The nature of organic color in water.

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Title:
The nature of organic color in water.
Uncontrolled:
Color in water, The nature of
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viii, 120 l. : ill. ; 28 cm.
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English
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Christman, R. F ( Russell F. ), 1936-
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s.n.
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Gainesville
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Water -- Analysis   ( lcsh )
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non-fiction   ( marcgt )

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Thesis:
Thesis--University of Florida.
Bibliography:
Bibliography: l. 114-119.
Statement of Responsibility:
By Russell Fabrique Christman.
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Manuscript copy.
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Vita.

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University of Florida
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THE NATURE OF ORGANIC

COLOR IN WATER











By
RUSSELL FABRIQUE CHRISTMAN


A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL OF
THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE
DEGREE OF DOCTOR OF PHILOSOPHY











UNIVERSITY OF FLORIDA
December, 1962











ACKNOWLEDGMENTS


The author wishes to express his earnest appreciation to his

committee chairman, Dr. A. P. Black, for guidance and encouragement

given during this investigation and throughout the author's graduate

training. He would also like to thank the other members of his

supervisory committee, Dr. W. S. Brey, Dr. J. M. Pearce, Dr. H. E.

Schweyer, and particularly Dr. J. D. Winefordner.

Special thanks are due many of the author's associates, princi-

pally Mr. Howard Latz and r.John Wilson for their many helpful sugges-

tions.

The author also wishes to express his gratitude to the National

Institutes of Health whose financial support made this project possible.

Finally, the author gratefully acknowledges his indebtedness

to his wife, Sylvia, for her support and understanding.












TABLE CF CONTENTS


Page


ACKNOWLEDGMENTS .

LIST OF TABLES .

LIST OF FIGURES .

CHAPTER

I. INTRODUCTION

II. HISTORICAL .

III. EXPERIMENTAL

IV. EXPERIMENTAL

V. SUMMARY .

APPENDIX .

LIST OF REFERENCES. .

BIOGRAPHICAL SKETCH .


* m 6 0 5 S

PROCEDURES

RESULTS .


* S S 0

f f ft


* 5 0 5 U 5 0 a a S 0 a

S U 6 0 o 0 a


S S S a a 0 O 0 S S 0 0 5 S .a


ii

iv

vi




1

2

22

27

71

74

114

120











LIST OF TABLES


Table Page

1 Classical Fractions of Soil Organic Matter 8

2 Molecular Weights of Some Humic and Hymatomelanic
Acids 14

3 Effect of Storage on Color in a Nitrogen Atmosphere 27

4 Effect of Storage on Color in an Osygen Atmosphere 28

5 Sources and Designations of Colored Waters 29

6 Chemical Analyses of Waters A Through E 31

7 Chemical Analyses of Waters F Through J 32

8 Variation of Iron Content in Colored Water Samples 33

9 Variation of Iron Content in Colored Water Samples 33

10 Variation of Color with pH 34

11 Electrodialysis and Millipore Filtration of Water A 36

12 Scattering Ratios of Latex Suspensions 37

13 Solubility Relationships of Fractions Obtained by
Shapiro Technique 44

14 Weight Distribution in Fractions Obtained by
Shapiro Technique 44

15 Ultimate Analyses of Fractions Ia and Ib 48

16 Weight of Organic Matter in Fractions 49

17 Percentage of Total Organic Matter in Fractions 50

18 Infra-Red Absorption Bands of Fulvic Acid Fractions 53

19 Ultimate Analyses of Fractions of Water A 55

20 Paper Chromatographic Separation of Fulvic Acid
Fraction of Water A 56





LIST OF TABLES (Continued)


Table Page

21 Paper Chromatographic Separation of Humic Acid
Fraction of Water A 57

22 Paper Chromatographic Separation of Hymatomelanic
Acid Fraction of Water A 57

23 Equivalent Weight and Quantitative Methylation
Data of Fulvic Acid Fractions 60

24 Quantitative Methylation of Model Compounds 61

25 Effects of C12, ClO2 and 03 on Color Removal 64

26 Effects of C12, C10 and 03 on Chemical Oxygen
Demand of Colored Water Concentrates 64

27 Infra-Red Absorption Bands of Ether Soluble Chlorine
Oxidation Products 65











LIST OF FIGURES


Figure Page

1 Effect of pH on Color of Waters A and G 75

2 Effect of pH on Color of Water B 76

3 Effect of pH on Color of Waters D and E 77

4 Dialysis and Millipore Filtration of Waters A Through E 78

5 Dialysis and Millipore Filtration of Waters F Through J 79

6 Light Scattering Properties of Water B at 546 ma 80

7 Light Scattering Properties of Water B at 436 m4 81

8 Effect of pH and Constant Alum Dosage on Coagulation
of Water B 82

9 Effect of pH and Constant Alum Dosage on Coagulation
of Water I 83

10 Variation of pH Width of Optimum Color Removal with
pH Width of Charge Reversal 84

11 Spectrum of Mercury Lamp Source with Distilled Water 85

12 Fluorescence Spectrum of Water B 86

13 Fluorescence Spectrum of Water F 87

14 Effect of pH on Fluorescence Emission of Water B 88

15 Ultraviolet Absorption of Water B as Function of
Color Value 89

16 Ultraviolet Absorption of Water I as Function of
Color Value 90

17 Absorption of Colored Waters at 220 m) 91

18 Absorption of Colored Waters at 300 mn 92

19 Absorption of Colored Waters at 350 mu 93





LIST OF FIGURES (Continued)


Figure Page

20 Effect of pH on Ultraviolet Absorption of Water B 94

21 Effect of pH on Ultraviolet Absorption of Water C 95

22 Infra-Red Spectrum of 'Free Acid" Fraction II 96

23 Fractionation Scheme 97

24 Relation Between Color Value and Total Organic
Content of Colored Waters 98

25 Relation Between Color Value and Concentration for
Humic and Hymatomelanic Acid Solutions 99

26 Infra-Red Absorption Spectrum of Fulvic Acid Fraction
of Water B 100

27 Infra-Red Absorption Spectrum of Fulvic Acid Fraction
of Water C 101

28 Infra-Red Absorption Spectrum of the Methyl Ester of
Fulvic Acid Fraction of Water A 102

29 Potentiometric Titration of Benzoic Acid in
Ethylenediamine 103

30 Potentiometric Titration of Phenol in Ethylenediamine 104

31 Potentiometric Titration of Resorcinol in Ethylene-
diamine 105

32 Potentiometric Titration of o-Hydroxybenzoic Acid
in Ethylenediamine 106

33 Potentiometric Titration of p-Hydroaqbenzoic Acid
in Ethylenediamine 107

34 Potentiometric Titration of Fulvic Acid Fraction
of Water B in Ethylenediamine 108

35 Potentiometric Titration of Fulvic Acid Fraction
of Water I in Ethylenediamine 109

36 Infra-Red Absorption Spectrum of Chlorine Oxidation
Product from Water A 110

37 Infra-Red Absorption Spectrum of Chlorine Oxidation
Product Extracted from Miami Tap Water 111


vii





LIST OF FIGURES (Continued)


Figure Page

38 Infra-Red Absorption Spectrum of a Fulvic Acid
Oxidation Product 112

39 Infra-Red Absorption Spectrum of a Fulvic Acid
Oxidation Product 113


viii











I. INTRODUCTION


Waters containing color resulting from their natural environ-

ment are found throughout most parts of the world. While organic color

is most frequently found in surface waters, it is occasionally found in

shallow or deep wells in limestone regions where solution topography

prevails.

The United States Public Health Service (1) has set a maximum

limit for drinking waters of 15 color units on a platinum-cobalt scale.

This limit has been imposed mainly for aesthetic reasons as the materials

responsible for color in water are not known to be physiologically

harmful. In practice, most municipal treatment plants strive to maintain

finished waters with color values of 10 or less.

Almost all of the research to date involving naturally colored

water has been concerned either with empirical methods of treatment for

color removal or with the mechanism of color coagulation. Little

attention has been directed to the fundamental nature of the organic

materials that are responsible for color in water. Therefore, the pri-

mary purpose of this investigation is to study various physical and

chemical properties of organic color in water, in the hope that such basic

information may eventually lead to a more complete understanding of

treatment processes.

Because of the widespread occurrence of this type of water, it

is an additional objective of this research to determine the uniformity

of these properties in colored waters from a variety of sources.











II. HISTORICAL


Nature of Organic Matter in Water


Aschan (2) carried out the first extensive research on these

substances with an analysis of six Finnish lake and river waters. The

humus material was precipitated with FeC13 and subjected to ultimate

analysis with the following results: carbon content was found to vary

from 44.99 to 54.10 per cent, hydrogen 3.86 to 5.05 per cent, nitrogen

1.46 to 4.23 per cent and oxygen from 38.76 to 47.93 per cent. He con-

cluded that these materials were acids of strength comparable to phenol.

Birge and Juday (3) observed that the carbon/nitrogen ratio was

larger in more highly colored lakes, a fact which has led many authors

to regard nitrogen as an impurity.

Saville (4) reported from the results of some qualitative

electrophoretic experiments that organic color is present in water as

a negatively charged colloid. Behrman, Kean and Gustafson (5) found

that little color in water would dialize through a parchment membrane,

indicating a colloidal nature. These authors also noted the indicator

action of colored water and that most of the color could be oxidized by

chlorine.

Christman (6) determined the rate of migration in cm/sec/volt

of organic color bodies in 32 natural waters and of two aqueous dye

solutions in an electrophoresis cell. All colloids present in the 32

waters were found to be negatively charged.







In 1910, Dienert (7) detected fluorescent materials present in

natural waters and noted that this property was lost by mild oxidation.

The use of absorption spectra analyses for the investigation

and routine checking of contaminated surface water was first described

in 1936 by Denmering (8). This author found the absorption bands pro-

duced by lignin substance offered a means of checking changes in the

concentration of organic material. Demmering applied this as a

criterion for amount of chemicals to be used for chemical treatment.

Datsko (9) carried out photometric investigations of color re-

actions of sulfuric acid with organic substances in natural waters. He

observed an intense coloration produced in the reaction and noted that

the intensity followed Beer's law. No description was furnished as to

the manner of preparation of his humic acid.

Keiser (10) has reported a differentiation of organic substances

in natural water. According to this author alum removes more of the

oxygen oxidizable material than the chlorine oxidizable material.

More recently, Flaig and Otto (11) have carried out investiga-

tions of humic materials and model substances on plant life. These

authors added polyhydroxyanthraquinone to spring water at a concentration

of 10-5 gms/ml and found that the plants grew normally, using this

material precisely as plants use natural humic acid.

Goryunova (12) has investigated the character of dissolved

organic substances in water of Glubokoe Lake, Russia. The water of this

lake was found to contain 96 gms/m3 of organic matter, the bulk of this

material being made up of high molecular weight fatty acids and colloidal

complexes. He also noted small amounts of substances that would reduce

Fehling's solution.







According to Isachenko and Egorova (13) the earthy smells in

reservoirs are caused by actinomycetes. Odors of this type have been

noted frequently in concentrates of organic colored water.

In 1954 Ponomareva and Ettinger (14) obtained dark, easily

precipitated substances of the humic-ulmic acid type from organic matter

in waters of the Neva River. This material was found to comprise only

5 to 6 per cent of the total organic matter. In addition, apocrenic

acid concentrations were found to reach the same percentage. The

remainder of the organic material was thought to consist of lightly

oxidized fractions of crenic acids.

Goryunova (15) has isolated and identified 50 per cent of the

organic matter in the water of Beloe Lake, Russia. This author concluded

these are humus materials but fatty acids and polysaccharides are also

present.

Skopintsev and Krylova (16) have stated that aqueous humus of

terrigenous origin is responsible for the yellow color of water. These

authors found a linear relation between color intensity determined

colorimetrically and the coefficient of extinction determined spectro-

photometrically. No such relation was found to exist between the total

organic content and either color intensity or extinction coefficient.

Raudnitz (17) isolated a water soluble, surface-active phosphoric

ester of humic acid from rhododendron leaves. Subsequent hydrolysis

yielded a uniform humic acid free of nitrogen and sulfur, insoluble

in common organic solvents but soluble in pyridine and aniline. Infra-

red data showed the presence of hydroxyl and carbonyl groups.

Wilson (18) stated that colored organic materials in water are

fulvic acids, a fraction of natural soil humus. The author believed no







specific chemical test for these materials will be developed because

of the diversity of functional groups present, but a quantitative

estimation of these acids can be obtained from their ultraviolet absorp-

tion spectra. Using a standard solution of fulvic acids from

phragmites peat, Wilson observed that for a given concentration absorption

at 300 m)n was constant in the pH range 1.0 to 5.0 increasing over pH 6.0.

Application of this technique to several natural waters led to the

following conclusions (a) All samples contained appreciable quantities

of fulvic acids. (b) All absorption spectra were similar. (c) Since

fulvic acid is the most water soluble fraction of natural soil humus it

would be expected to be found in natural waters in higher concentrations

than either humic or hymatomelanic acids.

Fouling of anion exchange resins by organic materials in colored

water is a problem of some industrial significance. Numerous authors

have observed organic fouling of strongly basic anion exchange resins

of the quaternary types I and II and have ascribed the differences in

their fouling rates to the structural differences in their quaternary

ammonium substituents. The foulants are held on the resins rather

strongly and are not removed by normal regeneration. Excellent reviews

of this problem are provided by Frisch et al. (19) and Frisch and

Kunin (20). In addition, Ungar (21) has reported fouling of weakly basic

anion exchangers and that resins of this type exhibit a limited exchange

capacity for humic acids.

One of the most recent publications in this area is an extensive

investigation of the yellow organic acids of lake water carried out by

Shapiro (22). This author has outlined a separation scheme for the

isolation of a "free acid" fraction of the organic materials present in







the water. Data obtained by infra-red and ultraviolet spectroscopy and

by various organic spot reactions led this author to report that these

acids are aliphatic unsaturated polyhydroxy dicarboxylic acids with an

approximate molecular weight of 456. No evidence was found in this work

of the existence of aromatic rings and dialysis studies indicated a lack

of colloidal character. Shapiro suggests the name "humolimnic" acids

for these materials to indicate their probable relation but improbable

equivalence with natural humic acids of the type found in soils.



The Nature of Organic Material ig Soils


Because of the relation of these coloring substances to the

organic matter present in soils, it is necessary to review the work in

this area of chemistry. No attempt will be made to critically outline

the early developments in this field as this has been done by Waksman (23).

However, the major findings will be discussed particularly where there is

an appreciable correlation with data obtained on the organic materials in

water systems.

The term humus dates back to the Roman era when it was generally

applied to the complete soil substance. From this time to around the

17th century little was known about the nature or functions of soil.

At the start of the 18th century Linnaeus classified soils in a manner

similar to his classification of plants, humus daedalea (garden soil),

humus ruralis (field soil), etc. In the latter half of the 18th century,

humus began to be considered in terms of decaying organic matter. How-

ever, the general ideas concerning the chemical nature of humus were

very vague.







Unfortunately, the term humus came into general use when organic

chemistry was in its infancy and was not always used to designate the

same organic materials. This was responsible for some degree of con-

fusion among workers in the field of soil science, a situation which

has never been entirely remedied.

Perhaps the earliest scientific classifications of humus are

chose of Oden (24) and Ramann (25). Ramann regarded humus as "colloidal

complexes of varying composition, consisting of unchanged colloids of

the original plant substance mixed with carbon-rich decomposition

products." Ramann recognized the importance of micro-organisms in the

synthesis of humus, a fact disregarded by Oden. The latter regarded

humus as "those dark colored bodies of unknown constitution which origi-

nate in nature through the decomposition of organic substances...; they

possess a definite affinity for water and show, if not true solution, a

distinct swelling."

Oden also supplied the first generally accepted system of naming

the four common fractions or preparations of soil organic matter. Up

until this time there were in common use 43 different names for the four

preparations.

Oden's classification scheme is shown in Table 1.

In 1930 this system was modified by Page (26), and resulted in

the most widely used classification system to date. He suggested the

name "humin" for preparation 4 and humicc acid" for preparation 2.

The multitude of names and classification systems during this

early period was due to a large extent to the variety of sources from

which the humus materials were obtained. The early theories about the

simplicity and uniformity of the substances in soil were gradually







TABLE 1

Classical Fractions of Soil Organic Matter


Preparation Name


1. Soluble in alkali, not precipitated Fulvic Acid
by acid

2. Soluble in alkali, precipitated by HC1, Huus Acid
insoluble in alcohol

3. Soluble in alkali, precipitated by HC1, Hrmatomelanic Acid
soluble in alcohol

4. Insoluble in alkali Humus Coal



replaced by the 20th century theories that the composition of humus

material and to some extent the properties, were dependent on the source,

manner of extraction, and length of formation time of the organic sub-

stances.

One of the earliest and perhaps the most persistent theory of the

origin of humus was that it was produced by the action of mineral acids

on sugars. This trend in research prompted many investigators to produce

artificial humus materials in the belief that these were formed along

the same pathways that nature employs in producing humus. Boullay (27)

calculated the chemical constitution of an artificial preparation as

C30H30015. Malaguti (28) suggested that HN03 changes sugars into oxalic

acid whereas HC1 and H2SO4 give humus matter, Under the influence of

acid, the sugar loses water giving first a compound of the structure

C12H28014 which changes to humic acid. The mechanism was as follows:


C12H22011 > C1281206 + 5 H20
Sucrose Humic acid







Stein (29) suggested the formula C24H1809 for the same material

and Berzellius (30) adopted yet another formula, namely C32H12016 for

this preparation.

Berzellius was one of the first workers to isolate humus

material from a natural water. He believed these compounds were also

present in plants and soils. He suggested the name crenic acid for his

preparation which he believed could be oxidized to apocrenic acid by the

action of air. These preparations had formulas of C24H12016 and

C24H6012, respectively. In addition he noted the salts of these acids

with alkalies. Magnesium and iron were soluble in water and could be

transported from soil to water in this manner.

A most extensive study of dark colored compounds extracted from

soil, peat, coal and prepared artificially in the laboratory was made by

Hermann (31). His system of classification only added to the complexity

of the problem but he did note the precipitative effect of heavy metal

acetates on these substances and the fact that although most humus

materials contain nitrogen, many do not.

Around the turn of the century investigators began to doubt the

validity of explaining natural processes of humus formation by results

obtained from artificial materials produced by treatment of sugars and

other complex organic substances. Attention turned towards investigation

of the physical and chemical properties of the natural humus. Thus,

Sestini (32) demonstrated in the humic complex the presence of ethereal

and anhydride substances, as well as hydrorl, alkyl and ketonic groups.

The recognition of the mixed and indefinite nature of the various

humic acids and the turning of attention towards the elucidation of the

complex chemical composition of the natural materials was the most import-

ant advance in the history of research in this area.






Van Bemmelen (33) was one of the first researchers to investi-

gate the colloidal properties of soil organic matter. He believed that

the humus complexes were amorphous and colloidal in nature and origi-

nated from the colloidal substances in plant life.

Hoppe-Seyler (34) investigated both natural and artificial

products and found that treatment with alkali gave dark solutions in

the presence of air, but acid treatment gave the same result even in a

hydrogen atmosphere. Included in his organic compounds were pyrogallol,

pyrocatechuic acid, pyrocatechin, and quinone, which all gave similar

humic materials on oxidation. He noted xylan, or wood gum, in soil in

an unchanged state, and found that lignin complexes take an active part

in the formation of humic acids.

Miklauz (35) pointed out that the compositions of the alkali

and alcohol soluble fractions are affected by the strength of extract-

ment and the extracting temperature; and that the highest carbon content

is usually found in the alcohol soluble fraction.

At this time a new concept appeared which has been called the

colloid-chemical approach, mostly due to the outstanding works of Van

Bemmelen, Baumann and Gully (36). This new trend was based upon two

main ideas. The first was that the acidity of humus is not due to free

humic acids but is associated with the reactions of humus with mineral

salts, bases being adsorbed on the colloidal surfaces of the organic

matter releasing mineral acids. Second, these workers concluded from all

existing data on natural and artificial preparations that humic acids are

not definite chemical compounds but mixtures of plants and animal

residues, partly decomposed, partly conserved because of resistance to

decomposition, and combined in a colloidal state. The possibility that







pure organic compounds are present in the complex was not excluded but

such a compound as humic acid was believed to be entirely absent from

soils.

Baumann and Gully later reported the electrical conductivity

of humic acid to be much less than that of a 0.5 per cent solution of

acetic acid and took this as proof of a lack of free acids.

In 1907-1913, Schreiner and Shorey (37) reported what seemed to

be conclusive proof of the absence of single humic acid structure when

they reported isolating from the humic and the hymatomelanic acid

fractions of soil several pure organic compounds. Among these were

vanillin, salicylic aldehyde, 2-hydroxy- and dihydroxystearic acid,

trithiobenzaldehydes, and various resin acids. The amounts of materials

extracted (50 mg of dihydroxystearic acid from 25 Kg soil) were very

minute and for this reason failed to convince many workers in the field.

Interest in humic acid as a pure substance persisted as the

works of Maillard (38), Samuelly (39) and Gortner (40) postulated humic

acid as a complex of sugars and proteins.

Oden (41) showed by potentiometric titrations that organic acid

ions are present in soils and peats. He attributed the low pH of water

solutions or dispersions of these materials to a release of hydrogen ions

from an acid group rather than as a surface phenomenon. Oden first

reported that these acids were tribasic and later tetrabasic with an

equivalent weight of 339. As had other authors before him, Oden recog-

nized the tendency of these substances to readily yield colloidal dis-

persions in water.

Oden proposed the formula C60H52024(COO014 of molecular weight

1332. He suggested that hymatomelanic acid is a hydrolysis product of







humic acid and forms only during the extraction of humic acid itself.

He reported an equivalent weight of 250 for hymatomelanic acid and noted

that it has a higher carbon content. Because he obtained decreasing

nitrogen content in successive purification procedures he concluded that

the nitrogen in humic and hymatomelanic acids is an impurity. One of

the first authors to quantitatively determine functional groups in these

organic molecules was Eller (42), who concluded from his data on the

relative amounts of carbonyl oxygen in the preparation that the natural

products had little relation to the products obtained by the action of

mineral acids on sugars. Leopold (43) agreed with this theory but found

that, on oxidation of certain phenolic substances, products were obtained

which were similar in physical and chemical properties to the natural

humic acids. Leopold's artificial products prepared by the oxidation of

phenol, quinone, and hydroquinone in alkaline solutions contained 58.05 per

cent carbon of the formula x(C6H4 3)2 with a basic structure of,

0 0
0------~0

lHO H H OOH
Ha'o H H H
0 0
0----------0


The process of oxidation was outlined as follows:


hydroquinone -+ quinone

Phenol oxyquinone -- humic acid

pyrocatechein -+ oxyhydroquinone


His natural humic acid contained 59.6 to 60.2 per cent carbon,

3.2 per cent hydrogen and 1.7 to 2.0 per cent nitrogen, which was






considered as an impurity. The natural products were considered to

arise from the following reaction:


C6H1206 + 0 = C6H403 + 4 H20

The acidity of these materials was attributed to the phenolic hydroxyl

groups.

Fuchs (44) and Leopold (43) demonstrated that humic acids pre-

pared from peat contain both carboxyl and phenolic hydroxyl groups. The

carboxyl groups form salts and the hydroxyl groups are capable of com-

bining with bases. These authors prepared salts of humic acid by shaking

the acid with alcoholic potassium acetate. The base content of the salt

was found to correspond quantitatively to the methoxyl content of the

saturated methylated humic acids. The methoxyl content of the natural

humic acid was found to vary considerably with the source of the material.

Fuchs concluded that humic acid contains three to four hydroxyl groups,

three to four carboxyl groups, one methoxyl and one carbonyl group. The

hydroxyl groups were assumed to be subject to methylation. He considered

humic acids then as "oxy-carbonic acids originating in the decomposition

of dead organic material as dark amorphous substances capable of giving

off hydrogen ions and of forming salts and possessing base-exchange

capacity." The following formulae were suggested for the complexes

isolated from peat.


C49H52010 (OCH3)2 (COOH) (CH3CO) (OH)2

C3H46011 (OCH3) (COOH) (OE)4

C59H41017 (COOH)4 (OH) 3(CH2CO)







Fuchs regarded the presence of nitrogen in different humic acids

as due to the bonding of a protein molecule to a nitrogen free compound

or to the replacement of oxygen by ammonia in the cyclic compound.

Leopold (43) determined the carbonyl number of several humic

acid preparations by reacting them with phenylhydrazine hydrochloride.

Fehling's solution was employed to decompose excess phenylhydrazine,

the amount in excess being measured be determining the amount of evolved

N2. His results varied from 0.22 per cent to 2.31 per cent carbonyl

and were at best inconclusive.

Reports of equivalent weights and molecular weights for these

acids vary widely and depend on the source of the material. Molecular

weight data obtained by Samec and Pirkmaier (45) are shown in Table 2.



TABLE 2

Molecular Weights of Some Humic and Hymatomelanic Acids


Source
Acid Lignite Peat Brown Coal


Humic 1445 1235 1345 gms/mole

Hymatomelanic 855 761 739 gins/mole




Plunguian and Hibbert (46) were the first to consider the possi-

bility of an enol-keto tautomerism in the structure of natural humic acid,

and ascribed the differences in properties of the natural and artificial

preparations to the relative amounts of aromatic and aliphatic hydroxyl

groups in each. The natural material was found to have mainly aromatic







hydroxyl while the artificial had a predominance of aliphatic hydroxyl

groups.

Fuchs (47) studied the relation between humic acids and lignin

and found them to be very similar structurally by testing with various

organic spot reactions. He found the products of nitration to be par-

ticularly similar and concluded they are isonitrosoketohydroxy-carbonic

acids of high molecular weight. According to Fuchs, the difference

between the lignin and humic nitro derivatives is the manner of N-

linkage in each. Molecular weights (determined in (CH3)2CO) ranged from

1410 to 1465 gms/mole for methylated nitro derivatives of lignin and

1535 for the corresponding derivative of humic acid.

Fuchs used bromination as proof that a tetrahydrobenzene ring

is present in natural humic acid and Fuchs and Stengel (48) were success-

ful in identifying various organic acids and nitrophenols by oxidation

of humic acid preparations with HNO3.

Ludmila (49) has presented evidence that humic acids are cyclic

compounds to which are attached carboxyl, hydroxyl, and carboxyl groups

and have an average molecular weight of 1400 gms/mole.

Fromel (50) extracted humic acid from coal and peat with NaOH

and NaF and proved that the Freundlich adsorption isotherm is followed

in the medium concentration range and light absorption in the 2300 to

5000 Ao range follows Beer's law.

Ubaldini (51) determined acid numbers of humic acid by suspending

humus material in ethanol and titration with alcoholic KOH. Total acid

numbers were between 492 and 500 and carboxylic acid numbers between 263

and 267. Methylation with diazomethane showed not all acid groups are







methylated. Carboxyl, phenolic hydroxyl and carbonyl groups were found

in a ratio of 8:7:2.

Scheele and Steinke (52) reported molecular weights of humus

material as high as 9000, and found a slow decomposition takes place

in the presence of alkali giving a water soluble substance that cannot

be precipitated by mineral acid.

Galle and Lodzik (53) investigated the functional groups of

humus material by reacting model substances with MeOl containing HC1o

The acidic MeOH effected an esterification of carboxyl groups and an

esterification of phenolic groups. The authors treated phenylmethanol,

diphenylmethanol and triphenylmethanol in this manner and found a 90.3

per cent methylation. In model compounds containing both carboxyl and

phenolic hydroxyl groups, no distinction was made as to which group was

more completely methylated.

Later Scheele (54) used conductometric titrations to show that

the equivalent weight of natural humic acid is around 200. Molecular

weight determinations by diffusion constants were said to reveal that

the molecular weight of humic acid is twice as large in acid solution

as in basic solution.

Esh and Guha-Sircar (55) prepared humic acids in the following

manner. Soil was washed consecutively with hot alcohol-benzene, water,

2 per cent HCI, and 4 per cent KOH. This filtrate was then precipitated

with HC1. Humic acid prepared in this manner was oxidized with H202

(12 to 15 per cent ) in NaOH at 30-40. Extraction with ethyl acetate

gave acidic and phenolic components. Oxidation with chlorine dioxide

for 20 hours at room temperature yielded butyric and oxalic acids in small

quantities.







Niklas and Genninger (56) also performed extensive studies on

the oxidation products of humic acids. They found ferric hydroxide

"of a given type and concentration" to act as catalysts in the oxidation

of humic acid with H202. No data were given by the authors but they

claim that a gas, a liquid and a solid were obtained in this manner.

The gas was described as a pungent, yellow acid which could be condensed

to a volatile liquid. A light brown wax was reported as being distilled

under pressure at 2500 from the residue.

Forsyth (57) subjected humic materials to ultimate and group

analyses and arrived at the conclusion that although various humic acids

have similar molecular structures, the number and kind of structural

groups depends entirely on the conditions and length of time of forma-

tion of the humus.

Evstigneev and Nikiforova (58) reported spectral data on hydroxyl-

methylfurfural obtained from a decomposition of glucose. This material

exhibits absorption maxima at 282.5 and 228-230 mu Humus substances

from sugars were found to show an initial absorption in the visible

region increasing rapidly in the ultraviolet.

Kukharenko (59) described a semimicro titrimetric determination

of the functional groups in humic preparations. This process involved

heating the humic acids at 1000C with calcium acetate. The number of

milligram equivalents of NaOH necessary to titrate the AcOH form from

the reaction from one gram of sample with calcium acetate equaled the

number of carboxyl groups in the humic molecule. No data were reported.

Hadzi (60) also reported spectral data on humic acids. These

authors found the infra-red absorption band at 3.25 microns (due to

phenolic hydroxyl) was shifted to 3.88 microns when the material was







deuterated (phenolic OD). In addition they found the carbonyl band at
-1
1700 cm and stated that the absence of marked absorption bands in the
-1
900 cm" indicates a lack of acid dimers and proves that the acidic

groups in humic acid are mainly phenolic.

Botteri (61) describes lengthy extractions of meladoin, humic

acid and lignin. This author concludes on the basis of aromatic sub-

stances obtained from alkaline fusion of these materials that there is

no similarity in their molecular structure.

Higuchi and Shibuya (62) obtained a reaction product with humic

acid in NaOH and the diazonium salt of sulfanilic acid. The purified

product was an orange-yellow dye of the following composition:

R(N:NC6H4SO3Na-p)7 where R equals


C106H85016 (CO2Na) (OH) 7 (CO) 4te.


Recently a good deal of attention has been focused on"the use of

model substances for determination of the structural elements in humic

materials.

Ploetz (63) produced 1,4-diquinonylbenzene by adding an aqueous

solution of the tetrazonium salt of p-C64(NH2)2 to a cooled aqueous

or aqueous alcohol solution of quinone. A coupling was noted by the

evolution of nitrogen gas. The author concludes the coupling product

is an intra-molecular quinhydrone complex, and this same structural type

is present in natural humus substance.

Welte et al. (64) used polymers of polyvalent phenols to give X-

Ray amorphous humic materials. Auto-oxidation reactions resulted in the

formation of oxalic acids, acetic acid and CO2 by ring cleavage.







Flaig (65) reported that orthoquinones are intermediate products

in the formation of humic acids. Substituted hydroxyquinones were used

as model substances in investigations on the preliminary stages and

decomposition products of humic acids.

The isolation and identification of organic compound in humus

is described by Wedgwood (66). Polycyclic hydrocarbons, pyrene, fluo-

anthrene, 1,2-benzanthracene, 3,4-benzopyrene, perylene and anthracene

have been detected.

Bremner (67) points out the need for isolative investigation in

soil analysis. Extraction of soil organic matter is usually done with

alkali, however some research has been done using mineral acid as the

extractant. Because, according to this author, acid or base may harm

organic matter extractions)new extraction procedures using neutral salts

were investigated. These results indicate soil organic humus is intimately

tied up with metallic cations. The efficiency of a neutral salt was found

to depend on its ability to remove cations as precipitates or soluble

complexes. The following order in decreasing efficiency was reported:


oxalate> malonate) salicylate> succinate> adipate)

maleate) trichloroacetate> phthalate.


The cations exhibited the following capabilities:


Li> Na) NH4) K) g.


The most effective neutral extractants were found to be those

capable of forming complexes with iron.

According to Bremner there are two distinct theories regarding

the origin of humic materials in soils. The first suggests that hunus






arises from an alteration of plant lignins entering the soil, while the

second suggests they are products synthesized by or formed by autolysis

of soil micro-organisms. Both of these theories have their proponents

among modern soil chemists.

According to Burges (68) all the available evidence suggests that

the humic acid fraction is either a single chemical substance or a group

of very similar substances. He believes humic acid is primarily non-

nitrogenous and nitrogen present in many preparations is due to a

secondary combination with protein or amino acids, the main structural

element being the quinone group.

Steelink et alZ (69) studied alkaline degradation products of

soil humic acid and isolated catechol (1), protocatechuic acid (II) and

resorcinol (III) in a mixture from a KOH fusion.


OH OH OH
OOH (O3OH


COOH

I II III


Felbeck (70) detected about 60 components by gas chromatographic

separation of pyrolysis products of soil organic matter. The pyrolysis

was carried out at a reaction temperature of 3000C and in a H2 stream,

to reduce polymerization of phenolic materials. These pyrolysis products

were ether soluble and were further separated by extraction of the ether

with aqueous alkali. Roughly 20 per cent of the total organic matter

charged to the furnace was accounted for in the product mixture.




21

Steelink at al. (71) have reported the presence of at least two

free radical species in natural humic acid from data obtained in electron

paramagnetic resonance spectra. The authors believed one could be a

semiquinone radical and the other a quinhydrone-type radical.












III. EXPERIMENTAL PROCEDURES


Collection ang Concentration Techniques


The colored water samples secured for this study were collected

in five-gallon polyethylene bags or drums with the exception of the one

local water which was collected in five-gallon polyethylene-lined cans.

Delhez (72) has reported that water stored in polyethylene bottles for

several days may show absorption in the ultraviolet due to a continuous

release of organic matter from the polyethylene. Accordingly, distilled

water samples were stored in each type of container used in the work for

periods up to three months. No spurious absorption was noted in any

sample from 200 to 300 my)

The drums were shipped by truck to the State Boards of Health of

the states of Massachusetts, Connecticut, Virginia, North Carolina and

Wisconsin for sample collection and returned to this laboratory in the

same manner. Numerous samples from the states of Florida and Georgia

were collected by the investigator or members of the laboratory staff.

Because the material responsible for the coloration of water is

present in minute quantities in the stream, concentration of the organic

materials was necessary. Although recovery of the colored materials was

possible by several techniques, only two were used in this study, namely

vacuum distillation and freeze concentration. Gross solid matter was

removed from all water samples prior to concentration by either technique

by filtering through a suitable grade of filter paper.







All concentrated samples were swept with nitrogen and stored

in pyrex containers at 250C + 20 until use.

The distillation apparatus was a laboratory batch-type evapo-

rator manufactured by Precision Scientific Co., Chicago, Ill. The

vacuum source was an ordinary water aspirator which provided an internal

pressure of 79 im hg at 450C. The distillation was carried out between

45 and 500C to lessen any possibility of decomposition or combination

of chemical species by oxidation or other means. In a usual concen-

trating run, 40 liters of water was concentrated to a final volume of

1.0 liter in six hours of continuous operation.

Freeze concentration was carried out in a 16 ft3 Crosbey home

freezer, fitted with a 25-gallon polyethylene drum and an electrical

stirrer. In normal operation, 20 gallons of stream water was concen-

trated to a final volume of 1.0 liter in approximately 110 hours of

continuous operation.



Instrumental Methods


All ultraviolet absorption spectra were obtained on a Carey

Model 14 recording spectrophotometer from 200-800 my at a scanning

speed of 25 I/sec, using 1.0 cm quartz cells.

Infra-red absorption spectra were determined on a Perkin-Elmer

Model 137 Infracord. Solid samples were run by the potassium bromide

technique while liquids were determined on sodium chloride plates.

Fluorescence spectra were obtained with a fluorimeter assembled

at the University of Florida. It consisted of a converted Perkin-Elmer

Model 12 B monochromator fitted with a lithium fluoride prism and a







55 watt Oaram mercury lamp. The detector, an IP 21 photomultiplier

tube, was situated at an angle of 900 from the incident beam and was

connected to an El Dorado Model RH 200 photometer and a Varian Model

G-11-A recorder. A color filter that transmitted the 365, 405 and

436 m mercury lines was placed between the source and the sample.

All color values were determined on a Lumetron Model 450 filter

photometer, a product of Photovolt Corporation, New York, New York. The

instrument was fitted with a blue filter having maximum transmission at

420 ma, and provided a 15 cm light path. The reference standard for

color tests, according to the procedure recommended by the APHA for the

sanitary examination of water, is a solution of 1.246 gias potassium

chloroplatinate and 1.0 gm crystallized cobaltous chloride in 100 ml

concentrated hydrochloric acid, made up to 1.0 liter with distilled

water. This solution has an assigned color value of 500 color units.

A standard color curve was obtained by measuring the optical densities

of appropriate dilutions of the color standard on the Lumetron at

420 m It was necessary to correct the optical density values for

turbidity present in the color samples. The general procedure described

in the Lumetron manual (73) was employed.

All pH values were determined on a Beckman Model G pH meter with

glass and saturated calomel electrodes. The potentiometric titrations

in nonaqueous solvents were carried out with the same titrimeter equipped

with platinum and saturated calomel electrodes.

Light scattering data were obtained with a Series 1965 Brice-

Phoenix Light Scattering Photometer. A General Electric D. C. galvaw

nometer with a sensitivity of 0.0015 ""a/an was used with this instrument.







A Welsbach Type T-23 Laboratory Ozonator was used to supply the

gaseous ozone required in the oxidation experiments. Operated at 7 psig

and 100 VAC, this instrument produced 80 zg/1 03 in an oxygen stream at

a flow rate of 0.03 S.C.F.M.

Electrophoretic mobility data were obtained with a Briggs

microelectrophoresis cell using the technique and accessory apparatus

described by Pilipovich et al. (74) and Black and Smith (75).



Methylation Procedure


The methylating agent used throughout this study was diazomethane

and was prepared in the following manner.

A 125 ml distilling flask was fitted with a condenser and a long

stem dropping funnel. To the flask was added 6 gms of potassium hydroxide

dissolved in 10-ml of water, 35 ml of methanol, 10 ml of sodium-dried

ether and a teflon stirring bar. A solution of 21.5 gis (0.1 mole) of

p-tolylsulfonyl-methylnitrosamide (diazalid) dissolved in 125 ml of ether

was added to the dropping funnel. The distilling flask was stirred in

a water bath at 70-750C, while the contents of the funnel was added at

such a rate that the 125 ml volume was delivered in 15 to 20 minutes.

The distillation was continued until the distillate was colorless. The

usual yield was 2.7 to 2.9 gns of diazomethane.

The diazomethane was standardized by reacting an aliquot with a

weighed excess of pure benzoic acid in absolute ether. The excess

benzoic acid was titrated with standardized sodium hydroxide to a phenol-

phthalein endpoint with rapid stirring.







The procedure for the quantitative methylation of the fulvic

acids and model compounds was as follows: A weighed sample was placed

in an excess of standardized diasomethane solution and allowed to stand

in the cold for 48 hours, or for such time as was necessary to achieve

maximum methylation. At this time, a weighed excess of benzoic acid

was added to react with the remaining diazomethane and the excess

titrated with standard aqueous alkali. From these data it was possible

to calculate the moles of CH2N2 consumed per mole of sample, or, in the

case of the fulvic acids, per unit weight of sample. The latter value

is hereafter referred to as n. That is


S=, moles CH2N2 reacted
100 gms fulvic acid


and if it is assumed that both the carboxylic acid and phenolic hydroxyl

groups are subject to methylation, and that these same groups constitute

the titratable acidity,

= ,nga of acid groups (-COOH and -QOH)
100 gis fulvic acid











IV, EXPERIMENTAL RESULTS

Characteristics of Ea Waters


As mentioned in Chapter III, all water samples were passed

through a roughing filter before concentration. Several qualitative

filters were tried and no measurable effect on color or pH of the

samples was noted. Centrifugation at 30,000 rpm also produced no loss

in color or change in pH. Prolonged storage of colored water has only a

slight effect on the color value, as indicated by the following experi-

ment with water A. Approximately 2.5 liters was swept with nitrogen and

stored in a cold room. The data in Table 3 show only a slight loss in

color over a six-month period.



TABLE 3

Effect of Storage on Color in a Nitrogen Atmosphere


Storage time O.D.b O.D.r
(weeks) (x 10) (x 10) Color


0 6.4 0.75 204

1 6.0 0.70 188

4 6.0 0.70 188

8 6.1 0.70 192

16 6.0 0.71 188

24 6.1 0.68 188








In the presence of oxygen, however, the color was not quite as

stable. Approximately the same volume of water A was swept with oxygen

and stored in the same manner, with the results as listed in Table 4,



TABLE 4

Effect of Storage on Color in an Oxygen Atmosphere


Storage time O,D.i O.D.
(weeks) (x 10) (x 10c Color


0 6.4 0.75 208

1 6.1 0.74 200

2 6.0 0.74 185

4 5.8 0.72 172

8 5.7 0.73 170


All samples used in this study were analyzed within three days

of collection and concentrated within one week.

In general, 25 gallons of each water was obtained. The source

of each water and the designations used to describe them throughout this

report are listed in Table 5.


Chemical Analyses

A partial mineral analysis was performed on each colored water.

These data are presented in Tables 6 and 7.

It is evident that the majority of these colored waters were of

low mineral content with the exception of the two waters from the Florida

Everglades. Several other facts are evident from these analyses. First,








TABLE 5

Sources and Designations of Colored Waters


Water Source


A Creek near Newnan's Lake, Gainesville, Florida

B *Suwannee River at Fargo, Georgia

C Rice Creek near Palatka, Florida

D Lumber Creek near Raleigh, North Carolina

E Juniper Creek near Raleigh, North.Carolina

F Florida Everglades about 20 miles NW of Miami

G Drainage canal near Belle Glade, Florida

H Emerson Brook near Tewksbury, Massachusetts

I Great Dismal Swamp near Norfolk, Virginia

J Unidentified stream near Hartford, Connecticut


*River rises in the Okefenokee Swamp and samples collected in
the swamp near the source.



all of the colored waters contained an appreciable iron content.

Attempts were made at the beginning of this work to obtain a water free

of this metal. Eight Florida waters and one Georgia water, all from

different sources and types of drainage areas, were secured in this en-

deavor and all were found to possess substantial iron contents as seen

in Table 8. In addition, Nordell (76) lists the results shown in Table 9.

Second, there is no relation between iron content and color

value for waters from different sources, although for a given water this

might be expected with seasonal variations.








Third, the chemical oxygen-demand values do not vary linearly

with color values in waters from different sources, although this

relation might hold for a given water. Fourth, biochemical oxygen-

demand values are extremely low in all cases. If the organic substances

responsible for color in water are products of microbiological decom-

position, they are apparently in their final state.

Finally, the amount of dissolved matter in a water sample may

often be estimated by multiplying the specific conductance by an

empirical factor, that varies from 0.55 to 0.90. The data in Tables 6

and 7 show that for soft waters containing organic color this factor is

appreciably higher, varying from 1.47 to 2.21.


Effect of pH on Color

One of the most interesting chemical properties of natural colored

water is the variation of color intensity with pH. This "indicator"

action has been noted by many authors, but again little data on the extent

of this effect has been presented in the literature.

Each colored water sample used in this study was diluted 1:1 with

a Clark and Lub's buffer solution of desired pH value, and the resulting

color determined on a Lumetron. These data for waters A, B, D, E and G

are presented in Figures 1 through 3, and Table 10. In every water used

in this study the variation of color with pH was reversible, but not

linear. In most cases, the slope of the crye in acid solution was con-

siderably less than in basic solution. The data in Table 10 show that

the per cent increase in color value over the pH range employed was quite

variable. However, the greatest increase was obtained in waters with







TABLE 6

Chemical Analyses of Waters A Through E


A B C D E


Total hardness
ppm as CaC03 9.6 14 68 40 30

Ca ppm as CaCO3 -- 8.0 40 2.5 2.0

g1 ppm as CaCO3 --- 6.0 28 1.5 1.0

Alkalinity
ppm as CaCO3 10 0.0 45 2.0 1.3

Chloride ppm 10 16 116 3.0 3.0

Total iron ppm 0.58 1.0 0.31 1.5 0.40

pH 6.75 4.26 7.28 5.85 5.50

Color 240 352 156 108 68

Turbidity ppm trace 12 trace trace trace

Conductivity
Mimhos/cm at
25C 44.0 42.5 256 22.0 17.0

Total residue
at 180oC ppm 76 94 250 40 25

COD ppm from
dichromate 142 83 46 96 28

BOD ppm
5 days at
200C 1.24 -- 0.99 0.60 1.6









TABLE 7

Chemical Analyses of Waters F Through J


F G H I J


Total hardness
ppm as CaCO3 295 505 30 16 8.0

Ca ppm as CaC03 250 365 24 10 6.0

* ppm as CaCO3 45 140 6.0 6.0 2.0

Alkalinity
ppm as CaCO3 230 370 10 4.0 2,0

Chloride ppm 70 183 5.0 4.0 2.0

Total iron ppm 0.15 0.30 0.12 1.05 0.70

pH 7.85 7.95 6.90 5.10 5.00

Color 76 264 70 424 240

Turbidity trace 4.0 trace 10 4.0

Conductivity
Amhos/cm at
250C 770 1490 88 61 23

Total residue
at 1800C ppm 469 990 65 113 61

COD ppm from
dichromate 51 142 24 66 64

BOD ppm
5 days at
200C 1.7 4.7 0.6 1.1 1.4







TABLE 8

Variation of Iron Content in Colored


Water Samples


Water Source Fe (ppm) Color (ppm)


Suwanee River at Santa Fe 1.08 102

Suwanee River at Fargo, Ga. 0.78 316

Fort Pierce 0.50 182

Palm Beach Canal 2.60 88

U. S. 27-South Bay 1.39 65

Pumping Station-U. S. 27 0.36 76

Miami Canal at Bend 1.11 60

Taniami Trail Canal 0.97 45

Newnan Lake tributary 0.58 220




TABLE 9

Variation of Iron Content in Colored Water Samples


Water Source Fe (ppm) Color (ppm)


Great Dismal Swamp
Elizabeth City, N. C. 1.8 1200

Haw River
Burlington, N. C. 1.2 300

Satilla River
Waycross, Ga. 0.11 130

Charles River
Shannuck, R. I. 0.20 70

Black water River
Burdette, Va. 0.17 80

Rancocos Creek
Birmingham, N. J. 2.00 100







TABLE 10

Variation of Color with pH


Color at Color at Per Cent Increase Over
Water pH 2.0 pH 10 Color at pH 2.0


A 160 270 67

B 237 412 74

C 102 182 78

D 71 130 83

E 50 82 64

F 54 82 52

G 190 290 53

H 42 86 105

I 330 440 30

J 150 265 77


lowest original color, and the smallest increase was obtained in waters

with the highest original color.

Because of this pH-color variation, all color values reported in

the course of this work were measured at a constant pH of 8.4.


Particle Size Estimation

Natural color in water has been reported as being in true solution

by some authors (20) and as existing mainly as a colloidal dispersion by

others (77). In all cases, the proposed theories have been backed by

little or no experimental evidence on particle size.








Early in this study it was observed that the smallest com-

mercially available maebrane filter, a 10 my millipore filter, would

retain only 13 per cent of the original color value of water A. Accord-

ingly, two electrodialysis cells were built. They were constructed

from plexiglass, fitted with a copper gauze cathode and a graphite block

anode. These cells were essentially rectangular plastic boxes cut

apart at approximately one-third and two-thirds of their total length

for the insertion of membranes and held together with steel bolts

through plastic cross-braces at either end. The only difference in the

two cells was their capacities, namely one and three liters.

In operation the cell was wired directly to a 125 volt direct

current line in the laboratory. A millimeter was connected in series

with the anode and a voltmeter was tapped across the electrodes. The

electrode compartments were filled with distilled water and changed

hourly.

The maximum amount of color that would dialyze through a fairly

coarse parchment paper membrane was found to be 82 per cent of the

original color after 16 hours of continuous operation. Unfortunately,

the pore size of this coarse parchment could not be determined. However,

when water A was dialyzed through a cellophane membrane with a pore

size of 4.8 mg, 87.5 per cent of the color was retained after 24 hours.

Collodion membranes held 91 per cent of the color in the same length of

time. These data are summarized in Table 11.

These data indicate that most of the color in this water is

colloidal in nature, and while there is a distribution of particle sizes,

most of the particles are between 3.5 and 10 m, in average diameter.







TABLE 11

Electrodialysis and Millipore Filtration of Water A


Maximum Per Cent
Material Pore Size Color Retained


Millipore filter 10 mp. 13.0

Cellophane membrane 4.8 m& 87.5

Collodian membrane 3.5 mul 91.0




All of the waters used in this study were vacuum filtered

through millipore filters of graded pore size down to 10 m,. In

addition, each water was dialyzed for 24 hours using the 4.8 mA cello-

phane membrane. The effects of each on the residual color values of the

waters are presented in Figures 4 and 5.

It is apparent from these figures that a natural distribution

in particle size existed in all of the colored waters used in this study.

Whereas the per cent of original color retained by the 100 m/ filter

varied from 4 to 38 per cent, that retained by the 4.8 mpA membrane

varied only slightly and was approximately 90 per cent. In every

colored water the majority of the particles were between 4.8 and 10 msA

in average diameter. Furthermore, approximately 10 per cent of the color

of each water passed through a membrane with a pore size of 4.8 mny.

The light-scattering properties of turbidity suspensions have

been extensively investigated by Black and annah (78). These authors

characterized various natural and synthetic dispersions by measuring the

intensity of scattered light from 150 to 1370 per unit intensity of

transmitted light at 00. In a given suspension, the ratio of scattered







intensity at 150 per unit scattered intensity at 900 per unit volume

was defined as a scattering ratio. For Latex dispersions, the scatter-

ing ratio increased with increasing particle size as shown in Table 12.



TABLE 12

Scattering Ratios of Latex Suspensions


Particle Size Scattering Ratio


365 mp. 136

264 m) 25

188 m)u 10

88 myA 8




The upper limit of this ratio for substances in true solution is

around 2.0.

If organic color in water is particulate rather than dissolved,

a scattering pattern should be obtained with a sensitive light scatter-

ing photometer. Accordingly, water B was filtered consecutively through

a 100 m,) and a 10 my millipore filter, as this procedure was found to

produce an optimum turbidity-free water by Black and Hannah. Scattering

intensities of a sample prepared in this manner were measured from 15o

to 137 with light sources of 436 and 546 mA Fluorescence filters were

supplied with this instrument for each wavelength and were color filters

that transmitted no source radiation. The results are shown in Figures 6

and 7, and describe the variation of scattering intensity at any angle per

unit transmitted intensity at 00 with the angle 0. It is evident that








these materials show appreciable scattering properties at 546 mk and

that the scattering ratios are not those of dissolved materials.

Scattering at 436 m,& is more intense and relatively constant

from 400 to 1370 This is caused by fluorescence of the organic color

at this wavelength and is shown in Figure 11 to be very intense and

independent of the angle 9. No fluorescence was observed at 546 mg .

It is interesting to note the effect of pH on the scattering properties

of this water. As shown in Figure 6 raising the pH from 5.0 to 12.0

resulted in a decrease in the scattering ratio, whereas the scattered

intensity at any angle per unit transmitted intensity at 00 was higher.

These data indicate that raising the pH of this colored water effected

both a decrease in individual particle size and an increase in the number

of particles present. The fluorescence intensity was higher in alkaline

solution as had been determined previously.


Coagulation Behavor

Black and Willems (77) have shown that coagulation of organic

colored water proceeds by an electrokinetic mechanism. They state,

"When alum or ferric sulfate is added to a colored water whose

alkalinity, naturally present or added, is sufficient to produce a pH

value within the upper portion of the pH range of hydrolysis of the

respective salts, positively charged colloidal picelles .

are formed. These positively charged hydrosols neutralize the negative

charge on the colloidal particles of color resulting in electro-

kinetic coagulation."

Although the two colored waters used in their study were of

different total solids and alkalinity content, they were both local

waters and were obtained from areas with similar forest covers.







As a knowledge of the similarity of this behavior in waters

from widely different sources is desirable, jar test and electro-

phoretic mobility data were obtained on seven of the waters used in

this study. Water A was not run as it was one of the waters used by

Black and Willems and had not changed in quality since their deter-

minations. Waters F and G were not coagulated as normal treatment of

these waters would have been lime softening.

The remaining waters were coagulated with alum, adjusting the

pH when necessary with lime or HC1. Figures 8 and 9 show the results

for waters B and I.

The uniformity of coagulation behavior in these colored waters

is apparent from these figures. The zone of optimum color removal was

accompanied by a zone of charge reversal in every case.

In waters B, C, I, and J, both zones occurred within relatively

narrow pH ranges. Waters D, E, and H had considerably wider pH ranges

of optimum color removal. However, the alum dosages employed with

these waters were substantially higher than would be used in normal plant

practice for waters of such low color values.

It is interesting to note that the width of the pH range of

charge reversal for these waters was roughly proportional to the width

of the pH range of optimum color removal as is shown in Figure 10.

These data indicate that in this respect at least, organic

colored waters from various sources are remarkably similar.

The fact that optimum coagulation of colored waters occurs at

relatively low pH values has been ascribed entirely to the low pH

ranges of hydrolysis of the Al (III) and Fe (III) ions (77). It should

be mentioned here that light scattering data obtained in this study







have shown that the particle size of color particles is affected by pH.

In acid solution the color particles were found to be fewer in number

and larger in size than those in the same water in basic solution.

What effect this phenomenon has on coagulation is undetermined but it

should be considered if the exact mechanism of color coagulation is to

be understood.


Fluorescence Spectra

Shapiro used the fluorescence of organic color fractions to

locate bands after chromatographic separation on paper. However, the

fluorescense emission spectra of organic color have never been recorded.

Since it was observed in this work that the raw stream waters fluoresced

as brilliantly as some of the fractions isolated from them, the spectra

of both were determined. Only the raw water fluorescence spectra will

be presented here.

Figure 11 is the spectrum of the Osram mercury lamp source with

distilled water, while Figures 12 and 13 are the fluorescence spectra

obtained from waters B and F. It can be seen from these figures that

the emitted fluorescence has a maximum intensity at 490 my although

the bands are relatively broad and of appreciable intensity from 450 to

550 mu The spectra of the remaining waters varied only in the intensi-

ties of emission.

The different mineral content of the waters may account for the

fact that the relative intensities of fluorescence emission were not

proportional to the initial color values.

Adjusting the pH value of any of the colored water samples to

higher values would invariably increase the intensity of fluorescent

emission. This effect is shown for water B in Figure 14.







Ultraviolet Spectra

None of the colored waters investigated in this study showed

any absorption in the visible region and all showed only end absorption

in the ultraviolet region. Each water was diluted, if possible, with

pH 8.0 buffer to color values of 200, 100, 50, 30 and 10 color units as

determined on the Lumetron. These samples were scanned on the Cary

Model 14 from 200 to 600 my against the same concentration of pH 8.0

buffer. Two of these spectra are presented in Figures 15 and 16.

However, all were similar in two respects. First, each water showed

only end absorption and second, dilution of each water had no effect on

the nature of the absorption curve other than to decrease the absorbancy

value. A substantial variation was observed in the absorbancy values

of each water at 200, 300 and 350 mp as shown in Figures 17 through 19.

However, at 350 mA all of the waters with the exceptions of waters F

and G had essentially the same absorbancy value at all concentrations of

organic color. The two exceptional samples were both obtained from the

Florida Everglades, in a region of extremely uniform vegetative cover,

and were noticeably more yellow to the eye than the rest of the samples

used in this study. In addition, waters F and G exhibited very weak

fluorescence spectra, in comparison with the other colored waters.

The effect of pH on the ultraviolet spectra of waters B and C

is shown in Figures 20 and 21. Only two waters are shown here, as

identical results were obtained with every water. The absorbance value

in alkaline solution was always uniformly higher than that in acid

solution.

The facts that organic colored waters are invariably colored

yellow-orange to the eye yet show no absorption in the 450-550 mA







region, and that the general shape of the ultraviolet spectra is that

of a scattering rather than an absorption curve suggests that the

color in these waters is due to scattering rather than molecular absorp-

tion of light energy. The difference in absorbance values at the

wavelengths shown in Figures 17 through 19 are probably due to two

factors. First, a natural variation in particle size would account for

a variation in scattering intensity. Second, the color values for each

series were determined on a Lumetron filter photometer which detects

scattered, transmitted and fluorescent radiation. The Cary, however,

eliminates all light of different wavelength from the incident beam

with a second monochromator, and therefore would not detect fluorescent

radiation. Thus, the Cary spectrophotometer was not examining equal

"color" values as determined on the Lumetron.

The variation of absorbancy on the Cary, as well as the Lume-

tron, with pH of the colored water samples may also be explained by

changes in particle size and fluorescence intensity with pH. As men-

tioned in a previous section, both of these effects have been observed.

These data are consistent with the theory that color in water

is of a colloidal nature.



Characteristics of Extracted Organic Material


Nature of the Fractions

Any fractionation scheme that is to be used in a method of

classification of colored waters should include all of the organic matter

present in these waters. The only fractionation method reported in the

literature has been presented by Shapiro (22). However, this author








reported that some organic matter was excluded by this extraction pro-

cedure, which was intended to extract one relatively pure acid fraction.

To determine if this amount was significant, water B was concentrated

and fractionated using Shapiro's techniques.

Shapiro's fractionation scheme was essentially as follows:

The concentrate obtained by vacuum distillation was dried in an oven

at 600C. The solid thus obtained was mixed with 80 per cent ethanol

acidified with HC1. This mixture was taken to dryness, dissolved in

dilute HC1 and extracted with various organic solvents, the most effi-

cient of which was ethyl acetate. Shapiro referred to material obtained

in this manner as the "free acid" fraction. Usual concentrations of

this fraction were reported as 4.16-5.24 mg/1 of raw water.

The solid and liquid materials resulting from the concentration

of 25 gallons of water B were combined and evaporated to dryness in an

oven at 450C over a 48-hour period. This solid residue was treated with

80 per cent ethanol acidified with HC1 to remove mineral materials. The

major portion of the residue dissolved leaving a small but definite

amount of finely dispersed solid (I). The alcohol solution was then

evaporated at 45C leaving a dark colored solid which was mixed with

dilute HC1 and extracted with ethyl acetate thereby separating it into

two more fractions; (II)', acid soluble but insoluble in ethyl acetate

and (II) soluble in ethyl acetate. Fraction (II)' can be further

separated yielding a water-soluble portion (III) and a water-insoluble

fraction (IV).

Thus four distinct fractions were obtained on the basis of

solubility by Shapiro's extraction scheme.








Tables 13 and 14 suamarize the solubilities and relative dis-

tribution of material in these fractions.


Solubility Relationships of


TABLE 13

Fractions Obtained by Shapiro Technique


NaOH NaHCO3 EtOH
5 per 5 per 95 per
Fraction H20 cent cent cent Ethyl Acetate


I + + -

II + + + + +

III + + + +

IV +

+ soluble insoluble + partially soluble




TABLE 14

Weight Distribution in Fractions Obtained by Shapiro Technique


Weight Per
Fraction Cent Weight (from 25 gal)


I 7 0.36 gma

"free acid" II 26 1.34 gns

III 8 0.41 gmn

IV 59 3.06 gms


Thus, the free acid fraction described by Shapiro is seen to con-

tain less than 30 per cent of the organic material present in this natural

water.








All of these fractions dried to a brown to black solid, I and

III being noticeably lighter in color than II and IV. Fractions I,

III, and IV gave no melting points, i.e., they were stable up to 3600C.

However, fraction II softened at 90-1000C. This is a somewhat

narrower melting point range than that reported by Shapiro, but still

indicates a mixture of several constituents.

All of these materials were insoluble in non-polar organic

solvents such as benzene, petroleum ether and dioxane. I, II, and III

were only slightly acid soluble but were completely soluble in basic

solutions of 5 per cent NaOH and 5 per cent NaHCO3. Fraction IV was

only partly soluble in alkalis but was very soluble in dilute HC1.

The presence of enolic structures in all four fractions was

tested by reacting these materials with bromine water and potassium

iodide. All fractions gave a strongly positive test taking up bromine

immediately and yielding a blue color with starch.

A positive test for aromaticity was obtained on all fractions

with the Chloranil test. However, the Le Rosen test for aromaticity

was negative for each fraction. Comparative tests with pyrogallol and

anthraquinone gave the same results. In the presence of sodium car-

bonate, fraction II gave an Immediate red color with the diazonium salt

of sulfanilic acid, indicating a phenolic nature.

Aqueous titration of fraction II with NaOH yielded an equivalent

weight of 224 gms/eq. Shapiro reported an equivalent weight of 228 gms/eq

for the "free acid" fraction isolated from a Connecticut colored water.

The infra-red absorption spectrum of fraction II, the equivalent

of Shapiro's free acid fraction, is shown in Figure 22. This spectrum

was identical with one obtained in Shapiro's laboratory.







In general, the data obtained in this study on fraction II are

consistent with those reported by Shapiro for the free acid fraction.

As mentioned earlier in this report, many authors have pointed

out the vast variety of humus substances of varying elemental com-

positions that may be obtained from soil organic matter, the particular

compositions depending on the type of soil, manner of extractions and

strength of extractants. Since it is believed that these materials do

originate in the soil and in bottom deposits as well, it was decided to

try a manner of separation more similar to that outlined by Oden,

Waksman and Bremner which has since become the classical method of

fractionation of soil organic matter.

Whereas less than 30 per cent of the total organic matter

present was isolated by Shapiro's scheme, the classical separation of

humic material in water would, if applicable, separate the entire organic

content into generally recognized fractions of humus material as shown

in Figure 23.

The concentrate used in this part of the study was prepared by

both concentrating techniques so that a comparison could be made.

Many soils contain ether or chloroform-soluble fats and waxes,

the actual type and amount depending on the nature of the humus and its

abundance in the soil. Accordingly, one concentrate batch obtained by

freezing was extracted in the following manner. One hundred milliliters

of concentrate was shaken in a separatory funnel with 200 ml of sodium-

dried ether. The ether was drawn off, dried over Na2SO4 and allowed to

evaporate. The residue consisted of a greenish-yellow, odorless wax.

Two extractions of ether were sufficient to remove all the wax from this

volume of concentrate. Only 2.0 mg was obtained from 100 ml of concen-

trate,








The wax melted at 55-600C and gave no indication of decomposition

up to 130C. Attempts were made to identify this material in methanol

in a vapor phase chromatography apparatus. However, nothing was recorded

but solvent up to a retention time of 10 minutes.

Solid material was noted in the concentrate obtained by either

the freezing or evaporation methods although there seemed to be less in

freeze concentrates. This solid was at first included in the extraction

process since the entire concentrate was taken to dryness before

extraction.

Since the first step of the new fractionation scheme required a

precipitation of the liquid concentrate with HC1, it was decided to

remove this solid. A 0.45/A millipore filter was sufficient for this

purpose and 3.5 gms of solid was filtered from one concentrate batch

before the HC1 treatment.

When the filtrate was treated with HC1, a much finer solid was

formed. These solids, namely that produced in the concentration process

and that produced by HC1 treatment, were labeled fractions la and Ib

respectively.

Samples of each were subjected to an elemental analysis. The

results, on an ash free basis, are shown in Table 15.

Throughout the literature search for this work, all references

to the separation of the combined humic and hymatomelanic acids by the

addition of HC1 referred to the separation as a precipitation process.

The general appearance of both of these solids, and the manner in which

they were formed suggested that they were not true precipitates but

were formed from a colloidal fraction in the concentrate and were

actually flocculated particles. A solid obtained by adding HC1 to an








extract of soil obtained from the stream area could not be distinguished

from one obtained from the water.



TABLE 15

Ultimate Analyses of Fractions Ia and Ib


Per Cent C Per Cent H Per Cent N


Ia 35.50 5.40 1.86

Ib 43.90 4.75 1.96


Fraction Ia probably resulted from a compression of the electrical

double layer during concentration of the negative charged colloidal

particles in the water to a point where its stability was overcome.

In order to avoid these physical or chemical changes during

concentration, attempts were made to concentrate smaller original volumes

which would still yield sufficient organic material for subsequent

analyses. The concentration of 20 liters of waters having a color of

200 or more to 1.0 liter produced no settleable solid phase. For

waters having a color of less than 200, a 40 to 1.0 concentration ratio

was used with the same result.

All of the colored waters used in this'study were concentrated

in this manner. The weight of organic material in each fraction obtained

from either 20 or 40 liters of raw water is listed in Table 16, while

the weight percentage of each fraction in the raw waters is shown in

Table 17 .











Weight of


TABLE 16

Organic Matter in Fractions


Volume Weight in Fraction, grams
Water Concentrated (1) I II III


1.01

1.58

0.850

0.736

0.560

0.280

1.190

0.499

2.07

1.02


0.020

0.014

0.020

0.007

0.002

0.003

0.010

0.013

0.020

0.008


0.130

0.210

0.100

0.079

0.070

0.036

0.200

0.047

0.310

0.205


- -








TABLE 17

Percentage of Total Organic Matter in Fractions


Fraction Per Cent
Water Fulvic Acid Hymatmcelanic Acid Humic Acid


A 87.0 11.2 1.8

B 87.5 11.6 0.9

C 87.7 10.3 2.0

D 89.6 9.6 0.8

E 88.8 11.1 0.1

F 87.9 11.3 0.8

G 85.0 14.3 0.7

H 89.5 8.4 2.1

I 86.4 12.9 0.7

J 82.8 16.6 0.6








These data show that organic materials in colored waters are

predominately of the fulvic acid type, and that the relative amounts

in each fraction were fairly constant in all the waters examined.

A general relationship was observed between the total organic

matter in each water and the color value of the water. As is shown

in Figure 24, a straight line results if the total organic content,

expressed here as ng/1 of raw water, is plotted against the respective

color values of the waters. Furthermore, samples of humic and hymato-

melanic acids obtained from water A were found to follow Beer's law as

shown in Figure 25. It is apparent from these data that on a unit

weight basis, the fulvic acid fraction is the most color producing, but

only slightly more so than the hymatomelanic acid fraction. Humic acid

produces measurably less color per unit weight than either of the other

two fractions.

Because of the predominance of the fulvic acid fraction in colored

waters, most of the attention during this investigation was directed

toward it.


Evidence of Functional Groups and Structural Characteristics

The fulvic acid material dried to a brown solid which was slightly

hygroscopic and possessed a faint smell of caramel. The hymatomelanic

and humic acid residues, however, possessed no noticeable odor, were

non-hygroscopic and were of different color; the former drying to a light

brown, flaky solid, and the latter drying to a dense, shiny black solid.

All were stable in the solid form at temperatures up to 3600C, although

the fulvic and hymatomelanic acid materials darkened noticeably at

2700C.







All fractions gave colored mixtures in 5 per cent NaOH and 5 per

cent NaHC03 leaving no solid material in the bottom of a test tube.

Fractions I and III were partially dissolved or dispersed in EtOAc, and

completely so in ethylenediamine, pyridine and dimethylformamide. All

of these materials were insoluble in non-polar organic solvents such as

benzene, petroleum ether and hexane.

All enols take up bromine instantly with intermediate formations

of dibromo-enols which form labile oc-bromoketones on elimination of HBr,

Such o(-bromoketones oxidize hydroiodic acid and liberate free iodine.

Fulvic acid samples from every water used in this study gave a positive

reaction to this test yielding a blue color with starch. In most cases,

the reaction and subsequent color development was much slower than the

model compounds aceto-acetic and benzoylacetic esters.

The same organic spot tests used with the Shapiro fractions were

employed to determine the presence of aromaticity to the fulvic acid

fractions of each water. The chloranil test was positive for the fulvic

acid fraction of every water except waters C, F and G, while the Le

Rosen test with sulfuric acid and formaldehyde was negative in all cases.

However, the strongest evidence of aromaticity in this fraction was

obtained using the phosphotungstic-phospho-molybdic acid reduction test.

Aromatic hydroxyl compounds reduce these acids with the production of a

characteristic blue color. A strongly positive reaction was observed for

every fulvic acid sample obtained in this study.

Infra-red spectra of the fulvic acid fraction of each colored

water were obtained and were very similar as shown in Figures 26 and

27 and Table 18.

Carboxylic acids typically show broad absorption bands with a

series of minor peaks or satellite bands over the range of 3.33 to 4.0u .









TABLE 18

Infra-Red Absorption Bands of Fulvic Acid Fractions


Band location, ,


3.30-4.00

3.30-4.00

3.28-4.15

3.31-4.00

3.31-4.00

3.32-4.00

3.35-3.98

3.30-4.15

3.28-4.20

3.30-4.15


5.85

5.82

5.80

5.85

5.85

5.85

5.85

5.85

5.85

5.85


6.15

6.18

6.18

6.15

6.15

6.20

6.15

6.15

6.15

6.15


6.95

6,95

6.95

6.97

6.97

6.90

6.90

6.95

6.95

6.95


7.20

7.25

7.25

7.20

7.15

7.20

7.20

7.20

7.20

7.20


8,00-10.0

7.80-10.0

7.80-10.2

7.95-10.1

7.90-10.1

7.95-10.0

7.95-10.0

7.80-10.0

8.10-10.0

8.30-10.0


Water


__







This is a result of the relative proximity of the associated OH and

normal C-H stretching frequencies. These bands were observed an the

spectra of all fulvic acid samples.

The two strongest absorption bands on every spectrum were

located at 5.8 and 6.15)a The former is the carboxy group stretching

frequency while the latter is the characteristic stretching frequency

of doubly bound carbon atoms in conjugated phenyl groups.

Absorption in the range of 6.9 and 7.3AU in all samples was

probably due to alkane groups. The C-H deformation frequencies of both

the -CH2- and CH3-C- groups are in this region as are the C-H stretching

frequencies of the cis and trans isomers of RlCH=CHR2.

The very broad absorption regions between 8.0 and 10.0/u is

probably due to a combination of the C-C and C-0 stretching frequencies.

The fact that this region is very broad and lacks any more definite

absorption bands indicates that these materials are sructurally complex

and/or contain several carboxylic acids.

Figure 28 is the infra-red spectrum of the methyl ester of the

fulvic acid from water A. Diazomethane was used as the methylating

agent. The presence of the band at 2,9,u indicates that not all of the

aromatic hydroxyl groups were methylated. The increased absorption at

3.4)u must be attributed to the increase in methyl groups. The absorp-

tion at 6.9 is stronger than in the unmethylated fulvic acid while

that at 7.3)A is weaker. In addition several bands were obtained,

namely at 7.95, 9.15, 9.80 and 12.5 u respectively. The band at 12.5)1

may well be the only spectral information obtained to date on the nature

of the substitution of the aromatic nuclei in these molecules. The C-H







out-of-plane deformation frequencies of the three adjacent free hydrogen

atoms in 1,2,3-trisubstituted benzene characteristically occurs at

12.50-12.99L .

The infra-red data thus confirm the assumption that these

materials are hydroxy carboxylic acids, and give strong indications of

unsaturation and aromaticity. The negative charge on particles of

organic color may arise from ionization of the carboql and aromatic

hydroxyl groups.

Fifty milligram samples of the three fractions of water A were

sent to Galbraith Laboratories, Inc., for ultimate analysis. The results,

on an ash free basis, are shown in Table 19.



TABLE 19

Ultimate Analyses of Fractions of Water A


Per Cent
Fraction C H N O


I 41.50 5.72 1.98 50.80

II 29.30 5.94 1.85 62.91

III 49.29 5.11 1.24 44.36


A similar analysis of the fulvic acid fraction of water B showed

40.6 per cent C, 5.45 per cent H, 1.53 per cent N and 49.78 per cent 0.

The fulvic acids of these two waters are indeed very similar in this

respect. In addition, the fulvic acid fraction of water B showed a

methozyl content of 2.64 per cent.







Complexity of Fractions

Some knowledge of the complexity of these fractions was

obtained by the use of paper chromatography. Several solvent systems

were tried but the most useful had the following composition by

volume: pyridine, 50 ml; water, 50 ml; 2-methyl-2-butanol, 50 ml;

(C2H5)2NH, 2 ml.

An excellent elution solvent was composed of methyl ethyl

ketone, formic acid and water of the following volume composition:

methyl ethyl ketone, 37.5 ml; water, 15 ml; formic acid, 7.5 ml.

The following tables summarize the results obtained by chro-

matographing the three fractions of water A on Whatman paper number 3

by the descending technique. The bands were located by irradiation

with ultraviolet light.



TABLE 20

Paper Chromatographic Separation of Fulvic Acid Fraction of Water A


Band No. Rf Range Description


1 0.0 dark brown (origin)

2 0.10-0.16 dark yellow (major)

3 0.36-0.38 yellow (major)

4 0.39-0.63 yellow smear

5 0.67-0.80 yellow mear








TABLE 21

Paper Chromatographic Separation of Humic Acid Fraction of Water A


Band No. Rf Range Description


1 0.0 brown (origin)

2 0.40-0.50 yellow (major)

3 0.44-0.61 yellow

4 0.93-0.98 yellow




TABLE 22

Paper Chromatographic Separation of bymatomelanic Acid
Fraction of Water A


Band No. RE.Rage Description


1 0.0 brown (origin)

2 0.08-0.11 yellow


Several of the bands obtained by chromatographing the fulvic

acid fraction of water A were eluted for infra-red analysis. No signi-

ficant difference in the absorption spectra of any band from that of the

unseparated fulvic acid was noted.


Titration Non-Aqueous Solvents

The acid strength of a phenolic hydroayl group is influenced by

the nature and position of any other substituents which may be present.







Phenol itself, and the naphthols and their homologues are only weakly

acidic in aqueous solution. Similarly, the fulvic acids isolated from

colored waters in this study, which have been shown to contain both

carbowylic acid groups and phenolic hydroxyl groups, exhibited only a

slightly acidic nature in aqueous solution.

In order to increase the acidity of these groups and obtain

more well-defined equivalence points, it was necessary to employ basic

organic solvents. Dimethylformamide, butylamine and ethylenediamine

were tried but the most useful was ethylenediamine. As the solutions

of fulvic acid in this solvent were intensely colored, it was necessary

to titrate them potentiometrically. A platinum-saturated calomel

electrode combination was found to be very suitable when sodium amino-

ethoxide was used as the titrant. Standardization of the sodium amino-

ethoxide was accomplished with benzoic acid. A typical potentiometric

titration curve of this acid is shown in Figure 29.

All the solvents were redistilled before use, as even traces of

water were found to significantly decrease the magnitude of the potential

change at the equivalence point.

Several model compounds were titrated to evaluate the magnitude

of this break for the neutralization of carbonylic and phenolic acid

groups. Figure 30 shows the curve obtained by titrating 53 ag of phenol

with 0.148 N sodium aminoethoxide. A sharp break of 125 millivolts was

obtained and calculation revealed an experimental equivalent weight of

94.0 compared with the theoretical value of 94.11. Two gradual breaks

were observed for resorcinol as shown in Figure 31.

This procedure will distinguish between carboxyl and phenolic

acidity in the same molecule as shown by the ortho and parahydro~ybenzoic








acid curves in Figures 32 and 33. These data show that the o-hydroxy

acid is the stronger acid of the two, but that the phenolic hydrogen

on the p-hydroxy acid is more acidic than that on the o-hydroxy acid.

This difference in acidity may be attributed to hydrogen bonding in

and subsequent stabilization of the salicylate anion.

The fulvic acid fractions of waters A through J were titrated

in this manner and the calculated equivalent weights are shown in

Table 23. The shapes of the titration curves were very similar and each

contained one well-defined equivalence point, as shown in Figures 34 and

35, for the fulvic acid fractions of waters B and I. These data in-

dicate that if different types of acid groups exist in these materials,

they are essentially of the same acidity. The other titration curves

were similar in appearance although the equivalent weights calculated

from them varied significantly. This apparent natural difference in the

number of acid groups per mole was also shown from quantitative methyla-

tion data.

Methylation of several model compounds using the procedure

described in the chapter on Experimental Procedures gave the results

listed in Table 24. The values listed are the maxima that could be

obtained with this methylating agent. It is apparent from these data

that diazomethane will not completely methylate all of the phenolic

hydroxyl groups in these compounds. If the structural relationships

in the fulvic acid molecules are at all similar to the structures of

these model compounds, the validity of the assumptions listed above is

doubtful. Nevertheless, the former definition of the n value is valid

and should be an approximation of the number of acid groups in these ful-

vic acid fractions per unit weight. Table 23 lists the R values







TABLE 23

Equivalent Weight and Quantitative Mathylation Data of
Fulvic Acid Fractions


Equivalent Weight
Water gms/eq


114

138

135

133

122

105


1.34

1.35

2.02

0.48

1.18

0.20

0.41

0.45

1.12

1.40


-- I---









TABLE 24

Quantitative Methylation of Model Compounds


Moles CH2N2 Reacted
Hole Sample
Model
Compound Theoretical Experimental



hydrocinnamic acid 1.00 1.01

p-hydroybenzoic
acid 2.00 2.05

o-hydrozybenzoic
acid 2.00 1.80

hydroquinone 2.00 1.55

resorcinol 2.00 1.62


obtained for the fulvic acid fraction of waters A through J,

the equivalent weights of each as determined by titration in

diamine.


along with

ethylene-


It is interesting to note that the lowest equivalent weight was

obtained for the fulvic acid fraction of water C, and that this same

fraction showed the largest A value. In general, there was an inverse

relationship between the equivalent weight and n values, although this

relationship was not linear. The fact that the n values are not exactly

1.00 per unit equivalent weight indicates, for the 5 values less than

1.0, that some titratable groups are not methylatable, and for n values

greater than one, some methylatable groups are not titratable.







Effects of Gaseous Oxidants

The most common method of color removal consists of coagulation

with alum or ferric salts. However, several gaseous oxidants have been

employed in plant practice to reduce organic color in water. The most

extensively used of these is chlorine.

For over fifteen years, the city of Miami, Florida, has been

treating a hard, colored water by lime softening followed by heavy

chlorination. The softening process itself reduces the color value of

the water from approximately 80 to a color value of 25 to 30. A chlorine

dosage of 12 ppm reduces the remaining color to less than 10. Oxidation

with ozone was tried on a pilot plant scale for a period of six months

with excellent results. However, the power cost for the production of

the ozone was found to be prohibitive. Although ozone is rarely used

for water treatment in this country, over 200 municipal plants throughout

Europe employ ozone for water disinfection, and for the control of tastes

and odors.

Chlorine dioxide is used extensively in water treatment for

taste and odor control, and in the paper manufacturing industry for the

bleaching of wood pulp. This chemical, although more effective than

chlorine as a disinfectant at high pH values, is generally too expensive

for routine water treatment problems.

The effects of chlorine, chlorine dioxide and ozone on the color

values of the waters used in this study were evaluated.

The gases were bubbled through duplicate 200 ml samples of each

colored water concentrate until a constant value was obtained. The

chemical oxygen demand of one sample was then determined according to the







procedures listed in Standard Methods, and the residual color value was

measured on the other sample after it has been appropriately diluted.

It was found that color values determined in this manner agreed with

those resulting from oxidation of the raw water directly. In the cases

of oxidation with chlorine and chlorine dioxide, it was necessary to

destroy excesses of these chemicals with crystals of sodium thiosulfate

before the color values could be determined on the Lumetron. The

results are listed in Table 25.

It is apparent from these data that ozone was more effective

than chlorine dioxide or chlorine for removing organic color, and

chlorine dioxide was more effective than chlorine. Both chlorine and

chlorine dioxide produced their maximmn effects within a contact time of

5 to 10 minutes, whereas ozone generally required anywhere from 20

minutes to 1.0 hour. This is probably due to the relatively low concen-

trations of ozone employed.

In actual practice where these oxidants have been used for color

removal, it has always been assumed that color removal corresponded to

oxidation of the organic matter to carbon dioxide and water. That this

is not true is shown by the results listed in Table 26. These data show

that even in the absence of a color value the waters still exerted a

substantial dichromate chemical oxygen demand. In addition, ether

extraction of each concentrate sample that had been oxidized with chlorine

yielded approximately 10 to 30 mg of a yellow solid which gave the infra-

red spectra shown in Figure 36 and Table 27. These spectra are very

similar and much more definitive than those of the fulvic acid samples.

Sodium fusions of each solid showed that no C-C1 bonds were present.







TABLE 25

Effects of C12, C102 and 03


on Color Removal


Original Color after Oxidation
Water Color Cl2 C102 03


A 240 -- 8 0

B 352 98 14 0

C 156 72 10 0

D 108 30 3 0

E 68 25 0 0

H 70 15 15 3

I 424 22 18 3

J 240 5 0 0




TABLE 26

Effects of C12, C102 and 0 on Chemical Oxygen Demand of
Colored Water Concentrates


COD of COD after Oxidation, ppm
Water Concentrate C12 C102 03


A 2800 -- 750 615

B 1394 1060 80 220

C 996 118 40 152
D 1316 408 120 129

E 554 0 0 82

H 960 780 751 794

I 1320 830 652 742

J 1280 616 514 810




65




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To determine if similar materials were present in a municipal

water that had been treated with chlorine for color removal, 25 gallons

of Miami tap water was concentrated to a final volume of 1.0 liter and

extracted with ether. The extraction yeilded 65 mg of a solid which was

similar in appearance to those previously obtained, and gave the Infra-

red spectrum shown in Figure 37.

The infra-red region from 11.1 to 16.7 .& is a region of strong,

aryl, out-of-plane C-H deformation vibration absorption which are

valuable for identifying substitution patterns in aromatic ring systems.

Generally, monosubstituted ring systems show absorption at 13.0 to 13.7u.

and 14,1 to 14.5A o-disubstituted at 13.0 to 13.6,u, m-disubstituted

at 12.4 to 13.3,A and 11.1 to 11.6A and p-disubstituted at 11.6 to

12.5.u The data in Table 27 show that with the exception of water A,

all spectra gave bands that are characteristic of p-disubstitution.

Since these materials have given every indication of being structurally

complex, it must be concluded that this p-disubstituted ring is only

a fraction of the whole structure and is perhaps an end group. All of

these solids darkened at 285-3100C without melting and gave positive

spot tests for aromatic hydroxyl groups.


Degradative Oxidation Studies

It is noteworthy that despite the many extensive investigations

concerned with the structural nature of soil organic matter, no pure

organic compound has ever been isolated that has been proven to exist in

the natural organic complex. In addition, the small number of pure com-

pounds that have been isolated have always represented yields of 10 per

cent or less.







Similarly, the organic materials extracted from the colored

waters used in this study were of such complexity as to defy identifi-

cation by normal chemical procedures. It was observed that oxidation

of fulvic acid with KMh04, Ce(SO4)2, C102 and H202 produced ether-

soluble oxidation products which gave infra-red spectra similar to the

unoxidized fulvic acids. In all cases, mild oxidation resulted in a

loss of fluorescence. However, some useful information was obtained

by degrading these molecules by peroxy-acid oxidation. These studies

were performed on the fulvic acid concentrates of waters A, C, I and J

as these were the most highly colored.

Periodic acid has a selective oxidizing action on l:2-glycols

and on ot-hydroxy aldehydes and ketones. Compounds of this type are

cleaved at the C-C bond between the adjacent hydroxyl or carbonyl groups

with the formation of aldehydes or acids of shorter chain length.

The procedure used for the oxidation of the fulvic acid con-

centrates with periodic acid was as follows. Two hundred milliliters

of colored water concentrate was adjusted to pH 1.0 with dilute sul-

furic acid and filtered to remove fractions II and III. This volume

was then mixed with a solution of 1.0 gm of sodium paraperiodate in

50 ml of approximately 7 N H2S04 in a 500 ml round bottom flask equipped

with a reflux condenser. The solution was then refluxed at atmospheric

pressure for 1.0 hour, cooled and extracted with ether.

In general, the extraction with ether left an aqueous phase

that was only slightly colored whereas the ether solution was invariably

a deep, orange-red color. Evaporation of the aqueous phases to dryness

at 400C under nitrogen and extraction with EtOH yielded solids which







gave infra-red spectra typical of the unoxidized fulvic acids. These

solids represented approximately 10 to 20 per cent of the total weight

of organic matter present after oxidation.

Chromatographic separation of the materials in the ether extracts

was possible on a 10-inch column packed with adsorptive alumina. In

general, the extracts gave one rather diffuse yellow band near the top

of the column and one relatively narrow yellow band approximately 4

to 6 inches down the column. Neither of these bands fluoresced under

ultraviolet light.

Following elution with 95 per cent EtOH, the lower band dried

to a sweet smelling yellow oil while the upper diffuse bend dried to an

orange solid that decomposed in a Bunsen flame.

For water A, less than 15 ng of the oil was obtained, whereas

135 mg of material was obtained from the diffuse layer. The infra-red

spectra of these materials resulting from the oxidation of water A are

shown in Figures 38 and 39. The absorption data obtained with the yellow

oil indicate that it is not a carboxylic acid and does not contain

hydrosyl groups. The only absorption below 4,0 is due to the presence

of C-H. The shoulder at 3.25 may be due to aromatic hydrogen while

the strong peaks at 3.4 and 3.48 are the assymetric and symmetric C-H

stretching frequencies of the CHl-C group. The presence of the strong

carbonyl band at 5.80/M indicates that fulvic acids may contain these

groups outside of carboxylic acid configurations. The doublet at 6.30

and 6.40)1 is probably phenyl ring absorption. This doublet frequently

occurs in conjugated ring systems. The medium intensity bands at 6.90

to 7.30,A are probably due to C-H deformations of the -CH2- and CH -C

groups respectively. Absorption at 7.90/A is a result of C-C







deformations while that at 8.95pc is due to C-O stretches. The origins

of the bands at 9.40. 9.65 and 10.5 are uncertain while those at 13.5

and 14.25 are indicative of monosubstitution. Absorption in the 3.0-

4.0 A region in Figure 39 indicates that the solid materials obtained

from this oxidation are a mixture of carboxylic acids. The band assign-

ments for all bands up to 8.0 are the sane as for Figure 38. Absorption

at 9.3 and 9.60/p is probably due to -OH deformations. The large

band at 12.5)A may be a result of skeletal vibrations. Attempts at

further separation of these acids met with little success.

The data in Figures 38 and 39 suggest that fulvic acids contain

either oA, -dihydroxy, or an c -hydroxy ketone groupings of the type,

OH 0
I II
R C C -R
I
H

OH OH

R C C R'
I I
H H


where both R and R' contain aromatic nuclei. The facts that the oily

substance showed the presence of a carbonyl group but was not an acid,

and that such small quantities of the oil were obtained, suggest that

the following groupings may also be present in fulvic acids, namely

OH OH OH OH
I I I I
R C C R"' R C C R"
I I I I
R' R" R' H




70


or


OH 0
I II
R C C- R

R'


where R and R' contain monosubstituted phenyl rings.












V. SUMMARY


The physical and chemical properties of ten naturally colored

waters have been investigated and have been found to be relatively in-

dependent of the source of the water.

Angular light scattering data and ultraviolet spectroscopy

have shown that the materials responsible for color in water exist pri-

marily in colloidal suspension in the water. It is suggested that the

color in these waters is due to light scattering rather than molecular

absorption of light energy. The pH value of the waters was found to

affect both the particle size and number of particles in a given sus-

pension. This effect was proposed as an explanation for the variation

of color value with pH.

The particle size of the organic matter in each water was

estimated at constant pH by dialysis and millipore filtration and was

found to be generally less than 10 ma.

The coagulation behaviors of these waters with aluminum sulfate

were compared and it was observed that the pH zone of optimum color

removal was accompanied by a zone of minimum zeta-potential and typically,

of charge reversal, of the coagulated color particles.

The fluorescence spectra of these colored waters were recorded

and revealed in all cases a maximan energy of emission at 490 ma with

the combined excitation wavelengths of 365, 404 and 436 myn.

Organic materials were isolated from these waters and were shown

to be similar to natural soil organic matter. The predominate fraction







was found to be fulvic acid whereas the combined humic and hymatomelanic

acid fractions accounted for less than 10 per cent of the total organic

matter present in each water. For all the waters examined a linear

relationship was observed between the amount of organic matter and the

color value of the water.

Data obtained from solubility relationships, chemical spot tests

and infra-red absorption spectra indicated that fulvic acids are aromatic

polyhydroxy methoxy carboxylic acids. Ionization of the carboxyl and

aromatic hydroxyl groups was offered as an explanation of the negative

charge of color particles.

Equivalent weights of the fulvic acids were determined by

potentiometric titration in ethylenediamine and were found to vary from

89 to 133 gms per equivalent among the waters studied. Quantitative

methylations of the fulvic acids with diazomethane revealed that there

were generally 0.40 methylatable groups present per 100 gas for fulvic

acids of high equivalent weight and 1.35 methylatable groups per 100 gas

for fulvic acids of low equivalent weight.

Ozone was found to be more effective than either chlorine or

chlorine dioxide for the removal of color from these waters. It was

shown that color removal did not correspond to oxidation of the organic

matter to carbon dioxide and water. Chlorine oxidation was found to

produce ether-soluble oxidation products. Infra-red spectra of the

substances suggested that they are polyhydroxy carboxylic acids con-

taining disubstituted phenyl rings. Similar materials were extracted

from a municipal water system that routinely uses chlorine for color

removal.





73


Degradative oxidation of the fulvic acids produced ether-

soluble mixtures of complex carboylic acids from which pure compounds

could not be isolated. However, infra-red spectra of ether-soluble

oxidation products of the fulvic acids with paraperiodic acid indicated

that these acids contain cc, @, dihydroxy, or oc-hydroxy ketone con-

figurations.


































APPENDIX




75












350

A- water G
-- water A

300 -





250





200 -





150





100

0 2.0 4.0 6.0 8.0 10.0

pH


Figure L Effect of pH on Color of Waters A and G














































2.0


4.0


8.0


Figure 2. Effect of pH on Color of Water B


450 --


400 1-


350 I-


300 I-


250 )-


200


10.0
























A- water D
-A- water E


2.0 4.0 6.0 8.0


Figure 3. Effect of pH on Color of Waters D and E


10.0
















100
A water A
--- water B
-0-- water C
80 -0- water D
-&- water E



60





S 40 -





20 -


0


o I I I I I-l-I--I--I
10 20 30 40 50 60 70 80 90 100

Pore Size my


Figure 4. Dialysis and Millipore Filtration of Waters A
Through E


















100 -


--- water F
-- water G
80 water H

--- water I

-A- water J

60 -




40 -
0 40





20 -
0





0 II I I I I
10 20 30 40 50 60 70 80 90 100

Pore Size my


Figure 5. Dialysis and Millipore Filtration of Waters F
Through J














100 --


A Scattering at pH 12.0

--- Scattering at pH 5.0


10 L-


1.0 0 -


20 40 60 80 100 120


Figure 6. Light Scattering Properties of Water B at 546 m)m


w3i
Ex





81









100









10 -






Scattering at pH 12.0

-- Scattering at pH 5.0

1.0 Fluorescence at pH 12.0

-0- Fluorescence at pH 5.0







I I I I I I

0 20 40 60 80 100 120


Figure 7 Light Scatterng Properties of Water B at 36

Figure 7. Light Scattering Properties of Water B at 436 m)




















Alum Dose 120 ppm


5.0

pH


6.0


100




50




0
O


7.0


Effect of pH and Constant Alum Dosage on
Coagulation of Water B


0.50


(+)

0

(-)


0.5




1.0


4.0


3.0



Figure 8.











ALUM DOSE 120 ppm


4.0


5.0


2.0





1.5


6.0


Effect of pH and Constant Alum Dosage on
Coagulation of Water I


160


140


120


100


80


60 1

0

40

20


0


1.0





0.5





0


0.5


Figure 9.




84








3.0 -





2.5




2.0




S 1.5




S 1.0 O


0

0.5 -




0 I I I I I

0 0.5 1.0 1.5 2.0 2.5

pH Width of Charge Reversal


Figure 10. Variation of pH Width of Optimum Color Removal
with pH Width of CMarge Rversal












1.0










0.5


I








350 400 450

/, X ^

Figure 11. Spectrum of Mercury Lamp Source with
Distilled Water


500























1.0


400 450 500 550




Figure 12. Fluorescence Spectrum of Water B












































400 450


500


600


Figure 13. Fluorescence Spectrum of Water F


350






















1.0


400


450


Figure 14. Effect of pH


500 550 600




on Fluorescence Emission
,f Water B












1.3


1.2


1.1


1.0


0.9


As 0.8 200


07


0.6

100
0.


0 .4
OA


0.3 50


0.2 30


0.1L 10


0
200 250 300 350 400 450

X, m)A

Figure 15. Ultraviolet Absorption of Water B as Function of
Color Value

















1.10


1.0

200
0.9 -


0.8


0.7 -


As 0.6

100
0.5 -


0.4


0.3 -
50
0.2
30

0.1 -

010


200 250 300 350 400 450




Figure 16. Ultraviolet Absorption of Water I as Function of
Color Value









1.5
G

1.4


1.3


1.2


1.1


1.0
F
0.9
C

0.8


0.7


As 0.6


0.5


0.4


0.3 -


0.2


0.1


0 I I I 1
0 50 100 150 200

COLOR UNITS


Figure 17. Absorption of Colored Waters at 220 mu


















1.0


0.9 G


0.8


0.7


0.6


As 0.5 E, B, D, J


0.4 -


0.3


0.2


0.1


0
0 I I ----
0 50 100 150 200

COLOR UNITS


Figure 18. Absorption of Colored Waters at 300 mny




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