ml selenium bromide 

solution (1.00 ml=1.00 mg Se) to 1,000 ml with 0.05N HBr. Check the acidity by titration with 0.100N NaOH. The acidity should not be allowed to fall 

below 0.05N.

Selenium bromide solution, 1.00 ml=1.00 mg Se: Dissolve 0.4084 g H2SeO3) dried 

overnight in a desiccator, in 0.05N HBr and dilute to 250.0 ml with 0.05N HBr.

Hydrobromic acid, 0.05N: Dilute 10 ml 40 percent or 8.5 ml 48 percent HBr to 

approx 1 liter.

Hydrobromic acid-bromine solution, cone : Mix 10 ml Br2 with 990 ml 40 percent or 

48 percent HBr. Use caution when transferring liquid Br2 . Store in a 

tightly stoppered bottle. (This should be done under the hood.)

Hydrobromic acid-bromine solution, dilute : To 5 ml 40 percent or 48 percent HBr 

add 10 ml saturated bromine water and dilute to approx 100 ml.

Hydroxylamine hydrochloride, 10 percent: Dissolve 10 g NH2OH-HC1 in water 

and dilute to approx 100 ml.

Phenol, 5 percent: Dissolve 5 g phenol in water and dilute to approx 100 ml.

Iodine, O.OOltf: Dilute 50.00 ml 0.02N I2 to 1,000 ml. Before final dilution, 

add 20 g KI. Determine the thiosulfate equivalent of the iodine solution 

before each set of determinations.

Iodine, 0.02N: Dissolve 2.54 g I2 and 12 g KI in water and dilute to 950 ml before 

standardizing. Standardize against primary standard As2O3 as follows: 

Dissolve approx 0.04 g As2O3 weighed accurately to 0.0001 g in 10 ml IN 

NaOH. Add 15 ml IN H2SO4 and mix. Add 50 ml 4 percent NaHCOs. 

Titrate slowly with I2 solution, maintaining constant agitation until most 

of the I2 has been reacted (0.04 g As2O3 requires approx 40 ml 0.02N I2 ). Add 

starch solution for indicator and continue the titration until the initial pink 

coloration just passes to clear blue. Deduct from the volume of iodine 

consumed the amount required to produce the same color in a solution 

composed of the reagents added to 40 ml of freshly boiled and cooled water 

in which 5 g KI has been dissolved.

AT i«* rr g AS2 03 X20.220 Normality of I2=- % ml 12

Store in a glass-stoppered bottle protected from light.
D:33a-l SELENIUM 257

Sodium thiosulfate, 0.0012V, 1.00 mloO.0198 mg Se: Dilute 10.00 ml 0.100# 

Na2S2Os to 1,000 ml. Before adjusting to final volume, add 5 ml amyl alcohol 

and shake vigorously. Check the titer of the solution against the standard 

selenium solution (1.00 ml=0.010 mg Se). Add 3 ml dilute HBr-Br3 solution to 5.0 ml Se standard. Continue as shown under "Procedure," steps 

17-22, and "Calculations."

Sodium thiosulfate, O.lOOtf: Dissolve 25 g Na2S2Os -5H2O and 0.1 g Na8COs in 950 

ml carbon dioxide-free water. Allow the solution to stand for a day before 

standardizing. Standardize against 0.100JV KH(IOS ) 2 as follows: Dissolve 

3.250 g KH(lO3 )j in water and dilute to 1,000 ml. Dissolve 5 g KI in 100 

ml of water. Add 10 ml H2SO« (10 percent v/v) and 25.00 ml O.KHMV 

KH(IOS ) 2. Titrate with Na2SjOs using starch as indicator near the end 

point.

Starch indicator, 0.5 percent: Dissolve 0.5 g of soluble starch in 100 ml of water. 

Heating may hasten solution. Prepare fresh daily.

REFERENCES

Association of Official Agricultural Chemists, 1945, Official and tentative

methods of analysis of the Association of Official Agricultural Chemists:

Washington, D. C., 6th ed. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Klein, A. K., 1943, Determination of selenium in foods: Assoc. Official Agr.

Chemists Jour., v. 26. 

Moxon, A. L., 1937, Alkali disease or selenium poisoning: South Dakota Agr.

Expt. Sta. Bull. 311. 

Norris, J. P., and Fay, Henry, 1896, lodometric determination of selenious and

selenic acids: Am. Chem. Jour., v. 18. 

Trelease, S. F., and Beath, O. A., 1949, Selenium; its geological occurrence

and its biological effects in relation to botany, chemistry, agriculture,

nutrition, and medicine: New York [Published by the authors, Box 42,

Schermerhorn Hall, Columbia Univ.] 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public

Health Service Repts., v 61, no. 11.

D:34 SILICA

Silicon is the most abundant element in igneous rocks and some 

other types of deposits. Most silica in water is probably derived 

from the decomposition or metamorphism of silicate minerals rather 

than from solution of quartz, as quartz is one of the most resistant 

rock minerals to attack by water. Many waters contain less than 

10 ppm of silica; those that drain deposits high in silicate minerals, 

particularly feldspars, often contain up to 60 ppm; concentrations 

exceeding 100 ppm are not commonly found. The chemistry of 

silica in solution is not known with certainty. It is believed that 

most silica is present in a nonionized form, but ionized silicate (s) is 

undoubtedly present in some waters.

Silica is not physiologically significant to humans, livestock, or 

fishes, nor is it of importance in irrigation water.

Most industrial processes tolerate silica in the range normally 

found, but it is particularly undesirable in boiler feed water. A 

recommended upper limit for boilers operating at 400 psi or above 

is 1 ppm. Silica forms a hard coating on steam-turbine blades, and 

a limiting concentration of 0.1 ppm in the steam has been recommended (Minhoff, 1948, p. 438).

D:34a-l MOLYBDATE BLUE METHOD

The molybdate blue method is not recommended for normal waters 

whose silica content is more than 100 ppm.

Principle of determination

Silica in solution as silicic acid or silicate has the property of 

reacting with ammonium molybdate in an acid medium to form the 

yellow-colored silicomolybdate complex. The silicomolybdate complex is then reduced by sodium sulfite to form the molybdate blue 

color.

The possibility of having larger concentrations of so-called unreactive silica is greater in water containing high concentrations of 

silica than in water with a low silica content. One-hour digestion 

of a 50-ml sample with 0.2 g of silica-free sodium bicarbonate has 

been recommended as a means of making all the silica available 

for reaction with the molybdate reagent (APHA, 1955, p. 186).

Phosphate gives a similar molybdate complex under certain pH 

conditions. In the following determination, the conditions are such 

that the phosphate complex is not formed. There is also evidence 

that hydrogen sulfide and ferric and ferrous iron interfere with the 

determination. Hydrogen sulfide may be removed by boiling an 

acidified sample. The addition of disodium dihydrogen ethylenediamine tetraacetate eliminates the effect of high concentrations of

259
260 ANALYTICAL PROCEDURES D:34a-l

iron and also complexes calcium and prevents precipitation of 

calcium sulfite.

With the listed apparatus, results are accurate and reproducible to 

±0.005 mg below 0.1 mg and to ±0.02 mg near 0.5 mg.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 384-394) and Kahler (1941, 

p. 536).

Apparatus and reagents 

Spectrophotometer, Beckman Model B :

Wavelength : 700 m/i

Filter : Red

Cells : 10-mm or 40-mm optical depth

Phototube : Red-sensitive

Blank : Dilution water plus reagents

Initial sensitivity setting : 1

Slit width : 0.8 mm (approx)

The following absorbancies have been observed :

Cell depth 

mg SiOa (mm) Absorbancy

0. 05 40 0. 37

.10 40 . 73

.10 10 . 19

.20 10 . 36

.30 10 . 54

.50 10 .90 

Sodium silicate, 1.00 ml=0.050 mg SiO2 

Hydrochloric acid, 0.252V 

Ammonium molybdate, 5 percent 

fra2EDTA, 1 percent 

Sodium sulfite, 17 percent

Procedure

1. Pipet a volume of sample containing less than 0.5 mg SiO2 (10.00 ml max) 

into 50-ml beaker and adjust the volume to 10.0 ml.

2. Prepare a blank, sufficient standards, and adjust the volumes to 10.0 ml.

3. Add 5 ml 0.252V HC1.

4. Add 5 ml 5 percent ammonium molybdate.

5. Add 5 ml 1 percent Na2EDTA.

6. After 5 min have elapsed following the addition of the molybdate, add 10 ml 

17 percent Na2SO3.

7. Mix, and allow to stand approx 30 min. The color is stable for several hours 

after this time.

8. Determine the absorbancy of the test sample and standards against the blank.

Calculations

1. Determine mg SiO2 in test sample from a plot of absorbancies of standards 

containing known amounts of the constituent.

 ^ 2. ppm SiO2 =-j TT X r^ r-Xmg SiO2 in sampl density ml sample
D:34a-2 SILICA 261

Report silica concentrations of <10 ppm to 1 decimal place and of 

>10 ppm to 2 significant figures only.

Preparation of reagents

Sodium silicate, 1.00 ml=0.050 mg Sio2 : Stabilize the pentahydrate of approx 

3 g Na2SiO3 -5H2O by placing it in a CaCL> desiccator for 2-3 hr. Dissolve 1.765 

g in water and dilute to 1,000 ml. Dilute 100.0 ml of this intermediate solution to 1,000 ml. The concentration can be checked by the method described 

in sec. D ; 34a-2. Store in a plastic bottle.

Hydrochloric acid, 0.25#: Mix 22 ml cone HC1 (sp gr 1.19) with water and dilute 

to approx 1 liter.

Ammonium molybdate, 5 percent: Dissolve 52 g (NH4)gMo7O24 4H2O in water and 

dilute to approx 1 liter.

Na2EDTA, 1 percent: Dissolve 10 g Na2EDTA in water and dilute to approx 

1 liter.

Sodium sulfite, 17 percent: Dissolve 170 g Na2SO3 in water and dilute to approx 

1 liter.

D:34a-2 GRAVIMETRIC METHOD

The gravimetric procedure is recommended for the analysis of brine 

and for normal water whose silica content exceeds approx 100 ppm.

Principle of determination

Silica is precipitated by evaporation with strong acid, separated by 

filtration, ignited, and weighed. If the sample contains much nitrate 

or other oxidizing agents, the precipitation should not be carried out 

in platinumware because platinum is readily dissolved by oxidizing 

agents heated with hydrochloric acid. A porcelain dish is satisfactory, 

but glass should not be used. The anhydrous residue of the sample is 

first ignited to burn off the organic material, which might not dissolve 

in acid and thereby interfere with the determination. The redissolved, or resuspended sample is taken to dryness with hydrochloric 

acid to dehydrate colloidal silicic acid and convert it to a filterable 

mass. The digestion in dilute acid extracts the acid-soluble material. 

This extraction musl; be quantitative if the filtrate from the silica 

determination is used in other gravimetric determinations. The acid 

concentration in the digestion is important because the dilute acid 

decreases the amount of silica reverting to colloidal form, but concentrated acid increases the reversion.

The actual silica content of the acid-insoluble residue is determined 

by weighing the residue, volatilizing the silica as fluorosilicate, and 

subtracting the weight of the nonvolatile residue.

SiO2+6HF<=±H2 SiF8 +2H2O 

H2SiF<r»SiF4+2HF

The reaction proceeds to the right on heating. Sulf uric acid is added 

to assure that any iron or aluminum contamination in the precipitate
262 ANALYTICAL PROCEDURES D:34flr-2

will be weighed as the oxide in both weighings. Without the addition 

of sulfuric acid, iron and aluminum would be in the form of chloride 

in the first residue and fluoride in the second residue. In the presence 

of sulfuric acid, the contaminating metals are left in the residue as 

sulf ates, which decompose to oxides during ignition.

Fa, (SO*) r» Fe2O3 +3SO3

Inasmuch as the procedure is a single dehydration technique, the 

results are likely to be slightly low. For quantitative work it is necessary to recover the 1-5 percent of silica that may remain in solution. 

This may be accomplished readily by determining the silica in the filtrate colorimetrically (see sec. D: 34a-l).

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952, p. 384-394) and Hillebrand and 

Lundell (1929, p. 536-546).

Apparatus and reagents

Evaporating dish, porcelain or platinum

Platinum crucibles

Steam bath

Muffle furnace

Hotplate

Hydrochloric acid, 50 percent v/v

Hydrochloric acid, cone (sp gr 1.19)

Sulfuric acid, cone (sp gr 1.84)

Hydrofluoric acid, 48 percent

Procedure

1. Evaporate a volume of sample containing between 0.10 and 0.25 g dissolved 

material (1,000 ml max) in a platinum dish. The filtrate from the SiO2 

determination is used for the gravimetric determinations of calcium, 

barium, magnesium, and sulfate, and those requirements will have a 

bearing on the sample volume taken for the silica determination.

2. Ignite the residue over an open flame or in a radiator. See sec. D: 22a-l.

3. Cautiously wash down the sides of the dish with two 10-ml portions 50 percent HC1 followed by enough hot water to bring the solution level at least 

even with the highest sample level during the initial evaporation.

4. If the sample contains more than about 2 mg NO3 or other oxidizing agents, 

quantitatively transfer the solution or mixture to a porcelain dish; if 

higher concentrations of oxidizing reagents are not present, the following 

reactions can be carried out in the platinum dish.

5. Evaporate the acid solution to dryness on the steam bath, and leave the 

dish on the bath for at least 2 hr more.

6. Add 8-10 drops concentrated HC1 and wash down the sides of the dish with 

approx 50 ml water.

7. Cover the dish with a watch glass and extract the acid-soluble material 

by digesting on a steam bath for 1-2 hr.
D:34a-2 SILICA 263

8. Quantitatively collect the SiO2 and other insoluble material on Whatman 

No. 42 filter paper. Scrub out the dish with a rubber policeman to dislodge the SiO2, and wash the dish and the precipitate at least 3 times 

with hot water; retain the combined nitrate and washings for the gravimetric determination of calcium, barium, magnesium, and sulfate.

9. Transfer the residue and filter paper to a platinum crucible.

10. Moisten the paper and residue with 2 or 3 drops cone H2S(X

11. Ignite the paper slowly over a low oxidizing flame.

12. Transfer the crucible to the muffle furnace and ignite at 1,000° C for 0.5 hr.

13. Cool in a desiccator and weigh. Record the weight to the nearest 0.0001 g.

14. Add 2 drops cone H2 SO4 .

15. Add 1,4 in. 48 percent HF.

16. Under a hood, volatize the acid and SiF4 on a hotplate.

17. Reignite in the muffle furnace at 1,000° C for 1.5 hr.

18. Cool in a desiccator and reweigh. Record the weight to the nearest 0.0001 g.

Calculations

ppm SiOa=-i rr- X i 2 r~ Xloss of weight in mg *^ density ml sample

Report silica concentrations of <10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.

Preparation of reagents

Hydrochloric acid, 50 percent v/v: Mix 50 ml cone HC1 (sp gr 1.19) with water 

and dilute to approx 100 ml.

REFERENCES

American Public Health Association and others, 1955, Standard methods for

the examination of water, sewage, and industrial wastes: New York, Am.

Public Health Assoc., Inc., 10th ed. 

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis:

New York, John Wiley and Sons, Inc. 

Kahler, H. L., 1941, Determination of soluble silica in water photometric

method: Indus, and Eng. Chemistry, Anal. Ed., v. 13. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

Minhoff, C. E., 1948, Boiler water conditioning to prevent turbine deposits:

Petroleum Refiner, v. 27.

D:35 SODIUM

Sodium salts are very soluble, and most sodium leached from the 

soil or rocks tends to remain in solution. The use of sodium salts is 

common in industry, and industrial wastes may contain large quantities of the element.

Sodium is not particularly significant in drinking and culinary 

water except for those persons having an abnormal sodium metabolism ; the water supply in some areas contains sufficient sodium to be 

a factor in the planning of sodium-free diets (National Research 

Council, 1954). It has been established that more than 50 ppm of 

sodium plus potassium in boiler water causes foaming. For highpressure boiler feed water a limiting concentration of 2-3 ppm has 

been recommended. Water with a high ratio of sodium to calcium 

plus magnesium is deleterious to soil structure. Sodium tends to 

disperse the soil colloids with the resultant loss of good tilth and 

permeability (U.S. Salinity Laboratory Staff, 1954, p. 69-82). The 

sodium-adsorption-ratio (SAR) is a useful index of the sodium 

hazard of irrigation water.

SAR=

V

'epm Ca + Mg 

2

D:35a-l FIAME-PHOTOMETRIC METHOD

The exact procedure used in the flame-photometric method is 

governed principally by the design and performance of the particular flame photometer used; hence, no specific instructions can be 

given. See sees. C: Ib and C: 2e for a discussion of flame photometry and flame, photometers. Many helpful suggestions are provided by the manufacturers of the individual instruments.

Report sodium concentrations of <10 ppm to 1 decimal place, 

of between 10 and 999 ppm to whole numbers, and of >999 ppm 

to 3 significant figures only.

D:35a-2 ZINC TTRANYL ACETATE METHOD

The zinc uranyl acetate method is not as satisfactory for the 

determination of low concentrations of sodium as the flame-photometric procedure (see sec. D: 35a--l).

Principle of determination

Sodium ion reacts with a concentrated solution containing zinc, 

uranyl, and acetate ions to precipitate sodium uranyl zinc acetate 

hexahydrate.

Na+i + Zn+2 + 3UO2+2 + 9C2H3O2-' + 6H2O >

NaC2H3O2 Zn(C2H3O2). 3UO2 (C2HSO2), 6HaO

265
266 ANALYTICAL PROCEDURES D:35ar-2

The triple salt would ordinarily be considered a soluble salt. Solubility is 5.85 g per 100 ml, and precipitation is quantitative only 

in the presence of a very large excess of reagent. The sample 

volume must be reduced by evaporation to 1-2 ml, and the volume 

of reagent added must be 10 or 15 times greater.

The zinc uranyl acetate mixture possesses a high degree of 

specificity for sodium and is one of the more specific among all 

analytical precipitants. Lithium precipitates quantitatively. Potassium does not interfere. Iron, aluminum, and manganese do not 

interfere. Certain organic complexing ions such as oxalate and 

citrate interfere at high concentrations, but only rarely will their 

effect be found. Other interferences, including phosphate, are at 

a minimum.

In this determination, the major requirement for good results is 

good technique. Temperature control is rather important because 

the solubility of the triple salt varies with temperature. Ideally, the 

sample should retain the initial temperature of the reagent throughout the standing and filtration period. A rise above normal temperature causes low results and a fall in temperature causes high 

results. The effect is easily controlled without necessitating thermostatic control.

With careful technique, results are accurate and reproducible to 

±2 percent for water containing more than 20 ppm of sodium; in the 

lower concentration range the error may be appreciable.

Additional information on the principle of the determination is 

given by Kolthoff and Sandell (1952).

Apparatus and reagents

Steam bath

Crucible, fritted-glass, medium porosity. There may be a considerable variation in porosity between individual "medium" filters. A slow filter should 

be rejected.

Hydrochloric acid, cone (sp gr 1.19)

Zinc uranyl acetate solution, saturated with triple salt

Ethyl alcohol, 95 percent, saturated with triple salt

Acetone, anhydrous

Procedure

1. Pipet a volume of sample containing between 0.25 and 2.5 mg Na+1 into a 

beaker.

2. Add 1 or 2 drops cone HC1 and evaporate to dryness on a steam bath.

3. Dissolve the cooled residue in 1 ml water. Add more water if necessary, 

but note the volume used. A feathery residue of CaSO* may be tolerated 

because it usually dissolves in step 4 when the reagent is added.

4. Add zinc uranyl acetate reagent in the ratio of 15 ml reagent to 1 ml 

residue solution.

5. Allow to stand 1 hr at a fairly constant temperature. Stir occasionally.
D:35a-2 SODIUM 267

6. Filter off the precipitate under suction. Drain the beal.er and filter as completely as possible before proceeding to the next step.

7. Transfer the remaining precipitate into the filter crucible with 2-ml portions of saturated alcohol.

8. Wash the precipitate twice with 10-ml portions of saturated alcohol. The 

washed precipitate should be fine, crystalline, and bright yellow.

9. Wash the approx 10 ml acetone and draw air through fv. r 1 or 2 min.

10. Wipe the outside and inner bottom ring of the crucible with a moist towel 

if salts have crystallized there. Dry the crucible a ad precipitate in a 

desiccator for 1 hr.

11. Weigh the crucible and precipitate.

12. Transfer the crucible to the suction apparatus and dissolve the salt out with 

warm water. Wash with 10 ml acetone and dry as before and weigh.

Calculations

1. mg Triple salt= (wt crucible and salt) (wt washed crucible)

1 1 noo

2. ppm Na+'= .. X . ' , X 0.01495 X mg triple salt density ml sample

Report sodium concentrations of <999 ppm to whole numbers and 

of > 999 ppm to 3 significant figures only.

Preparation of reagents

Zinc uranyl acetate solution: Dilute 27 ml glacia? acetic acid with 1,000 ml 

dilution water. Dissolve in this solution 100 g UOi(CaH*Oi)i*2H*O, and 

300 g Zn (C*HiO2)i-2HiO. On cooling, a precipitate of fine, crystalline triple 

salt should form due to sodium impurities in the reagent. This salt should 

be allowed to remain in the bottle to maintain the sodium saturation. If 

a precipitate does not form, add a few milligrams of NaCl.

Ethyl alcohol, 95 percent, saturated: Saturate the alcohol with the triple salt 

and allow them to stand in contact at all times. Decant the saturated alcohol just before use.

REFERENCES

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed.

National Research Council, 1954, Sodium restricted diets: Pub. 325. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline

and alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60.

D :36 SOLIDS 

D:36a DISSOLVED SOLJDS

Theoretically, dissolved solids are the anhydrous residues of the 

dissolved substances in water. Note that dissolved solid is not a 

measure of the total weight of dissolved materials as they occur in 

solution. Neither are dissolved gases of volatile liquids a part of 

the dissolved solids. In reality, the term "dissolved solids" is defined 

by the method used in the determination. The methods used by 

the Geological Survey reflect experience with thousands of water 

samples of different total concentrations and chemical compositions, 

and the results obtained are believed to approach closer the theoretical definition than results by other methods.

All salts in solution affect the chemical and physical properties 

of the water and exert osmotic pressure. Water with several thousand parts per million of dissolved solids is generally not palatable, 

although those accustomed to highly mineralized water may complain that less concentrated water tastes flat. A change in the 

source of drinking water more often may cause gastric disturbances 

than the concentration of dissolved solids in the water itself. The 

U.S. Public Health Service (1946) recommends that the maximum 

concentration of dissolved solids not exceed 500 ppm in drinking 

and culinary water on carriers subject to Federal quarantine regulations but permits 1,000 ppm if no better water is available. Reports of livestock tolerances are extremely variable and range from 

3,000 ppm (Colorado Agricultural Experiment Station, 1943) to 

15,000 ppm (Heller, 1933). In Montana the following arbitrary 

limits are used to classify water for livestock: up to 2,500 ppm, 

good; 2,500-3,500 ppm, fair; 3,500-4,500 ppm, poor; and over 4,500 

ppm, unsatisfactory (I. Seghetti, 1951, written communication). 

The blood of fresh-water fish has an osmotic pressure equal to 

approximately 6 atmospheres, or about 7,000 ppm of sodium chloride. 

Although some fresh-water fish are adapted to live in more saline 

waters, for practical purposes any water that has an osmotic pressure 

greater than 6 atmospheres can be expected to be lethal.

Industrial tolerances for dissolved solids differ widely, but few 

industrial processes will permit more than 1,000 ppm. Salinity 

hazard is an important consideration for irrigation water. The 

water-uptake relations of plants are controlled by the osmoticpressure differential between soil solution and the plant solution. 

A plant cannot draw as much water from a concentrated soil solution as from a dilute soil solution. For most waters that could be

269
270 ANALYTICAL PROCEDURES D:36a-l

considered for irrigation, the following general relation is applicable :

Specific conductance X0.65±0.1=dissolved solids

The U.S. Salinity Laboratory Staff (1954) classifies the salinity 

hazard of irrigation waters, in terms of specific conductance, as 

follows: Less than 250 micromhos, low; 250-750 micromhos, medium; 

750-2,250 micromhos, high; and greater than 2,250 micromhos, very 

high. However, the satisfactory use of a particular water for irrigation depends on many factors other than water quality, such as soil 

characteristics, drainage, irrigation practices, and crops grown.

Two methods of determination are given below, but the results 

obtained are not strictly comparable. The method of determination 

should be included on the analytical statement.

D:36a-l RESIDTTE-ON-EVAPORATION METHOD

The residue-on-evaporation method is recommended for all waters 

that contain less than 1,000 ppm dissolved solids and is applicable 

to all waters regardless oiL concentration, provided that the residue 

layer in the evaporating dish is kept sufficiently thin.

Principle of determination

' A volume of sample that will yield less than 200 mg residue is 

evaporated just to dryness on a steam bath. The residue is dried at 

180°C for 1.0 hr, cooled in a desiccator, and immediately weighed. 

The weight of the residue is limited to 200 mg to insure subjection of all of the residue to the full effects of drying at 180°C. 

Voluminous residues will often seal over during the evaporation 

process and even entrap pockets of water that will not be completely 

driven off during the drying process. Massive residues also give up 

their water of crystallization more slowly than thin films of residue. 

The chemical composition of water has a marked effect on the 

dissolved-solids value obtained, but the percentage of error incurred 

for any given chemical type of water is independent of the total 

concentration if the residue film is kept thin.

Bicarbonate is converted to carbonate in the evaporation and 

drying process. The following other general observations have been 

reported (Howard, 1933, p. 4). The residues of carbonate-type 

water that contains considerable magnesium may often be less 

than would be expected. Eesidues of calcium and magnesium 

chloride water can be expected to lose some weight (up to 50 or 

100 ppm) during the drying process; however, such loss of weight 

is usually more than counteracted by water of crystallization tightly 

held by tne salts. Most of the water of crystallization is driven off 

from sulfates of sodium and magnesium when the residue is heated
D:36a-2 SOLIDS 271

at 180°C, but this temperature is insufficient to dehydrate calcium 

sulfate completely. Kesidues of water with a high nitrate content 

may lose as much as 30 ppm on heating.

Inasmuch as many of the salts in the residue are hygroscopic, it 

is imperative that an efficient desiccant be used. Alumina in which 

a moisture indicator is incorporated is recommended. Under no 

circumstances should the dried residues be allowed to stand for 

long periods of time before weighing.

With listed apparatus, results are generally reproducible to ±0.5- 

1.0 mg.

Apparatus

Platinum evaporating dishes, 75- to 125-ml capacity, weighing less than 50 g. 

Platinum is recommended for precise work because the change in weight 

of glass or porcelain dishes may introduce appreciable error into the 

determination.

Steam bath

Oven, 180°C, uniform temperature throughout

Procedure

1. Pipet a volume of sample containing 10-200 mg (500 ml max) into a tared 

platinum dish.

2. Evaporate the sample just to dryness on a steam bath.

3. Dry in the oven at 180°C for 1.0 hr.

4. Cool in desiccator and immediately weigh. Record the weight to the nearest 0.0001 g.

Calculations

ppm Dissolved solids = -3 rr-X r^ j-Xnig residue density ml sample

Report dissolved-solids concentrations of < 1,000 ppm to whole 

numbers and of > 1,000 ppm to 3 significant figures only.

D:36a-2 CALCULATION METHOD

Comparison of dissolved-solids concentration obtained by the 

residue-on-evaporation and calculation methods is discussed in sec. 

C:8b.

Principle of determination

The concentrations of all determined solid constituents are converted mathematically into the forms in which they would normally 

exist in an anhydrous residue. These quantities are then summarized.

The calculation method is applicable only to those analyses that 

are complete for all practical purposes. The chemist can never be 

certain of the completeness of the analysis, but for most alkaline 

waters other than brines the determination of silica, iron, calcium, 

magnesium, sodium, potassium, alkalinity, sulfate, chloride, and
272 ANALYTICAL PROCEDURES D:36a-2

nitrate is sufficient. The wide range of metals possible in acid 

water precludes assumptions as to the completeness of the analysis. 

If the water is grossly polluted it is usually necesary to determine 

the nitrogen components (see sec. D: 25), and perhaps estimate the 

carbonaceous material, before assuming ? relatively complete analysis. Concentration of carbonaceous material can be estimated by 

redissolving the residue on evaporation and treating it with several 

successive portions of hydrogen peroxide. The difference in weight 

between the nonoxidized residue and the oxidized residue is an indication of the carbonaceous solids. The estimated carbonaceous solids 

are not included in the calculated dissolved solids but are one 

measure of the differences to be expected between dissolved solids 

determined by residue on evaporation and by calculation.

The conversion of the analytical statement to the forms in which 

the constituents would normally exist in an anhydrous residue 

involve many variables that are not known. It is assumed that all 

bicarbonate in solution will exist as carbonate in the residue.

2HCO3 -1 -»CO3 -2

Therefore, the bicarbonate in solution is divided by 2.03 to give its 

equivalent weight as carbonate in the residue. If the bicarbonate 

alkalinity includes other ions (see sec. D: 2), the calculated dissolved 

solids will include these constituents twice. However, occurrence 

of such errors of significant magnitude is relatively rare. For 

some constituents, such as boron, arsenic, phosphate, and selenium, 

the ionic states in solution are not known, much less the forms that 

the elements will take in the anhydrous residue. Heavy metals may 

be present in solution either as cations or as colloidal hydroxides, 

and their forms in the anhydrous residues cannot be predicted 

either. Consequently, when summarizing the constituents, all material is arbitrarily assumed to be present in the theoretical anhydrous residue in the same form as reported in the analysis, with 

the exception of bicarbonate. Dissolved solids calculated by the 

Geological Survey include all determined material usually existing 

in solid form at normal temperature and pressure.

The accuracy and reproducibility of the results are dependent on 

the completeness of the analysis and on the validity of each reported 

constituent concentration.

Calculations

1. Convert reported bicarbonate to carbonate.

ppta co3 -^ppm

2.03

2. Add all determined dissolved-solid material reported in the analytical statement.
D:36b-l SOLIDS 273

Report dissolved-solids concentrations of < 1,000 ppm to whole 

numbers and of > 1,000 ppm to 3 significant figures only.

D:36b SUSPENDED SOLIDS

Suspended solids are those that can be separated from the sample 

by filtration. The determined value is fairly representative of the 

sample but does not accurately represent the suspended-sediment 

concentration of the stream; suspended-solids values should not be 

confused with sediment concentration, which is the more accurate 

measure of material in suspension.

Note that suspended solids are reported as milligrams per liter of 

mixture, whereas all other concentrations determined by methods in 

this manual are reported as weight per million weights of clear 

solution (see sec. B:5). Suspended-solids concentrations reported 

in relation to weight of clear solution would be meaningless; furthermore, parts per million based on the weight of mixture would be 

misleading in that the different bases are not differentiated. Milligrams per liter accurately describes what is measured, and the 

reported value is sufficient for uses of the measurement.

D:36b-l FILTRATION METHOD

The filtration method is similar in substance to part IIB, APHA 

(1955, p. 271) Standard Methods, and D 1069-54 T, method B-c, 

ASTM (1954, p. 323) Manual on Industrial Water.

Principle of determination

The precision and accuracy of the determination and the significance of the values obtained usually do not warrant the collection 

of a separate sample. If the suspended material is sufficiently 

settled to leave a clear supernatant solution for analysis, the suspended-solids determination can be postponed until test samples 

have been withdrawn for other determinations. The sample bottle 

is marked at the original waterline, and the volume of water removed 

for any chemical analysis is replaced by dilution water. This treatment is not recommended if a significant quantity of material is 

removed with the test samples.

The sample is mixed well and a suitable volume quickly poured 

off into a graduated cylinder. Withdrawal must be rapid to get a 

representative sample, and no attempt should be made to take an 

exact volume. The suspended solids are collected in a Gooch crucible and the insoluble material dried and weighed. A glass-fiber 

filter-paper circle is preferable to an asbestos mat because it withstands washing better.
274 ANALYTICAL PROCEDURES D:36b-l

Apparatus and reagents

Gooch crucible 

Glass-fiber filter-paper circle 

Suction filtration apparatus 

Oven, 110°C

Procedure

1. Mix the suspension well, and rapidly pour off a suitable volume into a graduated cylinder.

2. Record the volume.

3. Quantitatively collect the insoluble materials on the mat in a tared Gooch 

crucible.

4. Wash the insoluble material sparingly with dilution water.

5. Dry the crucible overnight at 110°C.

6. Cool in a desiccator and weigh.

Calculations

mg Suspended solids per liter= = L-r-r Xmg residue 

e ^ v ml mixture

REFERENCES

American Public Health Association and others, 1955, Standard methods for the 

examination of water, sewage, and industrial wastes: New York, Am. 

Public Health Assoc., Inc., 10th ed.

American Society for Testing Materials, 1954, Manual on industrial water: 

Spec. Tech. Pub. 148-A.

Colorado Agricultural Experiment Station, 1943, Mineral tolerances in livestock drinking water: 56th Ann. Kept.

Heller, V. G., 1933, The effect of saline and alkaline waters on domestic animals : Oklahoma Agr. and Mining Coll. Expt. Sta. Bull. 217.

Howard, C. S., 1933, Determination of total dissolved solids in water analysis: 

Indus, and Eng. Chemistry, Anal. Ed., v. 5.

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public 

Health Service Repts., v. 61, no. 11.

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline 

and alkali soils: U. S. Dept. Agriculture, Agriculture Handb. 60.
3>:37 SPECIFIC CONDUCTANCE

The specific conductance of an electrolyte is the reciprocal of 

specific resistance and is expressed in mhos. Specific resistance is 

the resistance, in ohms, of a column of solution 1 cm long and 1 sq 

cm in cross section. In most waters, the conductance is so low that 

micromho is used as the unit of expression. The specific conductance is a measure of the ability of the water to carry an electric 

current and is therefore an indication, within rather wide limits, 

of the ionic strength of the solution.

Many natural waters in contact only with granite, silicious sand, 

well-leached soil or other difficultly soluble material have a conductance of less than 50 micromhos. Specific conductance of most 

waters in the Eastern United States is less than 1,000 micromhos, 

but in the arid western parts of the country, a specific conductance 

of more than 1,000 micromhos is common. Some saline lakes and 

brines may even reach several hundred thousand micromhos.

D:37a-l WHEATSTONE BRIDGE METHOD

Principle of determination

Specific conductance is determined by using a Wheatstone bridge 

in which a variable resistance is adjusted so that it is equal to the 

resistance of the unknown solution between platinized electrodes of 

a standardized conductivity cell. The fraction of the current carried 

through the solution is a function of the relative velocities with 

which the different ions move; ions differ widely in their velocity 

and in their effect on conductance (see sec. C: 8c). Kohlrausch has 

shown that at infinite dilution, ions behave independently and that 

in this range the conductance-concentration relations of electrolytes 

of identical percentage composition is practically linear. However, 

as the concentration of the electrolyte increases, interionic attraction 

decreases the conductance per unit mass of electrolyte. This change 

in electric conductance of strong electrolytes with concentration is 

explained by the Debye-Huckel theory of interionic attraction.

A = A 0 a-\/c A. 0 

where

A = equivalent conductance, 

A 0 =equivalent conductance at infinite dilution,

c= concentration, and 

a and 6=theoretical constants.

For practical purposes the relation between specific conductance 

and concentration is linear below about 10,000 micromhos.

The temperature of the electrolyte affects the ionic velocities and 

consequently the conductance. Conductance increases about 2 per275
276 ANALYTICAL PROCEDURES D:37a-l

cent per degree centigrade, which is about the same as the temperature coefficient of viscosity of water. Specific conductance at 25 °C 

is reported.

In determining the specific conductance, it would be troublesome 

to prepare a cell having electrodes exactly 1 sq cm in area and 

exactly 1 cm apart. However, it is not necessary to do this because it 

is possible to determine a factor called the cell constant (C). The 

cell constant is determined experimentally with a standard solution 

of known conductance. A 0.00702^ potassium chloride solution has 

a specific conductance of 0.001000 mhos at 25°C. The relation 

between resistance (R), cell constant (6"), and specific conductance 

(K] is shown in the following equation, where K is known and R 

is determined.

RK=C

Thus, if the resistance of the cell when filled with 0.00702# KC1 is, 

for example, 350 ohms, the cell constant would be 0.35 for the conductivity cell used. If the conductivity cell having a cell constant 

of 0.35 is filled with the sample of water at 25°C and the observed 

resistance is 865 ohms, the specific conductance of the sample could 

be derived from the cell-constant equation.

R~c_KK

or, by substituting values from the example,

n ^

-^=0.000405 mhos at 25°C. obo

Unless a constant-temperature room or bath is available, adjustment of temperature (T) to exactly 25° C is difficult. For most work, 

specific conductance is computed from the following equation:

,Micromhos^ - T . , R of 0.00702AT KC1 at T of sample measurementX 1,000 j . ^- - 

R of the sample

New conductivity cells should be cleaned and the electrodes 

platinized before using. Subsequently they should be cleaned and 

replatinized whenever the readings become erratic or indistinct or 

inspection shows that any platinum black has flaked off. One 

platinization will usually suffice for a period of several months. 

To platinize the electrodes, clean them in chromic acid solution 

and rinse thoroughly in several changes of water. Place the electrodes in a solution of chloroplatinic acid and lead acetate (dissolve 

3 g H2PtCl6 in 10 ml of water to which 20 mg of Pb(C2H3O2 ) 2 

is added; commercial platinizing solutions are also available). 

Connect the electrodes with 2 dry cells (l 1/^ volts each) in parallel
D:37a-l SPECIFIC CONDUCTANCE 277

and reverse the direction of the current once a minute for 6 minutes, 

or until the shiny platinum surface is covered. Avoid deposition 

of amorphous platinum on the electrodes. Repeat the electrolytic 

process using 10 percent sulfuric acid to remove chlorine. When 

not in use the cell should be kept immersed in distilled water.

The accuracy and reproducibility of results obtainable are dependent to a large extent on the type of bridge used but may approach 

±2-3 percent with this equipment. Close attention to temperature 

is essential for reliable work.

Additional information on the principle of the determination is 

given by Getman and Daniels (1946, p. 390-396) and Scofield 

(1932).

Apparatus and reagents

Wheatstone bridge (see sec. C : 2a for discussion of conductivity meters) 

Conductivity cell (see sec. C : 2a) 

Thermometer, 0°-50°C, graduated in 0.1° 

Potassium chloride, 0.00702A7

Procedure

The manufacturer's instructions for operation of the bridge should 

be followed explicitly. A constant-temperature room or bath at 

25 °C simplifies temperature consideration. Where such facilities 

are not available, the sample should be brought to approximately 

room temperature before the determination. However, a group of 

samples standing in the laboratory are seldom at exactly the same 

temperature because of the influence of drafts, sunlight, radiators, 

ovens, open flames, etc. Whether the conductance determination is 

made in the field or laboratory, the temperature of each sample 

should be determined at the time of measurement.

1. Prepare a graph of resistance of 0.00702^ KC1 throughout the operating 

temperature range.

2. Rinse the cell with sample.

3. Measure the resistance of the sample and record the temperature at the time 

of measurement. Record temperature to the nearest 0.1 °C.

4. Determine the resistance of 0.00702# KC1 at the temperature at which the 

sample resistance was measured from the graph prepared in step 1.

Calculations

c . . , , . , . R of 0.00702AT KC1 , nnrt Specific conductance (micromnos) = ~ ;; ; X 1,000 R of sample

where R = resistance in ohms.

Record conductivities of < 1,000 micromhos to whole numbers and 

of > 1,000 to 3 significant figures only.

Preparation of reagents

Potassium chloride, 0.00702A7 : Dry approx 1 g KC1 in an oven at 180°C for 1 hr 

and cool in a desiccator. Dissolve 0.5232 g in water and dilute to 1,000 ml.
278 ANALYTICAL PROCEDURES D:37a-l

REFERENCES

Getman, F. H., and Daniels, Farrington, 1946, Outlines of physical chemistry:

New York, John Wiley and sons, Inc., 7th ed. 

Scofield, C. S., 1932, Measuring the salinity of irrigation waters and of soil

solutions with the Wheatstone bridge: U.S. Dept. Agriculture Circ. 232.
D:38 SULFATE

Sulfate is dissolved from most sedimentary rocks. Large quantities may be derived from beds of gypsum, sodium sulfate deposits, 

and some types of shale. Water from mines may be high in sulf ate 

as a result of oxidation of pyrite. Organic material containing 

sulfur adds sulfate to the water as a phase of the sulfur cycle. 

Sulfate is discharged in many industrial waste products. In natural 

waters, concentrations range from a few parts per million to several 

thousand parts per million.

Salts of sulfate are saline cathartics and a quantity equal to that 

in 1 liter of water containing 1,000-2,000 ppm sulfate constitutes 

an average dose. The U.S. Public Health Service (1946) recommends that the sulfate concentration not exceed 250 ppm in drinking 

and culinary water on carriers subject to Federal quarantine regulations. There is some report of a detrimental effect of sulfate to 

livestock, but it also seems to offset in part the toxicity of selenium 

toward cattle (Hurd-Karrer, 1934, 1935, p. 289).

The sulfate content of water is not very critical in many industrial 

processes. The significance is also somewhat dependent on the 

character of the cations (California State Water Pollution Control 

Board, 1952, p. 377-378). It is less toxic than chloride to crops.

D:38a-l VISUAL THORIN METHOD

Principle of determination

Thorin reacts with barium to give a red color. The intensity of 

the color is dependent on pH, indicator concentration, and nature 

of the solvent., The color is more intense in organic media than in 

water solutions. The color reaction can be utilized to titrate sulfate 

directly with barium chloride by adding a large volume of organic 

solvent to the sample and titrating in this mixed medium. The 

color development is also highly dependent on the nature of the 

organic solvent; dioxane is superior to either ethyl or methyl alcohol 

when the end point is detected visually. The initial color of thorin 

in the dioxane-water medium is yellow, and the change in the end 

point is to pink. The color change is enhanced when the solution 

is viewed through a pale-blue filter such as a didymium glass filter 

or glass-blowers' goggles.

The optimum pH for the titration is about 2.5, which permits the 

use of the method for the analysis of many water samples without 

danger of precipitation of some salts in the organic media. Thorin 

reacts with many metals, including calcium; therefore, it is necessary to remove all metal ions by cation exchange prior to titration. 

Phosphate interferes somewhat by coprecipitation; with 100 ppm of

279
280 ANALYTICAL PROCEDURES D:38ar-l

sulfate, 10 and 20 ppm of phosphate give a positive error of about 

2 and 3 percent, respectively. Color may interfere with the detection 

of the end point, and the method described in sec. D: 38a-2 is recommended for highly colored water.

Two strengths of barium chloride and thorin are provided, but 

there is a rather wide overlap in application. The dilute titrant and 

indicator are recommended if the sulfate content of the sample is less 

than 5 mg. The concentrated titrant and indicator give better results 

in the higher concentration ranges.

Results are accurate and reproducible to ±0.02 mg in the 0-5 mg 

range and ± 1-2 percent for higher concentrations.

Apparatus and reagents

Ion-exchange columns, charged with Amberlite IR-120 and operating on the 

hydrogen cycle. The column should be at least 10 in. long to assure complete exchange in moderately concentrated water. Highly mineralized 

water and brine require a longer column. Experience has shown that a 

strong acid solution (HC1 30 percent v/v) is required to get uniformly 

satisfactory regeneration. The frequency of regeneration depends on the 

mineral content of the samples; with average water, regeneration after 3 

or 4 passes is sufficient.

Pale-blue filter: A didymium filter or glass-blowers' goggles have been used 

satisfactorily.

Titration assembly with a white porcelain base and fluorescent light.

Buret, 25-ml

Sodium hydroxide, O.OoN

Thorin indicator, 0.04 percent

Thorin indicator, 0.1 percent

1,4 dioxane: The grade labeled "purified" is usually satisfactory, but distillation 

from glass may be required if the blank titration is excessive.

Barium chloride, 1.00 mloO.500 mg SO4~2

Barium chloride, 1.00 mlol.OO mg SOr2

Procedure

1. Rinse the ion-exchange column with 20-30 ml of sample and discard the rinse 

water. (This portion can be checked for calcium if desired.)

2. Pass a sufficient volume of sample through the exchanger to provide 25-30 

ml of effluent for the determination.

3. Pipet a volume of effluent containing less than 25 mg SO4~2 (25.00 ml max) 

into a 1.50-ml beaker and adjust the volume to approx 25 ml. The clissolvedsolids content should not exceed 125 mg for the high-range titration or 50 

mg for the low range.

4. Adjust the pH to between 2.2 and 5 with 0.05N NaOH if necessary.

5. Add 50 ml dioxane.

6. Add 1.0 ml thorin indicator.

7. Titrate with BaCL to the point where the color changes suddenly from yellow 

to orange when viewed through the filter. On titrating sulfate concentrations 

of less than 10 ppm, the end point will appear and then fade on stirring. At 

the true end point the color persists for several minutes.
D:38ar-2 SULFATE 281

Calculations

When BaCl2 , 1.00 mloO.50 mg SO4"2, and 0.04 percent thorin indicator are used: 

The titration is not strictly linear below 5 mg SOr2, and the SO^ concentration is determined by reading from a graph prepared by titrating standards. The following titrant volumes have been required: 

mg SO4-2 ml

0. 00 0. 05

.25 .55

1. 25 2. 50

2. 50 4. 95

5. 00 9. 90

ppm S04 -2 =, * X I'000 X mg SOr2 density ml sample

When BaOlz, 1.00 mlol.OO mg SOr2, and 0.1 percent thorin indicator are used:

ppm S0r2 = * X 1 1 ' 000 . X ml BaCl2 density ml sample

Report sulfate concentrations of <10 ppm to 1 decimal place, between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 significant figures only.

Preparation of reagents

Sodium hydroxide, 0.05N: Dissolve 2.0 g NaOH in water and dilute to approx 

1 liter.

Thorin indicator, 0.04 percent: Dissolve 0.20 g of thorin in water and dilute to 

500 ml.

Thorin indicator, 0.1 percent: Dissolve 0.5 g of thorin in water and dilute to 

500ml.

Barium chloride, 1.00 mlol.OO mg SOr2 : Dissolve 2.130 g anhydrous BaCU, dried 

overnight in an oven at 180°C, in water and dilute to 1,000 ml. This concentration is not strictly stoichiometric but has been determined empirically. 

Check the titer by titrating standard SOr2 solutions as directed in the 

procedure.

Barium chloride, 1.00 mloO.50 mg SOr2 : Dissolve 1.085 g anhydrous BaCl2 , dried 

overnight in an oven at 180°C, in water, and dilute to 1,000 ml.

Sodium sulfate, 1.00 ml=1.00 mg SOr2 : Dissolve 1.4787 g Na2 SO4, dried for 2 hr 

at 180°C, in water, and dilute to 1,000 ml.

D:38a-2 SPECTROPHOTOMETRIC THORIN METHOD

The spectrophotometric thorin method is useful for waters whose 

sulfate content does not exceed 200 ppm and for waters with high 

color that interferes with the visual detection of the end point in the 

visual thorin method (see sec. D: 38a-l).

Principle of determination

The chemistry of the determination is similar to that in section 

D:38a-l, with the exception that the preferable titration medium 

consists of 80 percent ethyl alcohol instead of 66 percent dioxane. The
282 ANALYTICAL PROCEDURES D:383r-2

pH is adjusted to and maintained at 5 by a sodium acetate buffer. 

The end point of the titration is detected instrumentally.

With listed apparatus, results are accurate and reproducible to 

± 0.005 mg, which is comparable in terms of parts per million to that 

for the visual thorin method (see sec. D: 38a-l).

Apparatus and reagents

Ion-exchange column (see sec. D : 38a-l)

Absorption cells, 50-mm. The cement in some cells reacts with thorin to give a 

red color. Cells should be tested for thorin reaction, and those that give 

a red color in 10-15 min should be rejected. 

Spectrophotometric-titration assembly (see sec. C : 2h). 

Spectrophotometer, Beckman Model B :

Wavelength: 520 m/j

Phototube: Blue-sensitive

Initial sensitivity setting : 1

End point: At 0.20 absorbancy 

Buret, 10-ml

Solvent-indicator solution 

Barium chloride, 1.00mlo0.20 mg SOr2

Procedure

1. Rinse the ion-exchange columns with 20-30 ml of sample and discard the 

rinse (this portion can be checked for the presence of calcium if desired).

2. Pass sufficient sample through the exchanger to provide a 10-ml effluent for 

the determination.

3. Pipet a volume of sample containing less than 1 mg SOr2 and 10 mg dissolved 

solids (10.00 ml max) into a 50-mm absorption cell and adjust the volume 

to 10.0 ml.

4. Add 40 ml solvent-indicator solution.

5. Start the stirrer and set the absorbancy to 0.100.

6. Titrate with BaCl2 , (1.00 mloO.20 mg SOr2 ) to an absorbancy of 0.20, which 

is stable for 30 sec.

7. Determine a blank correction by titrating dilution water. The blank is 

constant throughout the concentration range of the method. A blank of 

0.05 ml has been used.

Calculations

ppm SO4 -2 = 1 .. X Il0()0 x0.2X(ml titrant-ml blank) density ml sample

Report sulfate concentrations of <10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of > 999 to 3 significant 

figures only.

Preparation of reagents

Solvent-indicator solution: Dissolve 0.025 g thorin (the Hach Chemical Co.

product has been used satisfactorily), 0.5 g anhydrous sodium acetate in

10 ml of water. Add the solution to 1,000 ml of 95 percent ethyl alcohol.

Add 12 ml glacial HC2H3O2 (sp gr 1.049) and mix. 

Barium chloride, 1.00 mloO.20 mg SOr2 : Dissolve 0.434 g anhydrous BaCU,

dried for 1 hr at 180°C, in water, and dilute to 1,000 ml.
D:38a^3 SULFATE 283

D:38a~3 GRAVIMETRIC METHOD

The gravimetric method is similar in substance to part IA, APHA 

(1955) Standard Methods; D 516-49, ASTM (1954) Manual on Industrial Water; and method 83, U.S. Salinity Laboratory Staff (1954) 

Handbook 60.

Principle of determination

Sulfate is precipitated as barium sulfate, and the precipitate is 

ignited and weighed. Precipitation is carried out in acid media to 

prevent precipitation of barium carbonate, hydroxide, and phosphate, which are insoluble in alkaline media. Hydrochloric acid 

has the effect of promoting the formation of coarse crystals. The 

acid concentration is not important as long as it does not exceed 

0.05./V. The precepitation should be made at boiling temperature 

because the tendency toward supersaturation decreases at the higher 

temperatures. According to Kolthoff and Sandell (1952, p. 322-336) :

Barium sulfate provides a notorious example of a precipitate subject to contamination through the action of coprecipitation phenomena. Since the determination of sulfate is of great practical importance, a great deal of attention 

has been paid to the problem of obtaining a pure precipitate of barium sulfate, 

and hundreds of papers on the subject are to be found in the literature. In spite 

of the immense amount of work that has been done, the determination of sulfate 

remains one of the less accurate determinations of quantitative analysis.

The analyst is referred to authoritative texts, such as Kolthoff and 

Sandell (1952) and Hillebrand and Lundell (1929), for detailed 

discussions of potential errors involved. In water analysis the 

effect of coprecipitation is largely dependent on the character and 

concentration of the other solutes. Small quantities of ferrous iron, 

magnesium, zinc, and aluminum may be present. Calcium, ferric 

iron, and nitrate should not be present in appreciable amounts. 

Even though some of these constituents may be coprecipitated with 

the barium sulfate, prolonged digestion of the precipitate in the 

mother liquor at elevated temperatures makes for a precipitate 

approaching theoretical composition. Silica may also be precipitated from the hot acid solution. The interference of silica is eliminated if the sulfate is determined on the filtrate from the gravimetric 

silica determination (see sec. D: 34a-2). An alternative method that 

minimizes silica interference involves treatment of the ignited residue (procedure 9) in a platinum crucible with hydrofluoric acid and 

a few drops of concentrated sulfuric acid, followed by reignition. 

Best results are obtained if the weight of the precipitate is between 

0.025 and 0.150 g, about 10-60 mg of sulfate. With careful work, 

results are reproducible to ±0.2 mg; the accuracy is dependent on 

the purity of the precipitate.
284 ANALYTICAL PROCEDURES D:38a-4

Apparatus and reagents

Steam bath or hot plate

Porcelain crucibles

Muffle furnace

Methyl orange indicator solution

Hydrochloric acid, cone (sp gr 1.19)

Barium chloride, 10 percent

Filter-paper pulp

Procedure

1. Dilute or concentrate nitrate from the gravimetric silica determination 

(see sec. D : 34a-2, procedure 8) to a convenient volume. Pipet an aliquot containing between 10 and 60 mg S(V 2 into a 400-ml beaker and adjust the 

volume to approx 200 ml.

2. Acidify to methyl orange with cone HC1 and add a few drops in excess.

3. Heat the acidified aliquot to boiling and while hot add slowly dropwise, and 

with constant stirring, 10 ml hot 10 per cent BaCl2.

4. Cover the beaker with a watchglass and digest on the steam bath or hotplate 

(do not boil) for 2-3 hr, or preferably overnight.

5. Mix in a small quantity of filter-paper pulp and let stand for 1 hr on steam 

bath or hotplate.

6. Filter hot through Whatman No. 42 filter paper.

7. Quantitatively transfer the precipitate to the filter paper and wash thoroughly, scrubbing out the beaker several times with a rubber policeman and 

rinsing with hot water several times. The filtrate from the last washing 

should contain no chloride when tested with AgNO3 .

8. Slowly ignite the precipitate in a tared crucible over a low oxidizing flame 

until the filter paper is reduced to a white ash.

9. Transfer the crucible to the muffle furnace and ignite at approx. 800°C for 

1.0 hr.

10. Cool in a desiccator and weigh.

Calculations

ppm S04~2= * X I'000 X mg BaSO4 X0.4115 density ml sample

Report sulfate concentrations of <10 ppm to 1 decimal place, of 

between 10 and 999 ppm to whole numbers, and of >999 ppm to 3 

significant figures only.

Preparation of reagents

Methyl orange indicator solution: Dissolve 0.05 g methyl orange in water and

dilute to approx 100 ml. 

Barium chloride, 10 percent: Dissolve 100 g BaCl2 in water and dilute to approx

1 liter.

D:38a-4 TURBIDIMETRIC METHOD

The turbidimetric test is limited to rough approximations of sulfate. 

Its maximum utility is as a laboratory aid in the selection of proper 

sample size for the determinations described in sees. D: 38a-l, -2, -3, 

and as a field test for the approximation of sulfate when accuracy 

requirements are low.
D:38ar-4 SULFATE 285

Principle of determination

Sulf ate is precipitated as barium sulf ate, and the resultant turbidity 

is compared with that produced in standards of known concentration.

Apparatus and reagents

Test tubes

Sodium sulfate, 1.0 ml=0.10 mg SO4

Barium chloride, acidified, 10 percent

Procedure

1. Pipet a volume of sample containing less than 0.25 mg SCX"2 (10.0 ml max) 

into a test tube and adjust volume to 10.0 ml.

2. Prepare sufficient standards in the 0.00- to 0.25-mg range and adjust the 

volume to 10.0 ml.

3. Add 1.0 ml 10 percent BaCl2.

4. Shake vigorously and allow suspensions to stand 10 min.

5. Compare the turbidity of the sample with that of the standards. Comparison 

is best made by looking down through the tubes against a black background.

Calculations

i nnn

ppm SO4-2= , ' . X mg SOr2 in standard ^ ml sample

Preparation of reagents

Sodium sulfate, 1.0 ml=0.10 mg SCv*: Dissolve 0.148 g Na2SO4 in water and dilute

to 1,000 ml. 

Barium chloride, acidified, 10 percent: Dissolve 10 g BaCl2 in water, add 5 ml cone

HC1 (sp gr 1.19), and dilute to approx 100 ml.

REFERENCES

American Public Health Association and others, 1955, Standard methods for the

examination of water, sewage, and industrial wastes: New York, Am. Public

Health Assoc., Inc., 10th ed. 

American Society for Testing Materials, 1954, Manual on industrial water:

Spec. Tech. Pub. 148-A. 

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Hillebrand, W. F., and Lundell, G. E. F., 1929, Applied inorganic analysis: New

York, John Wiley and Sons, Inc. 

Hurd-Karrer, A., 1934, 1935, Selenium absorption by plants and their resulting

toxicity to animals: Smithsonian Inst. Ann. Kept. 

Kolthoff, I. M., and Laitinen, H. A., 1941, pH and electro titrations: New York,

John Wiley and Sons, Inc. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts, v. 61, no. 11. 

U.S. Salinity Laboratory Staff, 1954, Diagnosis and improvement of saline and

alkaki soils: U.S. Dept. Agriculture, Agriculture Handb. 60.

D:39 SULFIDES

Sulfide occurs in water as a result of bacterial and chemical 

processes. It usually is present as hydrogen sulfide. Variable 

amounts may be found in waters receiving sewage, and (or) industrial wastes, such as from tanneries, paper mills, chemical plants, 

and gas manufacturing work (California State Water Pollution 

Control Board, 1952, p. 271-272).

Waters containing sulfides, especially hydrogen sulfide, may be 

considered undesirable because of their odor. The U.S. Public 

Health Service (1946) states that water on carriers subject to 

Federal quarantine regulations shall have no objectionable taste 

or odor. The toxicity to aquatic organisms differs signifies rtly with 

the species.

D:39a-l IODOMETRIC METHOD

This iodometric method does not differentiate the forms in which 

the sulfide exists in solution. The procedure is recommended for 

use on water containing more than about 0.5 ppm of sulfide.

Principle of determination

Sulfide is reacted with an excess of iodine in acid solution, and 

the remaining iodine is then determined by titration with sodium 

thiosulfate using starch as an indicator (Kolthoff and Sandell, 1952, 

p. 585-605).

H+ S+2I-i 

I2 + 2S203-2

A blank is treated with exactly the same amount of iodine and 

titrated with sodium thiosulfate. The sulfide concentration is calculated from the difference between the volume of thiosulfate 

required for the blank and the volume used for the sample.

Reducing substances such as sulfites and heavy-metal ions will use 

up iodine and therefore contribute positive errors. Oxygen and 

other oxidants may react with hydriodic acid to liberate iodine and 

thus contribute negative errors.

When the sulfide concentration is below 0.05 mg, the method is 

accurate to only ±50 percent. Above 0.05 mg of sulfide the accuracy improves as the sulfide concentration increases, and, at 5 ppm 

of sulfide, the results are reproducible to ±0.05 mg.

Apparatus and reagents

Buret, 10-ml. 

Iodine solution, O.OKW 

Hydrochloric acid, cone (sp gr 1.19) 

Sodium thiosulfate solution, O.OKW 

Starch indicator, stable

287
288 ANALYTICAL PROCEDURES D:39a-l

Procedure

Samples for the determination of sulfide should be collected with a 

minimum of aeration and agitation as directed in sec. A: 3c and 

treated in accordance with directions given in sec. A: 4d.

1. Vigorously shake the sample to mix the contents thoroughly. Immediately 

pipet a volume of sample with ZnS in suspension containing less than 1.5 mg 

S (100.0 ml max) into a 250-ml Erlenmeyer flask and adjust the volume to 

approx 100 ml.

2. Prepare a blank of approx 100 ml water and carry it through the procedure 

with the sample.

3. Add 10.00 ml O.OIOA7 L. 

i. Without delay add 10 ml cone HC1.

5. Immediately titrate the excess L with O.OIOA7 Na2 S2O3 , adding 2-3 ml starch 

indicator as the end point is approached.

Calculations

c, , TT c, 1 1.000 n ,~ni /ml blank titrant \

ppmS~2 as H2S=^ r X r^ r x0.1704x( , , .., . )

density ml sample \ ml sample titrant/

Report sulfides concentrations of < 10 ppm to 1 decimal place and 

of > 10 ppm to 2 significant figures only.

Preparation of reagents

Iodine solution, O.OIOA7 : Dissolve 6 g iodate-free KI, in approx 25 ml water. 

Add 1.2691 g resublimed L. When solution is complete, dilute to 1,000 ml. 

Standardize with O.OIOA7 Na2 S?Oa using starch as indicator.

XT Normality ... , , 0.010Xml Na2S2 O3 of T2= :~T 

ml J-2

Sodium thiosulfate, O.OIOA7 .- Dissolve 2.482 g Na2 S2O3 -5H2O in carbon dioxide-free 

water and dilute to 1,000 ml with carbon dioxide-free water. Standardize 

against KIO3 as follows: Dry approx 0.5 g KIO3 for 2 hr at 180° C. Dissolve 

0.3567 g in water and dilute to 1,000 ml. Pipet 25.00 ml KIO3 solution into 

250-ml Erlenmeyer flask, then add successively 75 ml water and 0.5 g iodatefree KI. After solution is complete, add 10 ml cone HC1 (sp gr 1.19). 

Allow the stoppered flask to stand 5 min in the dark and titrate with Na2S2O3 

using starch indicator as end point is approached (light straw color).

0 25 Normality of Na2S2 O3 =-ml Na2S2O3

Potassium iodide, iodate-f ree: The KI can be tested for ICV1 by dissolving about 

0.1 g in 5 ml water, acidifying with 1 or 2 drops cone H2 SO} (sp gr 1.84) and 

adding 2-3 ml starch indicator. Immediate appearance of a blue color indicates the presence of HV1 ; slow color formation is due to atmospheric 

oxidation.

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3. 

Kolthoff, I. M., and Sandell, E. B., 1952, Textbook of quantitative inorganic

analysis: New York, Macmillan Co., 3d ed. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts., v. 61, no. 11.
D:40 TUBBIDITY

Turbidity is the optical property of a suspension with reference 

to the extent to which the penetration of light is inhibited by the 

presence of insoluble material. Turbidity is a function of both the 

concentration and particle size of the suspended material. Although 

it is reported hi terms of parts per million of silica, it is in no way 

synonymous with the weight of sediment per unit volume of water. 

The standard for turbidity (Newell, 1902) is

... a water which contains 100 parts of silica per million in such a state of 

fineness that a bright platinum wire 1 millimeter in diameter can just be seen 

when the center of the wire is 100 millimeters below the surface of the water 

and the eye of the observer is 1.2 meters above the wire, the observation being 

made in the middle of the day, in the open air, but not in sunlight, and in a 

vessel so large that the sides do not shut out the light so as to influence the 

results. The turbidity of such water shall be 100.

This definition in effect describes the size of the particles to be used 

as standard solutions. With this size particle, and only this size, 

1 ppm silica (turbidity) is equal to 1 ppm silica (weight).

The U.S. Public Health Service (1946) recommends that the 

turbidity not exceed 10 ppm in drinking and culinary water on 

carriers subject to Federal quarantine regulations. Some species 

of fish require clear water, but others thrive in somewhat turbid 

water. In general, turbidity adversely affects fish production by 

excluding light and thereby interfering with the growth of plants 

important in the fish-food cycle (California State Water Pollution 

Control Board, 1952, 1954, p. 226-228). Turbid water is abrasive 

to pipes, pumps, and turbine blades. In process water, turbidities 

much more than 1 ppm are not tolerated by several industries, but 

others permit up to 50 ppm or higher.

D:40a~l HELLIGE TURBIDIMETER

Principle of determination

The optical phenomena which cause the opalescent or "milky" 

appearance of a fine suspension are light scattering and absorption. 

Which of the two effects predominates in a particular case depends 

upon the size and the concentration of the particles. Inasmuch as 

the subjective reaction to these phenomena cannot be accurately 

duplicated by means of the turbidimeters on the market at the present 

time, it is quite possible to find instrumental measurements of this 

property at variance with the observers' visual reaction. Sometimes 

an instrumental measurement will designate one sample as having 

higher turbidity than a second sample, whereas the observers' 

reaction will be the reverse.

289
290 ANALYTICAL PROCEDURES D:40&-1

At the present time the property called turbidity is defined by the 

instrument used for its measurement. While the basic instrument is 

considered to be the Jackson candle turbidimeter, there are many 

other instruments in use. Readings from one to the other are not 

universally comparable. While all the common visual instruments, 

including the Jackson device, utilize both the absorption and scattering effect in making the measurement, there is wide variation in the 

degree to which each effect predominates in the various instruments. 

In order to achieve uniformity consistent with convenience, the Hel- 

\lige turbidimeter has been selected as the basis for the turbidity 

measurement herein described. See sec. C: 2g for a description of the 

Hellige turbidimeter.

Apparatus and reagents

Turbidimeter, Hellige

Illumination bulbs and calibration curves. New bulbs should be calibrated 

against standards prepared by the manufacturer and not against laboratoryprepared suspensions of fuller's earth or other similar material.

Procedure

Meticulous cleanliness of the sample containers and all parts of the 

instrument is mandatory, and the manufacturer's instructions in the 

operation and maintenance of the instrument should be followed 

explicitly.

Calculations

Determine the turbidity of the sample from appropriate calibration 

curves. Report turbidity as follows:

Turbidity 

(as ppm 8iOa )

0-1

1-10

11-100

101-400

401-700

^701

Record to the 

nearest 

0.1

1

5

10

50

100

REFERENCES

California State Water Pollution Control Board, 1952, Water quality criteria:

Pub. no. 3.

 1954, water quality criteria : Pub. no. 3, Addendum no. 1. 

Newell, F. H., 1902, Measurement of turbidity and color: U.S. Geol. Survey,

Div. Hydrography Circ. 8. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts.. v. 61, no. 11.
D:41 VISCOSITY

Viscosity is a characteristic property of fluids. It may be described 

in simple terms as an internal friction or an internal resistance to 

motion and is defined as the tangential force in dynes that must be 

exerted between 2 parallel layers, 1 sq cm in area and 1 cm apart, 

in order to maintain a velocity of streaming of 1 cm per sec of 1 

layer past the other (Glasstone, 1946, p. 95-155). Such a fluid would 

have a velocity of 1 poise (dynes per cm2 per sec). A more common 

term for reporting viscosity is the centipoise, which is 1/100 of a 

poise. Since the viscosity of water is approximately 1 centipoise 

at 20 °C, it seems desirable to report viscosity in centipoises at 20° C.

The experimental methods for the determination of viscosity are 

based on the fundamental equation of viscosity derived by J. L. M. 

Poiseville in 1844

relating the coefficient of viscosity, »/, to the volume, F, of gas or liquid 

which will flow through a narrow tube of L length and r radius in t 

time when under the influence of p driving pressure (Glasstone, 1946) .

D:41a-l BROOKFIELD VISCOMETER

Principle of determination

The principle upon which the Brookfield viscometer operates is 

the measurement of the drag produced upon a spindle rotating at 

a definite constant speed while immersed in the sample. The torque 

on the spindle rotating at constant speed is measured by a cilibrated 

spring. Viscosity changes appreciably with changes in temperature; 

therefore, the viscosity should be measured at constant or controllable temperature.

When measuring viscosities within the range of 0-10 centipoises, 

the results are reproducible to ±0.02 centipoises.

Apparatus

Brookfield Viscometer, model LV with UL adapter

Procedure

Follow the instructions and procedure that the manufacturer 

issues with the viscometer.

Calculations 

Report viscosity to 0.01 centipoise at 20° C.

REFERENCE

Glasstone, Samuel, 1946, The elements of physical chemistry : New York, D. Van 

Nostrand Co., Inc.

291

D:42 ZINC

Zinc is abundant in rocks and ores but is only a minor constituent 

in natural water because the free metal and its oxides are only sparingly soluble. In most alkaline surface and ground waters it is present only in trace quantities, but more may be present in acid water. 

Chlorides and sulfates of zinc are highly soluble. Zinc is used in 

many commercial products, and industrial wastes may contain large 

amounts.

Zinc in moderate concentrations is not known to have adverse 

physiological effects on man or stock, but zinc salts give water an 

unpleasant astringent taste and form a greasy film on boiling water 

(Howard, 1923, p. 411). The U.S. Public Health Service (1946, p. 

13) recommends that the concentration of zinc not exceed 15 ppm 

in drinking and culinary water on carriers subject to Federal quarantine regulations. Small quantities of zinc are toxic to various 

aquatic animals and plants. The element may also have such a toxic 

action on purifying bacterial flora of streams as to present serious 

sewage pollution problems. The concentration of 0.3 ppm is toxic 

to some fresh-water fish (Ellis, Westfall, and Ellis, 1946). High 

concentrations of zinc are reported to be detrimental to some crops 

(Kelly and Brown, 1938).

D:42a-l DITHIZONE EXTRACTION METHOD

The dithizone extraction method is similar in substance to part IA 

(tentative), APHA (1955, p. 212-214) Standard Methods.

Principle of determination

Zinc forms a keto complex with dithizone. This complex is soluble 

in carbon tetrachloride or chloroform and imparts a bright-red color 

to the solution. Many other metals react the same way with dithizone as does zinc. At a pH of 4.0-5.5 sodium thiosulfate largely 

prevents the reaction of copper, mercury, silver, gold, bismuth, lead, 

and cadmium with dithizone while allowing the reaction of zinc to 

proceed. In the presence of much nickel or cobalt, potassium cyanide 

can be used as the complexing agent.

The reaction may not go to completion because the presence of 

thiosulfate also retards the reaction between zinc and dithizone. 

Therefore, the determination is to some extent empirical, and it is 

essential to adhere to directions as to volumes, time of extraction, 

pH, etc. Even analytical-grade reagents often contain sufficient 

zinc to be detected by the procedure. Instructions are given for 

extraction of prepared reagents with dithizone before use. Some

293
294 ANALYTICAL PROCEDURES D:42a-l

batches of dithizone may require purification; when necessary, the 

reagent can be purified by the ammonia extraction method (Sandell, 

1950). Zinc may also be dissolved from glassware and some stopcock grease. All glassware should be rinsed with dilute nitric acid 

before use. Dithizone and dithizonates decompose rapidly in a 

strong light. Hence, samples and standards should be carried 

through the procedure simultaneously, and the color comparison 

should be made before an undue lapse of time.

In the absence of excessive quantities of other metals that react 

with dithizone at the pH of the test, results are accurate and reproducible to ±0.0005 mg.

Apparatus and reagents

Matched test tubes, glass-stoppered 

Separatory funnels, 125-ml 

Zinc chloride, 1.00 ml=0.010 mg Zn+2 

Hydrochloric acid, 0.052^ 

Sodium hydroxide, 0.052^ 

Methyl orange indicator solution 

Acetate buffer solution 

Sodium thiosulfate, 25 percent 

Dithizone, 0.001 percent in CCh

Procedure

Samples for the determination of zinc should be treated in accordance with directions given in sec. A: 4d.

1. Pipet a volume of sample containing less than 0.005 mg Zn+2 (50.0 ml max) 

into a separatory funnel and adjust the volume to 50 ml with metal-free 

water.

2. In separatory funnels, prepare a blank and sufficient standards, and adjust 

the volumes to 50 ml with metal-free water.

3. Adjust the sample and standards with 0.052^ HC1 and 0.052^ NaOH until they 

are just acid to methyl orange.

4. Add 25 ml acetate buffer.

5. Add 5 ml 25 percent Na2 S3O3, and mix.

6. Add 5.0 ml 0.001 percent dithizone.

7. Extract by shaking vigorously for 2.0 min.

8. Withdraw the CCU layer into a glass-stoppered test tube and stopper tightly 

to prevent loss of solvent.

9. Compare the sample and standards by viewing transversely against a white 

background. The color changes from green to red with increasing concentrations of zinc. Record the zinc concentration of the sample to the nearest 

0.0005 mg Zn+2.

Calculations

ppm Zn+2=-] T7- X '° . X mg Zn+2 in sample density ml sample *
D:42a-2 ZINC 295

Report zinc concentrations of <1 ppni to 2 decimal places and of 

> 1 ppm to 2 significant figures only.

Preparation of reagents

Zinc chloride, 1.00 ml=0.010 mg Zn*2 : Dilute 10.00 ml ZnCL (1.00 ml=0.10 mg

Zn+a ) to 100.0 ml with metal-free water. Prepare immediately before using. 

Zinc chloride, 1.00 nil=0.10 mg Zn+2 : Dissolve 0.100 g reagent-grade Zn° (30-

mesh) in a slight excess of dilute HC1 and dilute to 1,000 ml with metal-free

water. 

Hydrochloric acid, Q.Q5N: Dilute 4.2 ml cone HC1 (sp gr 1.19) to approx 1 liter

with metal-free water. 

Sodium hydroxide, Q.Q5N: Dissolve 2.0 g NaOH in metal-free water and dilute

to approx 1 liter with metal-free water. 

Methyl orange indicator solution: Dissolve 0.5 g methyl orange in water and

dilute to approx 1 liter. 

Acetate buffer solution: Mix equal volumes of 2N NaCzHaO^ and 2N HCzHaOz.

Remove heavy metals by repeated extraction with 0.005 percent dithizone

until the CC1« remains green. Rinse out the dithizone by extracting with

CCl,.

Sodium acetate, 2N: Dissolve 136 g NaC,H3(V3H2O in 500 ml metal-free water. 

Acetic acid, 2N: Dilute 57 ml glacial HC2H3O2 to 500 ml with metal-free water. 

Dithizone, 0.005 percent in COU. Dissolve 12.5 mg diphenylthiocarbazone in

approx 250 ml CC14. Stopper tightly and store in the refrigerator. 

Sodium thiosulfate, 25 percent: Dissolve 25 g NajSzQvSHzO in approx 100 ml

metal-free water. Extract repeatedly with 0.005 percent dithizone until the

CCU remains green. Rinse out the dithizone by extracting with CCl*. 

Dithizone, 0.005 percent in CCU. Dissolve 12.5 mg diphenylthiocarbazone in

1 liter CCU. Stopper tightly and store in the refrigerator.

D:42a-2 DIRECT COLOR-COMPABISON METHOD

The method is not as specific for zinc as D: 42a-l but is more rapid 

if copper is also determined.

Principle of determination

In a buffered solution, zinc forms a red keto complex with dithizone 

in acetone. Sodium thiosulfate inhibits the interference of other 

metals that react with dithizone. Sodium citrate is used to complex 

iron, which oxidizes dithizone. The keto complex is soluble in wateracetone media, and no extraction is required.

Copper also forms a complex which absorbs light in the same range 

as the zinc complex. The absorbancy of the zinc and copper complexes 

determined collectively are corrected for that attributable to copper.

In the absence of other interfering metals, results are accurate and 

reproducible to ± 0.0005 mg.

Additional information on the principle of the determination is 

given by Sandell (1950, p. 87-113).
296 ANALYTIC/ L PROCEDURES D:428r-2

Apparatus and reagents

Spectrophotometer, Beckman Model B : 

Wavelength : 550 m/t 

Cells : 40-mm optical depth, cylindrical 

Phototube : Blue-sensitive 

Initial sensitivity setting : 3 

Slit width : 0.3 mm (approx) 

The following absorbancies have been observed : 

mg Zn Absorbancy

0. 001 0. 068

. 002 . 133

.004 .260

. 008 . 515

Zinc chloride, 1.00 ml=0.004 mg Zn+2

Sodium thiosulfate, 25 percent

Acetate-citrate composite solution

Acetone

Dithizone, 0.05 percent in acetone

Copper solution, 1.00 ml=0.010 mg Cu*2

Procedure

Samples for the determination of zinc should be collected in accordance with directions given in sec. A : 4d.

1. Determine the copper content of the sample by the procedure described in sec. 

D : 14a-l.

2. Pipet a volume of sample containing less than 0.01 mg Zn+2 (10.00 ml max) 

into a 50-ml beaker and adjust the volume to 10.0 ml with metal-free water.

3. Prepare a blank and sufficient standards, and adjust the volumes to 10.0 ml 

with metal-free water.

4. Add 1.0 ml 25 percent Naa&Os and mix.

5. Add 1.0 ml acetate-citrate composite solution and mix.

6. Add 10.0 ml acetone and mix.

7. Add 2.0 ml 0.05 percent dithizone and mix.

8. Cover the beakers to minimize evaporation.

9. Immediately determine the combined absorbancy of the test sample and 

standards against the blank, and when necessary make correction for water 

color as directed in sec. C : la-2, method 1.

10. Prepare a correction curve for copper contributing to absorbancy by treating 

standard solutions by the procedure given above.

Calculations

1. Correct absorbancy for that attributable to copper.

2. Determine quantity of Zn+2 in test sample from a plot of corrected absorbancies of standards containing known amounts of constituent.

3. ppm Zn+2 =-5 rr~ X rj r~ X nig Zn+2 in sample 

density ml sample

Report zinc concentrations of <1 ppm to 2 decimal places and of 

> 1 ppm to 2 significant figures only.
D:42ar-2 ZINC 29-7

Preparation of reagents

Sodium thiosulfate solution, 25 percent: Dissolve 250 g Na2S2Cv5H2O in metalfree water and dilute to approx 1 liter. Extract with dithizone in CCU until 

titrant shows no evidence of zinc.

Acetate-citrate composite solution: Dissolve 340 g NaC2HsCv3H2O in metal-free 

water, add 1.5 g Na3CBH5Cv2H2O, and dilute to 1,000 ml. Extract with 

dithizone in CCU until titrant shows no evidence of zinc.

Dithizone reagent, 0.05 percent: Dissolve 0.100 g diphenylthiocarbazone in 200.0 

ml acetone. This reagent must not be refrigerated (dithizone will precipitate at low temperatures).

Zinc chloride, 1.00 ml=0.0004 mg Zn*2 : Dilute 10.00 ml ZnCl2 (1 ml=0.10 mg 

Zn+2 ) to 250.0 ml with metal-free water. Dilute 10.00 ml of this intermediate 

solution to 100.0 ml with metal-free water. Prepare immediately before 

using.

Zinc chloride, 1.00 ml=0.10 mg Zn+2 : Dissolve 0.1000 g reagent-grade Zn° (30- 

mesh) in a slight excess of dilute HC1. Dilute to 1,000 ml with metal-free 

water.

Copper solution, 1.00 ml=0.010 mg Cu*2 : Dissolve 0.0100 g Cu° in a minimum of 

dilute HNO3 and dilute to 1,000 ml with metal-free water.

REFERENCES

American Public Health Association and others, 1955, Standard methods for the

examination of water, sewage, and industrial wastes: New York, Am. Public

Health ASSOC., Inc., 10th ed. 

Ellis, M. M., Westfall, B. A., and Ellis, M. D., 1946, Determination of water

quality: U.S. Fish and Wildlife Service Reserve Rept. 9. 

Howard, C. D., 1923, Zinc contamination in drinking water: Am. Water Works

Assoc. Jour., v. 10. 

Kelly, W. P., and Brown, S. M., 1938, Boron in soils and irrigation waters of

southern California and its relation to citrus and walnut culture: Hilgaria,

v. 3. 

Sandell, E. B., 1950, Colorimetric determination of traces of metals: New York,

Interscience Publishers, Inc., 2d ed. 

U.S. Public Health Service, 1946, Drinking water standards: U.S. Public Health

Service Repts., v. 61, no. 11.

INDEX

[Boldface entries indicate water properties; boldface numbers, beginning of their major description]

Absorbancy, in spectrophotometry, denned_ 44

Acidity 87

immediate, volumetric method ______ 90

potential free, volumetric method ____ 91

total, volumetric method_________ 88

Alkalinity, potentiometric method_ ____ 93,94

Aluminum, ferron-orthophenanthroline

method______________ 97

Arsenic, Gutzeit-volumetric method_ ___ 101

Ammonia-free water, denned_________ 76

Analyses, methods of checking__ _____ 81

order of determinations__________ 78

Analytical methods, choice of._.______ 75

B

Balances, mechanical ____.______ 60 

Balancing analyses._____________ 81 

Ball-valve samplers_______ ___ ___ 12,13 

Barium, complexometric method______ 107,109 

gravimetric method...__________ 107 

Bausch and Lomb spectrophotometer, sensitivity.. ...-...-...--.-...__ 66 

Beckman spectrophotometers, accuracy___ 46 

cell carriers________________ 64 

flame attachments_ __________ 56,68 

fuel selection.. _______________ 56 

photometric titration assembly..... __.. 71,72 

photomultiplier attachment...._.___ 58 

sensitivity_._______________ 66 

Beer's law____. ___________. __ 45 

Biochemical-oxygen demand, sampling for

determination of.._________ 13

Boats, precautions in sampling from_____ 18

Boron, carmine method__ ________.. 113,118

dianthrimide method..._________ 113

potentiometric method___ ______ 115

Bottom sampler_____.___.___. 10,12,13

Bridges, precautions in sampling from____ 17

Brin^-sample volumes, accuracy in analyses,. 77

Bromide, bromine displacement method.... 121,124

oxidation method_______.____ 121

Cableways, precautions. . . ... 18

Calcium, eomplexometricmethod ._ . . 127

gravimetric method........-. . -. 131

permanganimetric method_ ._ .. 129

turbidimetric method_______._... 134

Carbon dioxide, calculation method _... 137

Carbon dioxide-free water, defined .... .. 76

Chemical balance, in evaluating analyses... 81

Chemicals, purity.. .____-_------- 77

Chloride, gravimetric method_ .,_....- 141,143

volumetric method_...__ _... 141

Chlorine residual, orthotolidine-arsenite

method- . .. 147 

special treatment for . . . .. 29 

Chromium. . . 151 

chromium, permanganate-azide method.. _ 153 

hexavalent, diphenylcarbazide method .. 151 

Closed systems, samplers 14 

Color, comparison method_.-_-. 155 

correction for in spectrophotometric analysis._...__.._____-_ 52 

effect on spectrophotometric analysis 48,50 

instruments for measuring ... . 69 

notation on sample tag. 24,27 

Colorado River sampler. 12,13 

Composite samples, discharge-weighted. ... 35,37 

length of period . 34 

preparation for analysis ... - 78 

time-weighted__._.__-..-... .- 34,36 

Conductivity meters... _.--_- 58 

Copper, diethyldithiocarbamate spectrophotometric method___.-.. 157

D

Density, gravimetric method 161 

Depth-integrating samplers . 10,11 

Diluted conductance, calculation.._.._- . 84 

Dilution water, defined. .-.. .. 75 

Dissociation, in balancing analyses 83 

Dissolved gases, sampling for determination of. 13 

Drinking-water standards. See description of 

particular water property.

E

Equivalents per million, computation 80

Evaporation, vessels for 73

Extinction, in spectrophotometry, defined... 45

Filtration....-. - 39,78

Fixatives for samples.. 27

Flame photometry, auxiliary equipment 69

fuel selection. .. 55

types of instruments 56,58,68

use of standards 57

Fluoride, isolation by distillation...-.-. 163,169

zirconium-alizarin method 166

zirconium-Eriochrome Cyanine R method 163

Foerst sampler__._...- - 12,13,14

Freezing of samples, effect on analyses. 9

G

Glassware, cleaning.. . 9,49,73 

laboratory equipment.... 72 

sample bottles_-... 8 

volumetric, calibration and tolerances 73,74

299
300 INDEX

Gravimetric analysis, disadvantages...

Ground water, representative samplessafety precautions in sampling..... 

samplers.._________._. 

sampling data required_.___._-. 

sampling sites _________..

H

Hardness. _................................... 173

carbonate and noncarbonate, calculation

method______.____.. 177

hardness, calculation method______ 174,176

complexometric method ____ 174

Hellige turbidimeter.... -..-... .........-.. 289

Hydrogen-ion concentration, relation to pEL.. 81

Ice, precautions in sampling through_____ 19 

sampling through for representative

samples _.__._______ 20

Iodide, oxidation method__ __ _ _. __ 179 

starch method________________ 181

Iron__________________________ 183 

ferrous and ferric, bipyridine method.__ 186 

iron, bipyridine method_________ 183,184

Jackson candle turbidimeter ...____ 70,290

Lakes, safety precautions in sampling........ 15

sampling sites.. _____.._____. 5,23

Lead, dithizone method___...._____ J89

Lithium, flame-photometric method.....___ 193

Loss on ignition, gravimetric method... ___ 195

M

Magnesium, calculation method___...._. 197

Eriochrome Black T method______ 201

pyrophosphate method__________ 198

Manganese, permanganate method...____ 205

tetrabase method_____________ 207

Metal-free water, defined._____..__.. 76

Multivalent ions, in balancing analyses.. _.. 82

N

Nitrogen-___________________.. 211

ammonia, Nesslerization method_..._ 211

special treatment for.________.. 28

test-paper method.._______ 215

titration method__.______.. 214

nitrate, phenoldisulfonic acid method..... 216

reduction method....________ 219

special treatment for_..__.._... 29

nitrite, diazotization method_______. 221

ppecial treatment for. See Nitrate

nitrogen.

organic, Kjeldahl method_....____.. 223 

special treatment for. See Ammonia

nitrogen.

Nonvolatile constituents, sampling for determination of..._.____..... 10

O 

Oils and waxes, wet-extraction method ..... 227

Optical density, in spectrophotometry, defined....______________ 45 

Oxygen consumed, permanganimetric method. 229 

Oxygen dissolved, Alsterberg (azide) method. 233 

special treatment, for_.____._..... 29

Partial analyses, checking.....__.____ 86

Perkin Elmer flame photometer, fuel selection. 55

pH value, field instrument method__.__ 237,238

instrument method_._________.. 237

meters____.__..__.__.. .__ 60 

paper method_______.__.__... 238 

relation of hydrogen-ion concentration.... 81

relation of orthophosphate mixtures.__ 82 

Phenolic material, Gibbs reaction___...... 239

preservative for.... -.... . _... 30

Phosphorous_____.___..__.____. 245 

orthophosphates, phosphomolybdate

method_...__.._.____ 245 

relation to pH_____________ 82 

special treatment for__ _. ___ _ 30 

phosphorous, phosphomolybdate method.. 247 

Photometric titration______________ 71 

Photovolt Lumetron spectrophotometer, photometric titration assembly___ 71,72 

sensitivity ------------------------- 66

Plastic ware, laboratory equipment___-_ 49,72 

sample bottles. __.._-__-____ 8,20 

Platinum ware, care and use _______ 73,74 

Point samplers.. _______________ 10,12,13 

Porcelain ware, cleaning...................... 73

Potassium, flame-photometric method __ 251 

Precipitation, effect on analysis .. .._... 78 

Preservatives for samples._...___...___ 27 

Pycnometers- _----- ... ~. ... 77

Pyrex ware, containers for reactions ____ 49,73 

sample containers._ --------- ....- 8

R

Range lines, sampling along...--- - 23

Reagents, accuracy in measurement . 76

effect on spectrophotometric analyses. -- 49

specifications for purity . ... _ 77

storage ____-_____-__ ___ 72

Redistilled water, definition_____ - ___ 75

Reservoirs, precautions in sampling. .._._ 15

sampling sites ._ . 5,23

Residual sodium carbonate _...... ... 93

Residue on evaporation, relation to calculated

dissolved solids 83

relation to specific conductance _. _. _.. 84

Sample bottles, cases .. .... 32

closures.___- -____. ... 10

effect of composition on analysis..-------- 8

Geological Survey_ _ _ __- 9

Sample tag, ground-water________... 25,26 

surface-water.-.------------ ----------- 21,22

Sample volumes, accuracy in analyses. 76,77

Samplers, closed systems..- ...... ..-.. 14

in determination of dissolved gases..__- 13 

in determination of nonvolatile constituents____________.___ 10 

suspension apparatus... ... ... 15
INDEX 301

Sampling, frequency____------ ----- 7

safety precautions.-- -_.... .- 15

stations along streams.---------.---- .. 21

stations on lakes and reservoirs__ .. 23

Selenium, Norris-Fay titration method. _ 253 

special treatment for. . . .. 30

Serfass Wheatstone bridge__________ 60

Significant figures in analyses . 76,79

Silica, gravimetric method . --...-- 259,261 

molybdate blue method- . .. .- 259

Silicone solutions, coating of glass..__----- 73

Sodium, flame-photometric method_-___. 265 

zincuranyl acetate method ------- 265

Solids .. . . . - --. 269

dissolved, calculation method.. _ _... 269,271

residue-on-evaporation method _ _. 270

suspended, filtration method..._...__ 273

Special treatment for samples _____._- 27

Specific conductance, Wheatstone bridge

method___ - --_-__. 275 

as a guide in compositing.-_____.__ 38 

in checking analyses_.---..---------_. 84

in selection of volume of sample. ... 78 

instruments___-.... ... - . 58

Spectrophotometers, absorption cells_____ 64 

accuracy_____..._ _ 46,67,68

filters____________________ 62 

photocells_ -. . . --. . .--. 64 

requirements... ..---- -. ........ 67

sensitivity__________._-_-__ 46,66 

types--.----..-._ - 63,66

Spectrophotometric analysis, advantages.__ 43 

effect of color.________-_-------- 50

effectof turbidity - -. 53

sensitivity . 48

use of standards.. ______ - _ - ____ 49

Storage of samples, effect on analyses.- 31 

precautions__ .. 32

Streams, representative samples of. ..-... 20 

safety precautions in sampling _ ___ . 15

Streams C ontinued

samplers.._._...._... . 10 

sampling sites_____.__ . 3-5,21

Sulfate, gravimetric method___._... 279,283 

Spectrophotometric thorin method __... 281 

turbidiraetric method.... . 284 

visual thorin method___._-.--.---... 279

Sulfides, iodometric method . .. 287 

special treatment for 30

Surface water, representative samples...___ 20 

samplers__________-- ------ 10

sampling data required_... 21 

sampling sites.______-_..._.. . 3

T

Tidal water, sampling...___._._____ 4 

Titration, photometric, instruments_____ 71 

Transmittance, in spectrophotometry, defined ... 44 

Transmittancy, in spectrophotometry, defined..______________ 44 

Turbidimeters____._. _ 70 

Turbidity, effect on Spectrophotometric analyses... ... - 48,53,289 

Hellige turbidimeter.. ... . .. 70 

notation on sample tag_____..... 24,27 

removal from samples_________ 39,54,78

Viscosity, Brookfield viscometer____-_- 291 

Volumetric analysis, advantages. . 43

W

Wading, precautions_..... 

Water-color comparators. 

Waxes. See Oils and waxes. 

Wheatstone bridge... ....

15

58

Zinc, direct color-comparison method ... 293,295 

dithizone extraction method. 293

O