Effects of Organic 

Solutes on Chemical 

Reactions of Aluminum

GEOLOGICAL SURVEY WATER SUPPLY PAPER 1827-G
Effects of Organic 

Solutes on Chemical 

Reactions of Aluminum

By C. J. LIND and J. D. HEM

CHEMISTRY OF ALUMINUM IN NATURAL WATER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1827-G

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1975
UNITED STATES DEPARTMENT OF THE INTERIOR

GEOLOGICAL SURVEY 

V. E. McKelvey, Director

Library of Congress Cataloging in Publication Data

Lind, C. J.

Effects of organic solutes on chemical reactions of aluminum. (Chemistry of

aluminum in natural water)

(Geological Survey Water-Supply Paper 1827-G)

Bibliography: p. 80-83

Supt. of Docs.: I 19.13:1827-G

1. Organoaluminum compounds. 2. Water chemistry. I. Hem, John David, 1916- 

joint author. II. Series III. Series: United States. Geological Survey. WaterSupply Paper 1827-G.

TC801.U2 [QD412.A4] 553'.7'0973s[551.4'8] 75-619196

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

Washington, D. C. 20402

Stock Number 024-001-02715-1
CONTENTS

	Page 

Abstract ______________________________________________________________ Gl

Introduction __________________________________________________________ 1

Water soluble organic matter _____________________________________________ 3

Concentration in natural waters ______________________________________ 3

General character of organic solutes _________________________________ 3

Influence of simple organic ligands on aqueous aluminum species ____________ 5

Occurrence and form of simple organic ligands _______________________ 5

Mechanisms of organoaluminum interactions ____________________________ 6

Theoretical evaluation of complexing effects _______________________________ 8

Influence of larger organic molecules on the behavior of dissolved aluminum __ 25

Choice of a representative compound __________________________ _____ 25

Preparation of quercetin stock solution ________________________________ 28

Precipitation in aluminum-quercetin solutions __________________________ 28

Aluminum-quercetin complexes _____________________________________ 30

Composition of complexes __________________________________________ 31

Modification of aluminum hydroxide precipitate by traces of quercetin ____ 37

Blockage of crystallization in acidic solutions _______________________ 38

Average size of aluminum hydroxide polymeric ions ______________ 40

Identification of crystalline species ____________________________ 42

Blockage of crystallization in basic solutions _____________________ 45

Development of bayerite________________________________________ 45

Intermediate products ________________________________________ 48

Total effect ______________________________________________ 49

Effects of organic solutes on aluminosilicate precipitation _______________ 50

Modification of the precipitation of aluminosilicates by traces of quercetin__ 57

Crystallinity of precipitates ____________________________________ 59

Solubility of precipitate ____________________________________________ 63

Stream sediment __________________________________________________ 67

Summary _________________________________________________________ 68

Implications of organic complexing in aluminum geochemistry ____________ 69

Anomalous aluminum concentrations in natural water _________________ 73

Effect of filtration ______________________________________________ 75

Examination of analytical data _____________-___--_-_-__-_---_-_ 76

Conclusions _____________________________________________________ 79

References cited _______________________________________________ 80

ILLUSTRATIONS

Page

FIGURES 1-10. Graphs showing:

1. Effects of oxalate complexes on equilibrium solubility of

aluminum in the presence of synthetic halloysite __ __G16

2. Line separating fluoride from oxalate complexing domains 17

in
IV CONTENTS

FIGURES 1-10. Graphs showing Continued Page

3. Effects of malonate complexes on equilibrium solubility of

aluminum in the presence of synthetic halloysite __ G18

4. Line separating fluoride from malonate complexing domains ________________________________________ 19

5. Effects of salicylate complexes on equilibrium solubility of

aluminum in the presence of synthetic halloysite __ _ 20

6. Lines separating fluoride from salicylate complexing domains ____________________________________________ 21

7. Effects of complexes with kojic acid on the equilibrium 

solubility of aluminum in the presence of synthetic 

halloysite ____________________________ 22

8. Line separating fluoride from kojic acid complexing

domains __________________________________________ 23

9. Effects of complexes with catechol on the equilibrium 

solubility of aluminum in the presence of synthetic 

halloysite ___________________________________ 24

10. Lines separating fluoride from catechol complexing domains __________________________________________ 25

11. Sketches showing representations of flavan-3,4-diol structure, 

quercetin molecule, and rearrangement of flavonol to quinoid 

structure in alkaline solution ______________________________ 29

12. Graph showing intensity of fluorescence in solutions having 

various aluminum to quercetin molar ratios at 565 nm, with 

total solute concentration constant ________________________ 32

13. Graph showing intensity of fluorescence in solutions having 

various aluminum to quercetin molar ratios at 490 nm, with 

total solute concentration constant _____________________ 32

14. Graph showing intensity of fluorescence in solutions having 

various aluminum to quercetin molar ratios at 565 nm, with 

quercetin concentration constant _________________________ 33

15. Sketches showing possible structures of aluminum-ouercetin

complexes ____________________________________________ 35

16. Graph showing effect of small amounts of dissolved que^cetin on 

the concentrations of monomeric aluminum species in aged' 

solutions ________________________________________________ 39

17. Graph showing monomeric aluminum concentrations in solutions 

containing 10~5 - 30 mole/1 of quercetin after 310 days aging 

compared with equilibrium values for microcrystalline gibbsite 

and well-crystallized gibbsite or bayerite _________________ 40

18. Diagram showing X-ray diffraction patterns of gibbsite, bayerite,

and nordstrandite ________________________________________ 43

19. Diagram showing X-ray diffraction patterns of solid material 

recovered from solutions Q21D-F and Q22A-C aged 311 and 312 

days _____________________________________________ 44

20-25. Electron micrographs of solids recovered from solution Q22E:

20. Aged 40 days _____________________ _____ 47

21. Aged 57 days (bayerite crystal) ______________________ 48

22. Aged 57 days (general view) _______________________ 49

23. Aged 126 days _________________________ 50

24. Aged 311 days (bayerite crystal and mesh) ____________ 51

25. Aged 311 days (possible nordstrandite crystals along with

bayerite crystals) __________________________________ 52
CONTENTS V

Page 

FIGURE 26. Diagram showing X-ray diffraction patterns for solids recovered

from solution Q22D-F aged 311 days ______________ G53

27. Diagram showing X-ray diffraction patterns for solids recovered

from solution Q23A-C aged 310 days ______________ 54

28. Diagram showing X-ray diffraction patterns for solids recovered

from solutions Q23D and Q23G aged 310 days _________ 55 

29-42. Election micrographs of:

29. Solids recovered from solution Q22E aged 126 days ____ 56

30. Kaolinite from Mesa Alta, N. Mex _____________ 60

31. Keokuk geode kaolinite ___________________________ 61

32. Halloysite from Eureka, Utah __________________ 62

33. Solids recovered from solution Q20G aged 155 days __ 63

34. Solids recovered from solution Q20G aged 155 days

(another view)_--____-________________________ 64

35. Solids recovered from solution Q24D aged 251 days ____ 65

36. Solids recovered from solution Q24C aged 358 days ____ 66

37. Solids recovered from solution Q24G aged 358 days ____ 67

38. Solids recovered from solution Q24C aged 481 days ____ 68

39. Solids recovered from solution Q24E aged 481 days ____ 69

40. Solids recovered from solution Q24G aged 481 days ____ 70_

41. Solids recovered from solution Q26B aged 451 days ____ 71

42. Fine-grained material separated from Colma Creek bed

sediment _________________________________________ 72

TABLES

Page

TABLE 1. Concentrations of aluminum dissolved from bauxite by 0.01 molar 

organic acids compared with that which could theoretically be 

dissolved in their absence at the same pH __________________ G7

2. Structure of complexing ligands and their content of organic

carbon at 1x10 ~ 5 molal _________________________________ 9

3. Chemical equilibria used for solubility calculations.--- ________ 10

4. Activity coefficients, y, at designated ionic strengths __________ 15

5. Potential for complexing by other functional groups ________ 26

6. Effect of quercetin on stability of aluminum species _______ 38

7. *Ks0 values of aged acidic solutions containing quercetin ______ 41

8. Composition during aging of solutions having rn values greater

than 2.75 _______________________________________________ 46

9. Compositions of acidic aluminum-silica-quercetin solutions before

and after aging ___________________________________________ 58

10. Composition of neutral and basic aluminum-silica-quercetin

solutions ________________________________________________ 58

11. Preparation of solutions Q25A-D and Q26A-D ________________ 65

12. Analyses of naturally colored water with high aluminum contents

(Fe=s500/Ag/l) ---_-___________________-___ 77

13. Analyses of naturally colored water with high aluminum contents

(Fe=500-l,000/Ag/l) - -- - - 78
CHEMISTRY OF ALUMINUM IN NATURAL WATER

EFFECTS OF ORGANIC SOLUTES Ol T 

CHEMICAL REACTIONS OF ALUMINUM

By C. J. LIND and J. D. HEM

ABSTRACT

Concentrations of organic matter in the general range of 1-10 milligrams- per litre 

organic carbon are common in natural water, and many naturally occurring organic 

compounds form aluminum complexes. The aluminum concentrations in nesr-neutral 

pH solutions may be 10-100 times higher than the values predicted from solut ility data 

if formation of such organic complexes is ignored. The processes of polymerization of 

aluminum hydroxide and precipitation of gibbsite are inhibited by the presence of the 

organic flavone compound quercetin in concentrations as low'as 10~5'3 mole per litre. 

Quercetin forms a complex, with a probable molar ratio of 1:2 aluminum to ouercetin, 

that has a formation constant (fa) of about 1012 . A complex with a higher 

aluminum-quercetin ratio also was observed, but this material tends to evoT ve into a 

compound of low solubility that removes aluminum from solution.

In the presence of both dissolved aluminum and aqueous silica, low concentrations of 

quercetin improved the yield of crystallized kaolinite and halloysite. Small amounts of 

well-shaped kaolinite and halloysite crystals were identified by electron microscopy in 

solutions with pH's in the range 6.5-8.5 after 155 days aging in one experimer t and 481 

days aging in a repeated experiment. The bulk of the precipitated material was 

amorphous to X-rays, and crystalline material was too small a proportion of the total to 

give identifiable X-ray diffraction peaks. The precipitates had aluminum-silicon ratios 

near 1, and their solubility corresponded to that found by Hem, Roberson, Lind, and 

Polzer (1973) for similar aluminosilicate precipitated in the absence of organic solutes.

The improved yield of crystalline material obtained in the presence of quercetin 

probably is the result of the influence of the organic compound on the aluminum 

hydroxide polymerization process.

Natural water containing color imparted by organic material tends to be higher in 

aluminum than would be predicted by pH, silica concentrations, and solubility data for 

inorganic aluminum species.

INTRODUCTION

Dissolved aluminum is involved in the important processes of rock 

weathering and soil development. Studies of the structure of 

aluminum hydroxide, the factors that affect concentrations of 

dissolved aluminum and its inorganic complexes, and the processes of 

formation of clay minerals have been described in earlier papers in this 

series by Hem and Roberson (1967), Hem (1968), Roberson and Hem

Gi
G2 CHEMISTRY OF ALUMINUM IN NATURAL WATER

(1969), Smith and Hem (1972), and Hem, Roberson, Lind, and Polzer 

(1973).

The effect of organic solutes on the geochemistry of aluminum is 

generally minimal owing to the normally low concentrations of 

organic matter, but under certain conditions, where concentrations 

are higher, it can be significant. Some examples are the strong 

complexes formed by humic acid and aluminum described by Pauli 

(1966), the overall transformation of crop soil by aluminuir pictured by 

Fiskell(1969), and the possible binding of a sizable amount of dissolved 

aluminum to dissolved organic matter, in soil water in the normal pH 

range, demonstrated by Kaurichev, Nozdrunova, and Yevseyeva 

(1969). In fact, Hoefs (1969) went so far as to say "that many geological 

model cases which have been treated as pure inorganic systems are not 

realistic at all, because, especially in sedimentary geology, there are 

no pure systems without organic matter or micro-organisms." Also, 

Schnitzer (1971) stated that there is increasing evidence tl at the most 

important component of "soil solution" is fulvic acid, as it likely affects 

practically all reactions occurring in soils.

The association of aluminum with water-soluble organic matter 

raises questions concerning its possible relationship to biota. 

Aluminum is believed by some investigators quoted by Bowen (1966, p. 

174) to be an essential element for certain plants. The aluminum 

concentration in the sap of plants that grow well in acid foils may be 

very high; according to Chenery (1948) the sap of such plants may have 

a pH as low as 3.9. Because of the increased solubility of aluminum at 

acid pH's, it can be taken more easily into these plants. Th e silky oak, 

Orites excelsa, a plant native to Australia, is reported to exude 

aluminum succinate (Hem, 1970). Excessive amounts of aluminum 

have been shown to be toxic to some plants, such as cottcn seedlings 

(Huck, 1972), potatoes (Lee, 1971a and 1971b), sugarcane (Sumner 

and Meyer, 1971), and alfalfa and barley (Hutchinson and Hunter, 

1970).

This paper considers the influence of dissolved organic matter on the 

processes of solution and precipitation of aluminum and aluminosilicates. The potential for increased concentration of aluminum due to 

the presence of low molecular weight natural organic compounds at 

possible concentrations is shown. A naturally occurring organic 

pigment, quercetin, which is a large molecule having many of the 

properties of the more soluble humic substance, fulvic acid, is used as a 

model compound of higher molecular weight organic matter. This 

compound is shown to react with aluminum, forming complexes and 

precipitates and altering the process of formation of aluminum 

hydroxides by retarding the crystallization and by facilitating the 

formation of an intermediate material prior to establishing a final
EFFECTS OF ORGANIC SOLUTES G3

solid product. It also is shown to aid in the development of crystalline 

kaolinite from aluminum silicate solutions.

WATER SOLUBLE ORGANIC MATTER

The amount and character of organic matter in natural water varies 

widely. These variations are caused by conditions influencing 

biological activity such as source, season, and climate and by the 

chemical, physical, and biological conditions existing between the 

time of initial formation and the time of sampling.

CONCENTRATION IN NATURAL WATERS

A detailed and accurate picture of the forms and amount of dissolved 

organic material in natural water is not presently available. However, 

some indication of the amount can be obtained from the reported 

concentrations in different types of water from various localities. 

Filtered surface and bottom waters in the Rybinsk Reservoir, U.S.S.R., 

had concentrations ranging throughout the year from 4 to 12 mg/1 

(milligrams per litre) of carbon as organic carbon (Skopintsev and 

Bakulina, 1966). (This reservoir impounds 8 rivers and is located 200 

miles (320 kilometres) north of Moscow.) Some relatively orgaric-rich 

surface waters of the United States, studied by Christmrn and 

Ghassemi (1966) and Gjessing and Lee (1967), contained from 7 to 27 

mg/1. Birge and Juday (1934) found from 1.2 to 28.5 mg/1, with a mean 

of 7.7 mg/l,in a study of 529 lakes. Public water supplies from veils in 

the southwestern United States contained 0.7 to 1.6 mg/1, whereas 

those from surface waters in southwestern United States and along the 

Pacific Coast contained 1,1-7.7 mg/1 (Nelson and Lysyj, 1968). 

Kaurichev, Nozdrunova, and Yevseyeva (1969) tested some soils for 

organic carbon during May, July, and October and found that after 

ultrafiltration the soil waters at the surface contained from 24 to 616 

mg/1 and those at roughly 45-60 centimetres in depth contained1 from 2 

to 73 mg/1.

GENERAL CHARACTER OF ORGANIC SOLUTES

The ever-present naturally produced organic substances ir water 

are the materials of principal interest for this discussion. The 

manmade organics generally represent only a small fractior of the 

soluble colloidal organic matter present in natural waters (Chrf stman 

and Minear, 1971), and the current emphasis on pollution abatement 

hopefully will decrease their concentrations still further. According to 

Degens (1965, p. 204) more than 500 organic constituents have been 

extracted from sediments, soils, and natural waters.

Reactions with aluminum can occur in living matter and during its
G4 CHEMISTRY OF ALUMINUM IN NATURAL WATER

decomposition even before the organic matter becomes a dissolved 

species within the water body itself. Thus, the source of dissolved 

organic matter and the concentrations at its sources are of significance. 

Christman and Minear (1971) stated that color in water is derived from 

aqueous extractable material of living woody tissue, from the 

dissolution of decomposition products of decaying wood and of soil 

organic material, or from some combination of these processes. 

Although these may be the more obvious sources, certainly algal 

colonies also make some contribution. Kroes (1972) showed that 

filtrates from two types of green algae were nearly lxlO~5 molar in 

steam-volatile acids and, assuming his suggested minimum of 100 as a 

molecular weight, l.l-1.6xlO~6 molar with regard to the phenolic, 

yellow, water-soluble pigments. (In addition the filtrates contained 

7.5-9.0 mg/1 high molecular weight substances and 0.2-0.5 mg/1 

lipophilic substances also with phenolic character.) Suggestive 

evidence of a flavonol or catechin-type of tannin has been found in 

excretes of certain marine algae (Craigie and McLachlan, 1964).

Many plant pigments are known to be strong complexers of metals. 

One of their basic structures frequently is the benzo-y-pyrone unit. 

This unit can also be found, as kojic acid, among the metabolic products 

of the decomposition of pentoses and hexoses by mold. (Besides being 

found in plants, sugars having the basic structure of pentose and 

hexose are known to be a large part of the less stable fraction of soluble 

organic matter in natural waters.)

Catechol-related structures are common in organic materials such 

as tannin, and a large variety of phenolic units is known to be included 

in the structure of organic compounds from living substances and their 

decomposition products. Under the right conditions, the phenols have 

a capacity to oxidize to quinones, which in turn may form 

semiquinones and dimers. A review of the sedimentary organics and 

their source materials was given by Degens (1965). (Sedimentary 

organics may be precursors to dissolved organics and (or) the result of 

precipitation of them.)

Colored organic solutes derived from soil are frequently described by 

the terms "fulvic" and "humic" acids. The terms are indicative only of 

the methods used for extracting and separating the materials and do 

not denote any definite chemical composition. All such materials are 

known to be complicated mixtures. The three groups of compounds 

considered to be the starting point for the development of humic 

substances are phenolics and quinones, breakdown products of 

proteins (possibly in combination with quinones or phenols), and 

breakdown products of carbohydrates (Duursma, 1965). Black and 

Christman (1963) found indication that fulvic acids, the more water 

soluble of the humic substances, are high molecular weight aromatic
EFFECTS OF ORGANIC SOLUTES G5

polyhydroxy methoxy carboxylic acids. Six years later, Schnitser and 

Skinner (1969) still considered that the nature of free radicals ir humic 

compounds is a matter of speculation and that the most likely 

possibilities are semi-quinone radicals associated with a condensed 

ring system. Schnitzer (1971) noted that 61 percent of the weight of a 

typical fulvic acid molecule and all of the oxygen is in its functional 

groups, largely COOH, OH, and C=0. He found that the major type of 

reaction with metal ions involved both the acidic COOH groups and 

the phenolic groups and that there was none with alcoholic OH groups. 

Rashid (1972) listed the carboxyl, hydroxyl, quinone, other carbonyls, 

and the amino groups on the periphery of the humic acid molecules as 

the functional groups responsible for the higher metal solubility in the 

presence of humic acid as compared with other organic structures.

Lamar and Goerlitz (1966) observed that the predominant organic 

materials in naturally colored surface waters are a mixture of organic 

carboxylic acids or salts of these acids. They concluded that the natural 

organic acids are predominantly nonvolatile polymeric carboxylic 

acids having aromatic and olefinic unsaturation, some of which could 

have been polymerized in aqueous solution. The arcmaticpolyphenolic nature of colored substances in natural water is firmly 

established by oxidative and reductive degradation studies (Christman and Minear, 1971).

Color intensity is one mode of defining the amount of organic matter 

present in natural water. However, the amount indicated to be present 

by this method is only suggestive, as not all species even have color and 

those that do vary in intensity under similar concentrations. Gjessing 

and Lee (1967) found that slightly colored waters contain principally 

the low molecular weight organics, while moderately or highly colored 

waters have a greater amount of high molecular weight species. Some 

of these species were found to have molecular weights as great as 

200,000.

INFLUENCE OF SIMPLE ORGANIC LIGANDS ON AQUEOUS 

ALUMINUM SPECIES

The degree of complexing of aluminum by organic ligands derived 

from the simpler low molecular weight acids can be estimated from 

published data on the stability of such complexes. In evaluating this 

influence it seems appropriate to consider first the circumstances of 

occurrence, the specific identity, and the concentrations of a few substances of this type.

OCCURRENCE AND FORM OF SIMPLE ORGANIC LIGANDS

The nature of these substances can more easily be grasped by an
G6 CHEMISTRY OF ALUMINUM IN NATURAL WATER

examination of the many more common individual organic ligands in 

the source materials, such as plant debris and other material in or on 

soils, than by viewing them at high dulution in large water bodies. 

Bruckert (1970) found that malic, citric, and oxalic acids were the 

principal organic acids in oak, beech, fescue, fir, pine, end callune 

litter; as much as 6,560,1,310, and 630 parts per million, respectively, 

were present in samples of litter dried at 105°C. Other organic acids 

present were succinic, lactic, glucoronic, quinic, vanillic, 

p-hydroxybenzoic, p-coumaric, cafeic, gentisic (2,6-hydroxybenzoic 

acid), gallic, chlorogenic, ferulic, and protocatechic (pyrocat echic) acid. 

Kaurichev, Nozdrunova, and Yevseyeva (1969) identified the simpler 

soluble organic acids in a soil extract as citric, fumaric, g'ycolic, and 

tartaric acids. These soil solutions contained from 2.12 to 69.7 mg/1 

carbon as organic acid, and this amounted to 0.88-65 percent of the 

total carbon present. In another soil solution, p-hydroxyber zoic, vanillic, p-coumaric, and ferulic acids were identified and measured 

(Whitehead, 1964). The concentrations of these acids were from 

0.07xlO~5 to 4.9xlO~5 molar. For example, p-hydroxyb?nzoic acid 

was present in concentrations of 0.67-3.3 mg/1 carbon. The soils tested 

were calcareous loam (with rough grassland vegetation), sand (with 

bracken), clay with flints (with beechwood), and clay loam (with black 

currants and cultivated). Extracts from a mor site (forest humus with a 

layer of largely organic matter abruptly distinct from the mineral soil 

beneath) was found to contain 1.28x 10~5 molar citric acid, 3.62 x 10~5 

molar malic acid, 8.6 x 10~5 molar oxalic acid, 1.0 x 10 ~ 5 molar malonic 

acid, 2.4xlO~5 molar acetic acid, and somewhere between 5.0 and 

6.6xlO~5 molar succinic and lactic acid combined (Brucl'ert, 1970). 

This adds up to a total of about 2.2 x 10 ~4 molar identified acids for this 

soil solution. Stevenson (1972) mentioned values for aliphatic acids of 

even as high as lxlO~3 to 4xlO~3 molar in soil solutions.

Acids recovered from the more volatile 10 percent of the organic 

acids in naturally colored surface waters were identified as acetic, 

adipic, butyric, caproic, fumaric, lactic, maleic, malonic, oxalic, 

phthalic, propionic, succinic, and valeric acids (Lamar and Goerlitz, 

1963,1966). They found the concentrations of these acids to range from 

traces (4 /txg/1 (micrograms per litre)) to 592 /txg/1 (valeric acid). The 

presence of carboxyl and in many cases phenolic hydroxyl groups is 

characteristic of these identified acids, and like fulvic acids, these 

groups play an important part in organometallic reactions.

MECHANISMS OF ORGANOALUMINUM INTERACTION T S

Aluminum minerals are moderately soluble at pH's of 5 or less; the 

resulting A13+ species reacts strongly with the negative organic 

ligands. Aluminum concentrations again increase above pH 7 with the
EFFECTS OF ORGANIC SOLUTES G7

production of A1(OH)4~. This anion, of course, will not be inclined to 

react with the predominantly negatively charged organic ligands. 

Complexing of aluminum by dissolved organic ligands would slso be a 

function of pH in another way. Because most such ligands £ re only 

weakly acidic, potential attachment sites for cations will be occupied 

by H+ and will not be readily available at low pH. Aluminum complexes should be most readily formed with rather strongly acid organic 

species that are considerably dissociated with regard to H+ at pH's less 

than 5. Some types of naturally occurring organic materials showing 

this behavior are oxalic acid and some of the similar aliphs tic carboxylic acids such as malonic and succinic acids. Some aromatic acids 

such as salicylic (ortho-hydroxybenzoic acid) release carboxylic H+ at 

pH's below 5. This may be of particular significance because many 

naturally occurring colored organics in river and lake water an stated 

by various investigators to be polymeric aromatic acids with carboxyl 

and phenol groups. Salicylic acid is a highly simplified monomeric 

structure of this type. Data in Sillen and Martell (1964) show that 

aluminum forms strong complexes with salicylic, oxalic, malonic, and 

kojic acid. Also, Huang and Keller (1972) observed th^t more 

aluminum was brought into solution by the action of citric, tartaric, 

and salicylic acids on bauxite than could be accounted for simply by the 

lowering of the pH by these acids. Table 1 presents the theoretical 

amount of aluminum possible due to pH decrease along with the 

amount reported by Huang and Keller (1972). Complexing effects for 

many of the simple acids with aluminum have apparently r ot been 

evaluated. It might be interesting to note here that Chenery (1948) 

reported that the sap of aluminum-accumulating plants hed a pH 

range from 3.9 to 5.2.

Phenolic sites alone are not favorable for attachment of aluminum 

cations because the phenolic hydrogen is not dissociated significantly 

at pH's where aluminum cations can exist in important proportions. 

However, the phenolic site is not to be completely ignored. For example, when this site is located such that it is favorable for hydrogen

TABLE 1. Concentrations of aluminum dissolved from bauxite by 0.01 molcr organic 

acids compared with that which could theoretically be dissolved in their absence at the 

same pH

Organic acid

Citric _ _

Tartaric

PH

3.14

3.13

3.35

Aluminum in 

presence of 

complexing acid

i mg/1) 1

68.4

81.0

84.4

Al iminum in 

absence of 

Complexing acid

(mg/1)

51.6

55.3

12.1

'Huang and Keller (1972).
G8 CHEMISTRY OF ALUMINUM IN NATURAL WATER

bonding with another functional group, the capacity for tl ^ molecule 

to dissociate is enhanced. This is shown by salicylic acid, 

o-hydroxybenzoic acid, whose dissociation constant is 10 ~3 - 00, as opposed to p-hydroxybenzoic acid with a constant of 10 ~4- 53 and to 

m-hydroxybenzoic acid with a constant of 10~5-85 (Sillen and Martell, 

1964).

The strong binding of metals by chelation (involving more than one 

point of attachment on a single organic molecule) should be acknowledged. Here again salicylic acid and also kojic acid are examples. The 

dissociation of kojic acid is small at neutral pH (Ki = 10 ~ 7- 85 ) and of the 

monovalent salicylate ion is extremely small (K2 = 10~ 13-4), and yet 

the total complexing effect on aluminum is sizeable as shewn later in 

figures 5 through 8. Martell (1971) stated each metal chelate ring 

formed increases the stability constant by a factor of about 100. He said 

that the tight binding of the metal ion by chelation prevents it from 

dissociating to any degree even at high dilutions, while similar complexes with quite high stability constants are well dissociated in dilute 

solutions.

THEORETICAL EVALUATION OF COMPLEXING EFFECTS

Some indications of the potential importance of organic complexing 

in aluminum chemistry is provided by the theoretical evaluation of the 

influence of various organic ligands on the solubility of aluminum at 

equilibrium with solid aluminum hydroxide and the more soluble 

aluminum silicates and the comparison of this influence vdth that of 

fluoride. The selected organic compounds were oxalic, malonic, 

salicylic, and kojic acids, and the dihydroxy phenolic compound 

catechol. The structural formulas of these compounds ard the total 

organic carbon concentrations represented by 10 ~5 molar solutions of 

each are given in table 2. Fluoride concentration is 0.19 mg/1 when 

10~5 molar.

Acidity constants and stability constants for the organoaluminum 

complexes are given with the appropriate mass-law equations in table 

3. Where necessary, the constants were adjusted to zero ioric strength 

using the Davies equation (Butler, 1964). The activity coefficients used 

are listed in table 4. Also given are solubility equilibria a nd equilibrium constants for microcrystalline gibbsite, gibbsite, bayerite, synthetic halloysite, and kaolinite. In additon, the equilibria and stability 

constants needed for computing activities of hydroxide and fluoride 

complexes of aluminum are given in table 3. In all these equilibria, the 

activity of water and the solid phases is assumed to be ecmal to 1.

The effects of each of the five organic compounds on aluminum 

solubility and the amount of the organic complexes as coir oared with 

inorganic ones are shown graphically in figures 1-10. Th° solubility
EFFECTS OF ORGANIC SOLUTES G9

TABLE 2. Structure of complexing ligands and their content 

of organic carbon at 1 x 10 ~5 molal

Name Structure Organic carbon 

(mg/1)

Oxalic acid COOH 

COOH

0.24

Malonic acid

Salicylic acid

COOH

CHa

COOH

.36

.84

Kojic acid. OHv c c

Catechol

.72

.72

plots (figs. 1, 3, 5, 7,9) show the total of aluminum species in solution 

from pH 4.0 to 9.0, assuming equilibrium at standard temperature and 

pressure, with synthetic halloysite. A concentration of 10 ~3 - 48 mole/1 

(mole per litre) or 20 mg/1 as SiO2 of dissolved silicic acid was arsigned, 

but other species in solution were assumed to be low enough tc permit 

using unit activity coefficients. The calculations also carry the inherent assumption that the organic aluminum compounds that might 

be formed are more soluble than halloysite throughout the pF range 

considered.

The solubility calculations involve three principal steps. First, the 

total equilibrium activity of monomeric aluminum ions (A13+ , 

A1(OH)2+ , A1(OH)2 + , and A1(OH)4 -) was calculated from the 

halloysite solubility expression for a selected silica concentrat ion and 

pH value. Second, the total activity of the organoaluminum complexes 

was calculated for that pH and the corresponding calculated Al 

activity, using the summation of the organic species concentration.
TABLE 3. Chemical equilibria used for solubility c

[Activity m moles per litre is represented by [ ] and concentration in moles per litre by ( ). Constants for standard c

Mass-action equations 

Microcrystalline gibbsite solubility produc

(1)

Hydroxide complexes of aluminum

(2)

(3)

(4) [A1(OH)4 ][H+] 4 = 10-2207 [A13 + ] _______________ .__

Gibbsite solubility product

(5) [Al3+ ]=10822 [H+ ] 3 _ __ _______ ____ ___________ ____. ._.___________ __

Bayerite solubility product

(6) [A1(OH) 4 ][H+ ]=1(T 1395 __.......... ._.... ___________________________

Dissociation of water

(7) [H + ][OH-]=l(r 1400 __________________ __ . __________ _ _ ___ ____._.....
Synthetic halloysite solubility

(8) [Al3+ ] 2 [H4SiO4u ] 2 = 10112M [H + ] 6 ___.._.. ___________ _________ ... 

(9) [Al(OH)4 -] 2 [H4Si04 °] 2 [H f ] 2 = 10 32 ' 9lS _ . ..... ________ _ _____ 

Kaolinite solubility

(10) [Al 3 + ] 2 [H4 SiO4°] 2 = 10673 [H+ ] 6 ______ ___ _ _ __ ____ _ ... 

(11) [Al(OH)4 -] 2 [H4Si04°] 2 [H f ] 2 = 10- 37:i4 ____________________ ... ___ _ 

Dissociation of silicic acid 

_______________________________________________________________

(12) [H3Si(V][H+ ]=10-977 [H4SiO4 °] ____________ __ _ 

Oxalic acid dissociation and oxalate complexes (L0 =oxa

__________________________________________________________

 

(13) [HL0 ']=104 ' 14 [H+ ][L02 -]__________._______________________ _______________ 

(14) [H2L0]=10125 [H+ ][HL0 -] __________________________________ _____________ 

(15) (A1(L0) -) = 105 ' 45(A1L0 + )(L02 -) (32°C, 7 = 1) ________ ___________________ 

(16) [A1(L0 ) 2 -]=10609 [A1L0 + ][L02 -] (32°C) __________________________________ _ 

(17) [A1(L0) 2 ~]=1013 [A1 3+ ][L03 ~] 2 (18°C)._______________ ..__________..____._ 

(18) [A1L0 + ]=10691 [A1 3+ ][L02 ^] ___-_____________________ ___ _____________ 
TABLE 3. Chemical equilibria used for solubility calcula

Mass-action equations 

(19) [A1(L0)3 3 ^]=1016-3 [A13+ ][L02 "] 3 (18°C) _ _-. 

(20) [A1(L0) 3 3 -] = 103 - 3 [A1(L0) 2 ~][L02 -] (18°C) ._______....___..__...._. ...__ 

 

Malonic acid dissociation and malonate complexes (Lm2 ~ = mal

 

(21) [HLm ~]=105 '67 [H + ][Lm2 ~] _.._..__._..________.._______.._________....____ 

(22) [HLm ^]=105 - 72 [H+ ][Lm2 -] (35°C)._..._____...__.____....__________________ 

(23) [H2Lj = 10285 [H+ ][HLm -] .__ __________.____...._____...._______________ 

(24) [H2Lm ]=10284 [H+ ][HLm -j (35°C)^_____.__-____________^_________^___.__ 

(25) [Al(Lm ) 33 -]=104- 06 [Al(Lm ) 2 -J[Lm2 -] (35°C,/=?)* ___________________________ 

(26) [Al(Lm) 3 3 -]=101584 [Al 3+ ][Lm2 -] 3 (35°C,/=?)* ____ _._._-_..__ _._.._. 

(27) [Al(Lm) 2 ~]=101178 [Al 3+ ][Lm2 ^] 2 (35°C,/=?)*_..._ _____._.__._._____^__. 

 

Salicylic acid dissociation and salicylate complexes (LS2 ~ = sali

 

(28) [HLs ""]=10134 [H+ ][Ls 2 -]_____^_________________._____^._________________ 

(29) [H2LS ] = 103 [H+ ][HLT] ......._.______....___._........_.__..__......______ 

(30) [AlLs + J = 1014- n [Al 3+ ][Ls2 -](26-28°C) ______________________________________ 
El- I- ECTS OK ORGANIC SOLL'TES G13

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TABLE 3. Chemical equilibria used for solubility calcula

Mass-action equations 

Catechol dissociation and complexes (Lpy2 ~ =dissociated cat

________________ ______ 

_____ ____ ______ 

________ _______________ _ 

(44)

(45)

(46)

Hydrofluoric acid dissociation and fluoride com

(47) [HF°] = 103 - 17 [H f ][F~] _________ _. _ ___ _.. ___________________ _______

(48) [A1F2+ ] = 107 - 01 [A1 3+ ][F~] _____ ______________

(49) [A1F2 + ] = 1012 ' 75 [A13 + ][F-] 2 ......... ___________ _______________________

(50) [A1F3 °]=1017U2 [A1 3+ ][F-] 3 ________________________________________________

(51) [A1F4 -]=1019- 72 [A13 + ][F-] 4 _______ _ . _____ _ _ _ ___________ _________

(52)

(53)

Sulfate complexing

(54) [A1SO4 + ]=103 - 2 [A1 3+ ][SO4 2 -] _ ___ ________

(55) [A1(SO4 ) 2 ^]=1051 [A13+ ][SO4 2 "] 2 ____________________ ________________
EFFECTS OF ORGANIC SOLUTES G15

.med from

TABLE 4. Activity coefficients, y, at designated ionic strengths

[Reference columns: 1, from Butler (1964, p. 435); 2, presumed from data in Butler (1964, p. 435); 3, presu 

data in Garrels and Christ (1965, p. 103); 4, presumed from Butler (1964, p. 438)]

*-> parameter Charge ^ *£**. Reference ^ *£«& Reference

OH~ ____

H+ ___ 

HaSiOj" HL0 -

Lo2 - _

HLm

Lm2 

HLS - _

Ls2 

Lk ~

py " " ^py2 ~ ---- F 

Ala +

A1(OH>"+"

A1(OH>2 +

AKOHl,-

AlLo+ -

A1(LO )2~

AllLgU 3 "

Al(Lm)2 -_

AlLsm3 _

AlLk2 + -

Al(Lkj« +

AlLny+ __

Al(Lpy) 

AKLnylo3 "

AlF2^" AiF2 +:::_

AlF4 -_ 

A1F5 2 - ._

3

1

4

3

5

3

5

6

6

6

6

6

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9

4

4

4

4

4

9

4

9

4

7

9

4

4

9

4

4

4

4

-1

+ 1

-1

-1

-2

-1

-2

-1

-2

-1

-1

-2

-1

+3

+ 2

+ 1

-1

+ 1

-1

-3

-1

-3

+ 1

+ 2

+ 1

+ 1

-1

-3

+2

+ 1

-1

-2

0.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

0.899

.914

.901

.899

.670

.899

.670

.907

.675

.907

.907

.675

.899

.445

.660

.901

.901

.901

.901

.445

.901

.445

.901

.685

.914

.901

.901

.445

.660

.901

.901

.660

1

1

2

2

1

2

1

1

2

1

2

2

2

1

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

2

0.2

T "

.1

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.2

1

1

1

1

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^___

 -

0.72

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____

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80

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.28

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.83

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____

4

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~2

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4

4

4

_^__

 -

The third step involved the summation of the aluminum species. The 

calculations were simplified by expressing the summation ecuations 

in terms of A13+ and H+ activities. The calculation is repeated for as 

many values of pH and total organic concentration as may be needed to 

outline the curves for total organic contents of 0, 10 ~400, and 10~5- 00 

molar.

A second type of graph was prepared for each of the organic 

compounds to compare the organic complexing effect with that of 

fluoride and hydroxide, the two inorganic ions common in natural 

water that complex aluminum most strongly. These graphs (fgs. 2, 4, 

6,8,10) are based on the same relationships that were used to compute 

solubility, with added equilibria and summations for aluminum 

fluoride complex species from table 3. The regression lines on each 

graph represent conditions under which the total complexed 

aluminum is equally divided between fluoride and organic aluminum 

species.

In the preparation of these latter graphs, it was first specified that 

aluminum solubility was controlled by equilibrium with synthetic 

halloysite, at a silica concentration of 20 mg/1 SiO2 (10~3 - 48 molar). A 

summation of organic complex species, in terms of the free ligand and 

uncomplexed aluminum activities (and of pH if necessary), vas then
G16

-1.0

LJJ 

DC 

H- -2.0

in 

LJJ _i

9 -3.0

CHEMISTRY OF ALUMINUM IN NATURAL WATER

-4.0

-5-0

O

O -6.0

O

-7.0

10'4 °° mole/1 oxalic acid 

10~500 mole/I oxalic acid 

oxalic acid absent

1.0

2.0

3.0

4.0

4.0 5.0 6.0 7.0 8.0 9.0

pH

FIGURE 1. Effects of oxalate complexes on equilibrium solubility of alurrinum in the 

presence of synthetic halloysite. Dissolved silicic acid 10~3 48 mole/1 (20 mg/1 as SiO2). 

Position of solubility intercepts for four other solids shown by numbers at right (see 

text).

equated to a summation of fluoride complex species in terms of free 

fluoride and aluminum activities. A pH is specified, along with a 

convenient value for activity of the free organic ligand, and a value for 

[F~] can then be obtained which satisfies the equation. Tl « pH that 

was specified earlier in this manipulation fixes the activity of free 

aluminum. From the activities of the free ligands and the other 

information, it is possible to calculate equilibrium totals of dissolved 

fluoride species and of the organic compound. The graphs use these 

totals as ordinate and abscissa. The final result of the calculation is one 

point on a curve for the pH selected. The whole calculation is repeated
EFFECTS OF ORGANIC SOLUTES G17

-1.0

-2.0

-3.0

-4.0 -

< -5.0 

O

(D 

O

-6.0

-6.0 -5.0 -4.0 -3.0 -2.0 -1.0 

LOG TOTAL DISSOLVED OXALATE, IN MOLES PER LIT^E

FIGURE 2. Line separating fluoride and oxalate complexing domains, pH 4.5-7.0; 

aluminurn concentration controlled by solubility of synthetic halloys; te.

for additional pH and free ligand values to give enough points for 

plotting of curves.

Some of the organic compounds showed no important pH effects 

between pH 4.5 and 6.0, and one regression line can be used to show the 

specified relationship. For other compounds, however, there is a strong 

pH effect, and several curves are required to demonstrate this. Figure 

1 shows a substantial increase in aluminum solubility, especially near 

pH 6, when the higher oxalate concentration is present. Figure 2 shows 

that the oxalate ion is comparable with fluoride in strength of 

complexing with aluminum. Neither species is able to compete very 

effectively with hydroxide in alkaline solutions, however, as shown by 

the convergence of solubility lines in figure 1 above pH 7.

The shape of the solubility traces in figure 1 is the same for any 

aluminum mineral whose solubility products in acid and alkaline 

solutions involve [H+ ]/[Al3+ ] or [H+ ]/[Al(OH)4 -] ratios the same as 

those of microcrystalline gibbsite. This permits a broader ure of the
G18

-1.0

-2.0

CHEMISTRY OF ALUMINUM IN NATURAL WATER

-3.0

-4.0

-5.0

OIO -6.0

O

-7.0

4.0

 10~4 ' 00 mole/1 malonic acid

 10~5 ' 00 mole/I malonic acid

 malonic acid absent

5.0 6.0 7.0 8.0

1.0

2.0

3.0

/ 4.0

9.0

PH

FIGURE 3. Effects of malonate complexes on equilibrium solubility of aluminum in the 

presence of synthetic halloysite. Dissolved silicic acid 10 ~348 mole/1 (20 mg/1 as SiO2 ). 

Position of solubility intercepts for four other solids shown by numbers- at right (see 

text).

diagram. By displacing the curves up or down, equivalent to shifting 

the vertical scale, the same drawing will represent the solubility of 

various forms of gibbsite or of 1:1 clay minerals such as halloysite and 

kaolinite. The amount the scale should be shifted can be determined 

from the difference in the molar free energies of the solids vhen their 

formulas are the same, as in the gibbsite or in the kaolinite species. 

The clay mineral solubilities also are a function of the amount of silica 

in solution, but the appropriate shift of the vertical scale can be readily 

computed for any H4SiO4 activity that is compatible with the assumed 

solid.
EFFECTS OF ORGANIC SOLUTES G19

-1.0

-7.0

-6.0 -5.0 -4.0 -3.0 -2.0 -1.0 

LOG TOTAL DISSOLVED MALONATE, IN MOLES PER LITRE

FIGURE 4. Line separating fluoride from malonate complexing domains, pH 

4.5-7.0. Aluminum concentration controlled by solubility of synthetic halloysite.

The intercepts at the pH 9.0 axis, by the right limb of the solubility 

plot, are for different sets of conditions, as follows:

1. Fresh A1(OH)3 precipitate at alkaline pH, microcrystalline gibbsite 

in acid, no silica.

2. Synthetic halloysite, dissolved silica, 60 mg/1 as SiO2-

3. Well-crystallized bayerite at high pH, gibbsite at low pH, no silica.

4. Well-crystallized kaolinite with AG°=-903 kilocalor; es/mole,

dissolved silica, 20 mg/1 as SiO2-

The wide spread of solubility implied by these conditions (nearly three 

log units) points up very clearly the importance of knowing th e type of 

solid one deals with when attempting to predict aluminum solubility. 

Increasing ionic strength also will increase the apparent solubility
G20

-1.0

CHEMISTRY OF ALUMINUM IN NATURAL WATER

-2.0

-3.0

-4.0

O

§ -6.0

-7.0

4.0

- 10'400 mole/I salicylic acid

- - 10"5 °° mole/I salicylic acid

- salicylic acid absent

1.0

2.0

3.0

4.0

5.0 6.0 7.0

PH

8.0 9.0

FIGURE 5. Effects of salicylate complexes on equilibrium solubility of aluminum in the 

presence of synthetic halloysite. Dissolved silicic acid 10 ~348 mole/1 (20 r^g/l as SiO2). 

Position of solubility intercepts for four other solids shown by numbers at right (see 

text).

of aluminum, but not to the same extent at all pH levels; hence a simple 

shift of scale will not show ionic strength effects accurately. A 

recalculation taking ionic strength effects into account can be made 

using activity coefficients in table 4. The effect is strongest at low pH 

where multivalent ionic species predominate.

Figure 3 shows the solubility of synthetic halloysite in th« presence 

of malonic acid and 10 ~3 - 48 mole/1 of silica. Malonic acid is similar to 

oxalic acid in having two carboxyl groups, but it has one more carbon 

atom than oxalic and is a somewhat weaker acid. The cimplexing 

action of malonic acid toward aluminum is significantly weaker than
EFFECTS OF ORGANIC SOLUTES G21

HI

cc hI] -2.0

CC111

Q_

-3.0

-4.0 -

-5.0

2 -6.0 h-

-7.0

-8.0 -7.0 -6.0 -5.0 -4.0 -3.0 -2.0 -1.0 

LOG TOTAL DISSOLVED SALICYLATE, IN MOLES PER LITRE

FIGURE 6. Lines separating fluoride from salicylate complexing domains at pH 5.00, 

6.00, and 7.00. Aluminum concentration controlled by solubility of synthetic halloysite.

that of oxalic acid. Its action is also weaker than that of fluoride, as 

indicated by figure 4.

Salicylic acid has a benzene ring structure, with one pherolic and 

one carboxylic hydrogen. The influence of this ligand on aluminum 

solubility is indicated in figure 5. This complex species does not 

influence aluminum solubility to a significant extent above about pH 

6.5.

Unlike the oxalate and malonate complexes, the strength of the 

salicylate complex seems to vary with pH in the pH 5-7 range. In figure 

6 the three lines representing equal totals of fluoride and salicylate 

species each have different trajectories.

Kojic acid has a rather substantial complexing effect at pH 7.00, as 

shown by figures 7 and 8. Here the complexing is less influenced by pH 

than in the case of the other organic compounds studies. This acid does 

not possess a carboxyl group like so many of the common low molecular 

weight organic acids. Instead it has a benzo-y-pyrone structure, and 

because of resonance, the carbonyl and adjacent OH groups present a 

chelation site.

The complexing effects of catechol on aluminum are relatively weak
G22

-1.0

LLJ

fE -2.0

O -3.0

-4.0

CHEMISTRY OF ALUMINUM IN NATURAL WATER

O

§ -6-0

-7.0

 10'4 ' 00 mole/l kojic acid

10~5 ' 00 mole/l kojic acid

kojic acid absent

1.0

2.0

3.0

4.0

4.0 5.0 6.0 7.0 8.0 9.0

pH

FIGURE 7. Effects of complexes with kojic acid on the equilibrium solubility of 

aluminum in the presence of synthetic halloysite. Dissolved silicic acid 10 ~3-48 mole/1 

(20 mg/1 as SiO2>. Positions of solubility intercept for four other solids shown by 

numbers at right (see text).

and summarized in figures 9 and 10. Like salicylic acid the strength of 

the catechol complexes is influenced by pH.

As noted earlier, it was assumed for the preparations of figures 1-10 

that organic aluminum species would be more soluble than s luminum 

hydroxide or clay minerals and that the influence of organic 

complexing on aluminum solubility would, therefore, always be in a 

positive direction. Data on actual solubilities of organic compounds of 

aluminum are not readily available. Some of these compounds, 

however, may have solubilities low enough to act as effective controls 

over aluminum concentration in solution, and the complexing action of
EFFECTS OF ORGANIC SOLUTES G23

o.o

-1.0

LU

tr

t -2.0

-3.0

LU

Q 

CC 

§ -4.0

-5.0

-6.0 -

-7.0 -

-8.0

Line represents pH's 5.00, 

6.00, and 7.00

-6.0 -5.0 -4.0 -3.0 -2.0 -1.0 

LOG TOTAL DISSOLVED KOJIC ACID, IN MOLES PER LITRE

FIGURE 8. Line separating fluoride from kojic acid complexing domains for pH 5.00 

7.00. Aluminum concentration controlled by solubility of synthetic halloysite.

such species could depress the solubility of aluminum rather than 

enhance it. The ligands of organic species with this kind of potential 

effect are themselves generally of rather low solubility in water, and 

their effect is commonly considered to be a type of chemisorption 

reaction.
G24

-1.0

CHEMISTRY OF ALUMINUM IN NATURAL WATER

 10"400 mole/1 catechol

catechol absent

-7.0

4.0 5.0 6.0 7.0 8.0

1.0

2.0

3.0

4.0

9.0

PH

FIGURE 9. Effects of complexes with catechol on the equilibrium solubility of 

aluminum in the presence of synthetic halloysite. Dissolved silicic acid 10 ~348 mole/1 

(20 mg/1 as SiC^). Positions of solubility intercept for four other solids shown by 

numbers at right (see text).

Besides these simple organic ligands, many others are possible in 

natural systems owing to substitution by other functional groups. 

These, too, are capable of complexing aluminum. A comparison of the 

stability constants of the simplest aromatic structures of several of 

them with those of salicylic acid will demonstrate the potential 

towards complexing that they offer. (See table 5.) Some of these 

structures might also be found as part of the fulvic and humic acid 

molecular structures.

In summation, it can be said that over much of the pK" range of 

minimum solubility of synthetic halloysite and microcrystalline
EFFECTS OF ORGANIC SOLUTES G25

UJ -1.0

o:

cc

a! -2.0

-3.0

-4.0

-5.0

-6.0

-7.0

pH 5.00 

pH 6.00 

pH 7.00

-7.0 -6.0 -5.0 -4.0 -3.0 -2.0 -1.0 0.0 

LOG TOTAL DISSOLVED CATECHOL, IN MOLES PER LITRE

FIGURE 10. Lines separating fluoride from catechol complexing domains at pH 5.00, 

6.00, and 7.00. Aluminum concentration controlled by solubility of synthetic halloysite.

gibbsite, complexing of aluminum by organic ligands of low molecular 

weight can be at least as important as that due to the strong inorganic 

complexes with fluoride or hydroxide ligands.

INFLUENCE OF LARGER ORGANIC MOLECULE? 

ON THE BEHAVIOR OF DISSOLVED ALUMINUM

Large organic molecules associated with natural coloring material 

in water may approach or even exceed colloidal dimensions. 

Combinations of aluminum with these organic species may change 

aluminum solubility or interfere with or alter mecharisms of 

formation of aluminum hydroxide or aluminosilicates.

CHOICE OF A REPRESENTATIVE COMPOUND

To study the behavior of large organic ligands toward aluminum, a 

series of experiments was performed using a single well-defined 

compound, associated with macromolecular structures, that possesses 

many of the characteristics discussed so far. Besides the lignins,
G26 CHEMISTRY OF ALUMINUM IN NATURAL WATER

O a 

2

 8

«

«E

1

tf

SPo

*?

MO

*~

§p

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EFFECTS OF ORGANIC SOLUTES G27

ffico
G28 CHEMISTRY OF ALUMINUM IN NATURAL WATER

flavan-3,4-diols and flavan-3-ols may be among the sources of colored 

macromolecules, as they undergo condensation polymerization reactions (Christman and Minear, 1971). The biosynthetic formation of a 

tannin polymer is by means of flavanone to flavan-3,4-diol to 

flavan-4-ol (Harborne, 1972). Thus, a structurally related compound 

could be considered. (See figure 11A.) The flavonoids are the largest 

group of naturally occurring phenols (Seikel, 1964); of the some 600 

flavonoids that have been isolated and fully characterized, quercetin, 

kaempferol, and flavadiol are the most common (Harborne, 1972; 

Bate-Smith, 1969). Quercetin, whose complexing capacity is well 

known, was chosen for the study. The structure of quercetin and the 

rearrangement of a flavonol to the quinoid structure in an alkaline 

medium are given in figures 11S and C. (It might be noted that the 

benzo-y-pyrone structure and adjacent OH's on a benzene ring are both 

found in the quercetin molecule. The complexing by each of these 

structural components alone has already been discussed.)

PREPARATION OF QUERCETIN STOCK SOLUTION

A stock solution of purified quercetin was prepared and used at 

various dilutions for tests described in this text. A commercially 

obtained material was purified by the method described by Johnston, 

Stern, and Waiss (1968), by separation in a methanol solution on a 

column of Sephadex LH-20. The elution aliquots to be saved were 

identified by measuring their absorbance by means cf a spectrophotometer. Measurements were made at 370 nm (nanometres), as 

mentioned by Porter and Markham (1970b). The quercetin was then 

crystallized from the methanol, dried in a vacuum dessicator, and 

further dried for 1 hour at 120°C with no loss of weight. (The quercetin 

dihydrate loses its two water molecules at 95-98°C and decc mposes at 

314°C, Strecher and others, 1960.) The purified quercetin is sparingly 

soluble at neutral pH. To prepare a sufficiently concentrated stock 

solution, a finely ground weighed amount was dissolved ir distilled, 

demineralized water by adding dilute sodium hydroxide by drops with 

constant stirring. The pH was kept below 8.65 during the preparation 

of the first stock solution and below 8.50 for later preparations. 

According to Robinson (1967), the oxidation of flavonols in alkaline 

solutions, open to the air, is slow enough that mildly alkaline solutions 

may be used in their preparation. Aeration of the solution wa s kept at a 

minimum. When the solid was completely dissolved, the excess base 

was neutralized with standardized perchloric acid, and the final 

solution thus contained some sodium perchlorate.

PRECIPITATION IN ALUMINUM-QUERCETIN SOLUTIONS

Aluminum may combine chemically with quercetin in the lower pH
EFFECTS OF ORGANIC SOLUTES G29

OH

H OH

B

n*

OH 0

OH

6' 5'

OH 0 OH O OH O"

FIGURE 11. Structural representations: A. Flavan-3,4-diol structure, B. Quercetin 

molecule, C. Rearrangement of flavonol to quinoid structure in alkaline solution (after 

Robinson 1967).

ranges to form a material of low solubility. Also, some type? of large 

organic molecules may form relatively stable colloidal cr larger 

aggregates with aluminum hydroxide. The potential importance of 

such effects was examined by a series of preliminary experiments.

Five equimolar solutions of aluminum and quercetin were prepared. 

The solutions had concentrations of lxlO~5, 1.5xlO~5 , 3.5xlO~5, 

5x 10~5, and 8x 10~5 molar with regard to each constituent and were 

0.01 molar in sodium perchlorate. These solutions were placed in 

125-millilitre boron-free glass bottles, having moisture-tight seal 

caps, and allowed to stand at 25°C. From time to time determinations 

were made of turbidity, aluminum concentration, fluorescence 

emission, and pH. So as not to alter the process of precipitation by 

agitation, a new bottle was opened at the time of each check, leaving 

the remainder of the solution undisturbed. The pH of these solutions 

was near 4.50 at first but increased to values near 5.00 during the 

aging period.

By the end of the 2-month aging period, the solutions with the two 

higher aluminum and quercetin concentrations had developed a 

substantial turbidity and had lost much of their original dissolved 

aluminum. These changes can probably be ascribed to the formation
G30 CHEMISTRY OF ALUMINUM IN NATURAL WATER

and precipitation of some chemical compound of aluminum with 

quercetin. However, the dissolved aluminum remaining at ths end of 2 

months is nearly the same as would be predicted for equilibrium with 

microcrystalline gibbsite. The polymerization process for production of 

gibbsite ordinarily entails a decreasing pH, but the quercetin in these 

solutions might have reacted with H+. No detailed study of the 

precipitates was made, but electron micrographs showed almost no 

gibbsite crystals were present.

Upon excitation by light, derivatives of benzo-y-pyrone emit 

fluorescence at specific wavelengths, depending on the structure 

involved. In the case of quercetin, the addition of aluminum to the 

structure changed the number of emissions from one to two, and they 

were at wavelengths different from that of the pure quercetin 

structure itself. Emission for these two centered at 490 and 565 nm and 

was observed throughout the aging period.

The behavior of the aluminum-quercetin solutions can be briefly 

summarized as follows:

1. All solutions emitted fluorescence at 565 and 490 nm, about in 

proportion to the initial aluminum and quercetin concentrations, 

but the fluorescence intensity decreased with time.

2. The loss of fluorescence was generally associated with development 

of turbidity and precipitation.

3. During the 2-month aging period, a decrease in dissolved 

aluminum and an increase in pH occurred in all the solutions, but 

it was observed after shorter aging periods in the solutions with 

the higher initial concentrations of reactants.

4. The precipitate was mostly an aluminum-quercetin compound, 

perhaps with a small admixture of gibbsite.

ALUMINUM-QUERCETIN COMPLEXES

The nature of the reacting species at the time of preparing the 

solutions can be deduced readily. The pH of the aluminum pe^chlorate 

solution was below 3.00, and that of the sodium quercetin above 5.00 

(5.90 for lxlO-5M to 5.03 for 8xlO~5M); both were 0.01 molar in 

sodium perchlorate. At the low pH of the aluminum solution, the 

aluminum is all in the A1(H2O)6 +3 state, but at the adjusted pH of the 

mixture (4.50), about 16 percent is in the A1(H2O)5(OH)+2 state and 

about 6 percent in the A1(H2O)4(OH)2 +1 state. Titrations of lxlO~3 

molar sodium quercetin solutions with standard acid and 1 ase gave 

indication of one inflection point below pH 7.00, and this was at pH 

5.56. The quercetin anion probably has one negative charge site 

available on most units, and in the mixing process aluminum cations 

or protons might become attached there. To evaluate the complexing 

reaction more closely, the properties of prepared solutions were
EFFECTS OF ORGANIC SOLUTES G31

measured over a period of time to establish the final complexing ratios 

and to determine if these ratios represented relatively stable soluble 

complexes or if the slowly precipitating solid that formei had a 

well-defined composition.

Several conclusions can be made from Porter and Markham (1970a, 

b, c). The amount of complexing and the ratio of aluminum to quercetin 

complexed depends on conditions such as pH, nature of the solvent, and 

the structure and concentrations of the reacting substances. From 

figure 11 it is evident the quercetin molecule has three places where 

oxygen and (or) phenolic OH are located at adjacent sites. Sites of this 

type might constitute favorable locations for aluminum ion attachment. In other words, aluminum might be attached at the 3'-4' 

position, the 5-hydroxy-4-keto-position, and the 3- 

hydroxy-4-keto-position, the latter probably being the most preferred 

site.

COMPOSITION OF COMPLEXES

Two approaches were tried in an attempt to identify the ratios of 

aluminum to quercetin in the complexes and precipitates. The first was 

the method of Job (1928), where the total of the two constituents is held 

constant in a series of solutions but the metal to ligand ratios differ, 

and the second method was a modification involving a series of 

solutions in which one constituent was held constant and the other 

varied. Fluorescence was measured at 565 and 490 nm for each 

solution.

The sum of the concentrations studied using Job's method was 

2xlO~5 molar in aluminum plus quercetin. The fluorescence 

emissions at 565 and 490 nm were measured after 1, 6, and 19 days of 

aging. The mole ratios of aluminum to quercetin examined were 0:10, 

1:9,1:4, 3:7,1:2, 2:3,1:1, 3:2, 2:1, 7:3,4:1,9:1, and 10:0. The maximum 

in fluorescence should be observed at the ratio corresponding to the 

composition of the complex causing the fluorescence. The pF's of the 

solutions increased somewhat during the experiment, but the 

turbidity even through the 19th day was in the range of pure water, 

indicating no significant precipitation. A small amount of fluorescence 

is given off by quercetin itself at the wavelengths observed for the 

fluorescence of complexes. Consequently, to assist in interpreting the 

fluorescence due to complexes, the readings were correctei for the 

effect of quercetin according to the method of Vosburgh and Cooper 

(1941). The final results are shown in figures 12 and 13. Fluorescence 

intensities are represented in arbitrary units on a linear scale of 0-100 

and are not directly convertible to ion concentrations. However, the 

conditions of measurement were the same for all samples and the 

intensity values, therefore, indicate relative concentrations of 

complexes and are a valid measure of changes or differences. The
G32 CHEMISTRY OF ALUMINUM IN NATURAL WATER

50

w 40 C/3 m

- 01 

?S 30

01 ID

§<10

u_

o first day

x nineteenth day

0

0/10 1/9 1/4 3/7 2/3 1/1 3/2 7/3 4/1

MOLAR RATIO ALUMINUM/QUERCETIN

9/1 10/0

FIGURE 12. Intensity of fluorescence at 565 nm in solutions having various aluminum 

to quercetin molar ratios. Total solute concentration constant.

50

- c/340 - 

w oi

?! oc

0 

0/10 1/9 1/4 3/7 2/3 1/1 3/2 7/3 4/1 

MOLAR RATIO ALUMINUM/QUERCETIN

9/1 10/0

FIGURE 13. Intensity of fluorescence at 490 nm in solutions having various aluminum 

to quercetin molar ratios after 1 day. Total solute concentration corstant.

565-nm curve had a maximum near an aluminum to quercetin ratio of 

2:3 on the 1st and 6th day and became sharper, centering around 1:2 by 

the 19th day. The intensity of fluorescence did not change greatly 

during the 19 days. The 490-nm curve had a maximum near 7:3 when 

the solutions were aged 1 day, 3:2 when aged 6 days, and became 

sharper at 3:2 after 19 days aging. The intensity of this fluorescence 

was noticeably greater at the first measurement than it was in later 

measurements, and the later two sets of data points are not shown in 

figure 13.
EFFECTS OF ORGANIC SOLUTES G33

The attempt at the method where one component is held constant 

was done with the quercetin concentration held at 1 x 10-5 molar. The 

soluble complexes were measured immediately after mixing and after 

6 and 14 days of aging. The ratios of aluminum to quercetin measured 

were 0:10, 1:4,1:3, 2:5, 9:20,1:2, 3:5, 2:3, 3:4,1:1, between 6:5 and 5:4, 

4:3, 3:2, about 5:3, 2:1, 9:4, and 5:2. A plot of fluorescence intensity 

versus the aluminum-quercetin mole ratio should give a curve with a 

change in slope at the composition of the complex causing the 

fluorescence. The pH's of the solutions immediately after mixing 

ranged from 4.49 to 4.51, on the 6th day from 4.49 to 4.60, and on the 

14th day 4.56 to 4.75. The turbidity of all the solutions at each of these 

times was well within the range of pure water. Because the quercetin 

concentration was held constant, no correction was applied for its 

presence. The 565-nm plot, figure 14, had a smooth curve when aged 

only 1 day. By the 6th day, however, the plot had a sharp break at the 

1:2 ratio. Measurements made at 14 days are not shown and gave more 

scatter; however, they also indicated a 1:2 ratio. The fluorescence at 

this wavelength did show some decrease with time, particularly at the 

higher ratios. The 490-nm readings did not give interpretable results. 

The intensity of the fluorescence at the 490-nm wavelength decreased 

substantially during the 14-day aging period. In conclusion, it can be 

said that quite likely the 565-nm fluorescence peak represents a

0 1/4 1/2 3/4 1/1 5/4 3/2 7/4 

MOLAR RATIO ALUMINUM/QUERCETIN

2/1 5/2

FIGURE 14. Intensity of fluorescence at 565 nm in solutions havin? various 

aluminum to quercetiu molar ratios. Quercetin concentration constant.
G34 CHEMISTRY OF ALUMINUM IN NATURAL WATER

complex with an aluminum to quercetin ratio of 1:2 and the 490-nm 

peak a complex, richer in aluminum, with a 3:2 ratio.

Observations during these tests indicate that the high aluminum to 

quercetin ratio complex, represented by the 490-nm fluorescence, is 

the less stable and may be a precursor of solid material. Tb*» basis for 

this statement is that this complex is the one that loses fluorescence 

when the turbidity becomes greater than that of pure water and also 

that the peak is decreased when these test solutions are filtered. The 

loss of fluorescence may be explained by polymerizatic n to macromolecules and subsequently to particles large enough to precipitate.

Even though it is recognized that Job's method is strictly applicable 

only to systems simpler than the one investigated here and cannot be 

used to determine accurate stability constants for these complexes, the 

qualitative indication that the method gives of complex composition is 

probably useful.

A second series of tests was made using Job's method with somewhat 

more concentrated solutions. The sum of the concentrations of the 

initial solutions was 4xlO~5 molar in aluminum plus quercetin. 

Before this experiment was completed, a noticeable amount of 

precipitate had formed in most of the solutions, and the fluorescence 

had largely disappeared. However, the results again indicated molar 

ratios of 1:2 and 3:2 aluminum to quercetin for the two complexes. 

From consideration of the results of Porter and Markham (1970c), who 

studied quercetin complexing and dissociation in a nonaqueous 

system, one can picture the more soluble 1:2 complex with quercetin in 

water at pH 4.5 as the structure represented by figure ISA. The second 

aqueous complex ratio, 3:2, might be represented by figure 15B. The 

complexes appear to be unstable or at best sparingly soluble, however, 

and one might question the significance of any effort to specify exact 

values of thermodynamic stability constants.The amount of ionization 

of quercetin is small in the final aqueous solutions. Taking the pKa of 

5.56±0.08 (fCa is the quercetin acid dissociation constant) and a final 

pH of 4.55 and assuming the total concentrations of quercetin as 

lxlO~5 and 2xlO~5 molar, the calculated concentration^ of H4L~ 

would be IxlO""6 and 2xlO~6 molar, respectively, or only about 10 

percent of the total quercetin. (The activity coefficient used, 0.912, was 

that for dibenzyl acetic acid (Butler, 1964), and the 0.01 ionic strength 

of the solution was considered.) However, there may be otl ^r factors 

influencing quercetin complex formation besides the previously 

discussed acid dissociation. Martell (1971) gave a rather thorough 

description of chelation and references for further pursuit of the 

subject. He indicated that the combination of a ligand with a metal ion
EFFECTS OF ORGANIC SOLUTES G35

A

HO

B

Al

FIGURE 15. Possible structures of aluminum-quercetin complexes. A, Aluminum to 

quercetin 1:2. B, Aluminum to quercetin 3:2.
G36 CHEMISTRY OF ALUMINUM IN NATURAL WATER

involves the displacement by the ligand of one of the water molecules 

from the fully coordinated metal ion hydration shell. The nature and 

size of the coordinated donor groups of the solvent and attached ligands 

determine the coordination numbers of the metal ions. If the number of 

donor groups matches the coordination number of the metal, the 

chelate or complex is stable, but if not, and the relatively active "aquo" 

form exists, several types of reactions can occur. Strong polarization of 

the donor atoms and adjacent parts of the organic molecules results 

when the metal and ligand combine, shifting electrons toward the 

metal ion. This change of electron distribution can lead to eolvolysis, 

rearrangement of the organic molecule, and even decomposition 

reactions. Thus, an accurate understanding of the final precipitation 

process and probably even the complex development requires more 

information than we have been able to obtain in our experiments.

Attachment of the aluminum ions at the keto and hydroxy] positions 

is somewhat similar to that of complexing by kojic acid and at the 

phenolic hydroxide positions to that of complexing by catechol, 

represented in figures 7 and 9 respectively. The low solubility of the 

product formed with quercetin, however, implies that the influence of 

quercetin on aluminum behavior might differ from the solubility 

enhancement predicted for the other organic ligands. It seems likely, 

for example, that a concentration of 10~4 molar for quercetin cannot be 

reached at pH 5 when the concentration of aluminum is of a rear order 

of magnitude.

Figures 12 and 14 can be used as a basis for estimating the strength 

of the 1:2 aluminum-quercetin complex which fluoresces at 565 nm. At 

the maximum fluorescence reached, the quercetin may, as a first 

approximation, be assigned entirely to the complex A1Q2 (where Q 

stands for quercetin). From the stoichiometry a concentration of 

uncomplexed aluminum can be assigned. The concentration of 

uncomplexed quercetin, however, can only be said to be small 

compared with the total amount and cannot be assigned any specific 

value. Nevertheless, one may now select a point lower on the 

fluorescence curve that represents an intensity less than the 

maximum by some specific factor and calculate from this the amount of 

quercetin complexed at that point. The amounts of free aluminum and 

quercetin can be approximately calculated from stoichiometry applied 

to the aluminum quercetin ratio indicated on the X axis of the graph.

Several repetitions of this calculation give values for /? > for the 

formation of A1Q2,

Al+3+2Q-=AlQ2+,

of about 1012 , which is applicable at the pH of the test solutions 

(4.5-5.5). This approach ignores possible effects from the other
EFFECTS OF ORGANIC SOLUTES G37

aluminum-quercetin complex that fluoresces at 490 nm, but this 

complex appears to be too unstable to treat as an equilibrium species. 

The method of estimation is subject to many uncertainties, and the 

value of /% for the quercetin complex given here is only a rough 

approximation.

Although the effect of quercetin on equilibrium solubility of 

aluminum near pH 5.00 probably is similar in magnitude to the effect of 

malonate shown in figure 3, there are important differences. Low 

solubility of quercetin and the tendency to form higher complexes that 

are even less soluble are important. Also, the chemistry of the 

quercetin complexes is not fully evaluated. The influence of pH, for 

example, is uncertain, and the tendency for the pH of solutions 

containing quercetin and aluminum to increase during aging is not 

fully explained.

MODIFICATION OF ALUMINUM HYDROXIDE PRECIPITATE BY TRACES OF

QUERCETIN

Organic complexes of aluminum may influence the chemical 

behavior of the element in other ways, besides their effect on 

equilibrium solubility. Where concentrations of the oganic ligand are 

too small to form complexes with a major part of the dissolved 

aluminum, there may still be threshold effects that can influence the 

mechanisms whereby aluminum hydroxide and related species are 

precipitated, and the form of the precipitate can be substantially 

influenced by such effects. The influence of small concentrations of 

quercetin on reactions of aluminum was studied, and result^1 will be 

presented in the following sections of this report.

Solutions that precipitated aluminum hydroxide in previou^ studies 

by Hem and Roberson (1967), Roberson and Hem (1969), Schoen and 

Roberson (1970), and Smith and Hem (1972) were duplicated with the 

addition of quercetin as an example of organic material in concentrations that might occur in natural waters. The ratios of aluminum to 

quercetin, 100:1 and 100:4, were such that aluminum-quercetin 

complexes themselves could comprise only a very small fraction of the 

reacting species. In spite of this, there was a definite modification of the 

rate and possibly the mode of formation of aluminum hydroxides.

Two series of solutions were prepared, one containing 5 x 10 ~6 mole/1 

quercetin (0.9 mg/1 organic carbon), designated the Q21 and Q22 

series, and one containing 2xlO~5 mole/1 quercetin (3.6 mg/1 organic 

carbon), designated the Q23 series. The solutions contained 4.71 x 10~4 

mole/1 aluminum and were 0.01 in total ionic strength. Tl ^y were 

prepared by the method described in Hem and Roberson (1967), except 

with the addition of the amount of quercetin mentioned to solution I 

and twice that amount to solution III.
G38 CHEMISTRY OF ALUMINUM IN NATURAL WATER

BLOCKAGE OF CRYSTALLIZATION IN ACIDIC SOLUTIONS

The development of gibbsite in acidic solutions is obviously hindered 

by the presence of quercetin. Table 6 gives results of the procedure for 

determining the three forms of aluminum, described by Smith and 

Hem (1972), applied to solutions aged 121,126, and 127 days and also 

to those aged 310, 311, and 312 days. Solutions containing quercetin 

had substantially higher pH and greater concentratic ns of the 

unstable Alb polymer than did the comparable solutions where 

quercetin was absent. Figure 16 shows graphically that monomeric 

aluminum was produced in lower concentration in the presence of 

quercetin. The process of conversion of Alb to crystallire gibbsite 

released protons, and inhibition of the process by quercetin appears to 

be the logical explanation of the observed differences. The "apparent 

solubility" of Alc species after 310 days is represented in figure 17 as 

circles designating the monomeric aluminum concentration. In earlier 

studies where quercetin was absent, the solubilities after comparable 

aging times were substantially lower.

In the work of Hem and Roberson (1967), the solid formed during 

relatively short aging periods (1 to about 60 days) was described as

TABLE 6. Effect of quercetin on stability of aluminum species

Aluminum-quercetin solutions

Sample

Q21D __ 

Q21E __ 

Q21F __ 

Q23A __ 

Q21G__

Q22A __ 

Q22B __ 

Q23B __

Age 

(days)

0 

127 

312

_ 0 

127 

312

_ 0 

127 

312

_ 0 

121 

310

_ 0 

127 

312

. 0 

126 

311

_ 0 

126 

311

_ 0 

121 

310

rn 

PH ((OH):(A1

4.65 1.02 

4.51 

4.49

4.72 1.28 

4.56 

4.53

4.76 1.47 

4.59 ____

4.57

4.79 1.53 

4.68 ____ 

4.65

4.78 1.79 

4.65 

4.64 ____

4.86 1.89 

4.68 

4.63 __-._

5.06 2.42

4.74 ____ 

4.66

5.16 2.72 

4.90 

4.81 __

Ala Alb Alc

(moles/1 xlO")

1.81 

1.92

1.55 

1.74

1.54 

1.49

1.15 

1.13

.92 

1.10

.75 

.90

.36 

.31

.23 

.22

2.30 

1.90

2.38 

1.93

2.47 

1.94

1.95 

1.55

2.82 

2.20

2.44 

1.58

.67

.44

.34 

.11

0.59 

.89

.78 

1.03

.71 

1.28

1.61 

2.03

.96 

1.40

1.51 

2.23

3.68 

3.96

4.14 

4.38

Aluminum only

Age Sample (da» s)

C ______ 0

82 

121

D ______ 0 

82 

121 

259

pH (

4.52 

4.48 

4.47

4.66 

4.26 

4.23 

4.12

rn Ala Alb Alc

(rroles/1 xlO4 )

0.94 

1.36

3.40 

3.20 

3.23

2.84 

2.34 

2.39 

2.30

0.65 

.97 

.79

1.06 

.34 

.06 

.02

0.49 

.37 

.52

.64 

1.86 

2.09 

2.22

E ______ 0

77 

116 

254

F ______ 0

82 

121 

259

G______ 0

77 

116 

254

H______ 0

82 

121 

259

4.79 

4.54 

4.52 

4.32

4.88 

4.30 

4.28 

4.20

5.02 

4.40 

4.35 

4.28

5.23 

4.51 

4.49 

4.45

1.84 

2.13 

2.47 

2.76

2.24 

1.53 

1.60 

1.50

1.46 

1.20 

1.21 

1.02

1.17 

.68 

.66 

.62

.78 

.29 

.284 

.235

1.50 

1.45 

1.20 

.12

1.78 

.26 

.10 

.02

1.97 

.13 

.043 

.007

1.36 

.05 

.011

.50 

1.56 

1.74 

2.92

1.30 

3.08 

3.23 

3.50

1.40 

3.63 

3.84 

3.91

3.40 

4.20 

4.24 

4.30
EFFECTS OF ORGANIC SOLUTES G39

D 

?. 2

O no quercetin, 254 259 days

x quercetin 5x10"6 mole/I 

311-312 days

1.0 2.0 

r, MOLE RATIO TOTAL BOUND OH TO AL

3.0

FIGURE 16. Effect of small amounts of dissolved quercetin on the concentrations of 

monomeric aluminum species in aged solutions.

macrocrystalline gibbsite, and its solubility product (*KSQ) in the 

reaction

Al(OH)3(c)+3H+ =A13+ +3H20

was determined to be 109- 35 . The better crystallized material obtained 

by Smith and Hem (1972) after much longer aging had a *Kso value of 

108 - 22 . Data in table 7 indicate that in the presence of small amounts of 

quercetin the calculated solubility product for A1(OH)3 is near the 

macrocrystalline gibbsite value even after 312 days of aging.
G40 CHEMISTRY OF ALUMINUM IN NATURAL WATER 

-2.0

LU

tr

t -3.0

tr

LU

-4.0

-5.0

O

E -6.0

LU

1 

O

z

O

1

-7.0

-8.0

'I r

gibbsite-bayerite

4.0 5.0 6.0 7.0 

PH

8.0 9.0 10.0

FIGURE 17. Monomeric aluminum concentrations in solutions containing 1Q~530 

mole/1 of quercetin (circles) after 310 days aging compared with equilibrium values 

for microcrystalline gibbsite and well-crystallized gibbsite or bayerite.

Smith and Hem (1972) attributed the decrease in pH observed 

during aging to a process of polymerization whereby Alb was converted 

to Alc and protons were released. The presence of quercetin evidently 

slows down this process. The pH of solutions containing quercetin only 

decreased slightly during 1 year of aging. The concentrations of the 

organic ligand were too small to have a direct influence; rather there 

must have been a "threshold" effect, whereby the quercetin interfered 

with nucleation and growth of aluminum hydroxide particles.

AVERAGE SIZE OF ALUMINUM HYDROXIDE POLYMERIC IOHS

The process of polymerization of aluminum hydroxide ionic species 

to form gibbsite has been described by Hem and Roberson (1967) and 

Smith and Hem (1972) as the formation of a network cf double 

hydroxide bridges linking aluminum ions together in a geometric
EFFECTS OF ORGANIC SOLUTES G41 

TABLE 7. *Ks0 values of aged acidic solutions containing quercetin

Solution idavs) rn va^ue '°B *^sO

Q21D ________

Q21E ________

Q21F ___ _ _

Q21G ________

Q22A ________

Q22B _ ______

Q23A ________

Q23B ________

127

312 

127

312

127

312 

127

312 

126

311

126

311 

121

310 

121

310

1.02

1.28

1.47

1.79

1.89

2.42

1.53

2.72

9.32

9.29 

9.39

9.36 

9.46

9.40 

9.39

9.45 

9.38

9.33

(mean)9.38 

9.21

8.95 

9.57

9.48 

9.39

9.15

pattern of coalesced hexagonal rings. Bridging hydroxide is essentially 

unreactive, and its change of status in the solution frees protons and 

lowers the pH. Aluminum hydroxide polynuclear ionic species have 

been postulated by some investigators, usually with no definite 

structural consideration being given. It seems evident, however, that 

the aluminum ions in such species would necessarily be linked by the 

double hydroxide bridges. The determination of Alb *s by a kinetic 

procedure that depends on the slow rate of disruption of this bridged 

structure. It is possible to make some inferences about the relative size 

of Alb units in the presence of traces of quercetin from data in table 6 

and from structural considerations.

Amounts of Alb in some of the aluminum-quercetin solutions were 

initially larger than in the similar solutions of Smith and Hern (1972), 

given in table 6 for comparison, but this is not considered particularly 

significant. The initial concentration of Alb is m Part a function of the 

way the solutions are mixed.

The ratio of bound hydroxide to aluminum in the Alb fraction was 

calculated from stoichiometry of the solutions. This value ranged from 

1.4 to 1.9 and did not get as high as 2.0 in any quercetin solutions even 

after more than 300 days of aging. In the comparable solutions of 

Smith and Hem (1972), the OH: Al ratio in the Alb fraction ws greater 

than 2.0 after only 1 day. If all the bound OH is in bridging positions,
G42 CHEMISTRY OF ALUMINUM IN NATURAL WATER

the minimum value of the ratio for a six-membered hexagonal ring is 

2.0. At least some of the bound OH in the Al^ units is likely to be held in 

nonbridging positions. In any event, the low values of the ratio indicate 

that most of the Al^ units do not become large enough to include even 

one ring and thus do not attain the gibbsite structure during the 300 

days of aging.

The net effect of quercetin would appear to be one of slowing the 

polymerization process. Some Al^ does disappear, but at the rate 

indicated for the solutions cited in table 6, many years would he needed 

to complete the process. Only a few of the centers of polymev growth 

formed when the solutions are mixed are large enough to sustain 

growth toward a crystalline state. The smaller polynuclear ions may 

attain a considerable degree of metastability in such a solution, but 

such species still are not thermodynamically stable.

IDENTIFICATION OF CRYSTALLINE SPECIES

Principal evidence as to the crystalline nature of the gibbsite 

obtained in studies by Hem and Roberson (1967) and Sclroen and 

Roberson (1970) included X-ray diffraction patterns and electron 

micrographs. The material characteristically formed thin hexagonal 

platelets 0.1/xm (micrometre) or less in diameter with a distinctive 

appearance in electron micrographs, and aggregates of these particles 

gave X-ray diffraction scans with sharp peaks characteristic of 

gibbsite (fig. 18).

The X-ray diffraction procedure may fail to detect crystallirity if the 

crystals are very small or form less than about 10 percent of the total of 

solid material. The electron micrographs permit examination of very 

small fractions of the total precipitate and small amounts of crystalline 

material in an amorphous matrix could thus be overlooked by either 

procedure. The identification of crystalline material by either 

procedure, however, is definite evidence of important amounts of such 

material being present.

Of all the solutions containing quercetin, only one (Q22B) yielded 

material that could be identified as gibbsite by X-ray diffraction. The 

pattern for this material is shown in figure 19. Other solutior s having 

similar concentrations of quercetin did not show this peak. Those 

solutions where the initial OH:A1 molar ratio was below 1.89 showed 

no peaks at all. Those having higher ratios and some of the solutions 

having higher quercetin contents contained material that gave a broad 

and intense peak for a c?-spacing between about 6 and 8 A (angstroms). 

No positive assignment of structure of composition can be made from 

this evidence; however, a broad ill-defined bulge in the pattern was 

observed in some earlier work where quercetin was absent, at about 

this c?-spacing. Material of this kind has been called boehmite or
EFFECTS OF ORGANIC SOLUTES G43

gibbsite

bayerite

nordstrandite

Degrees 2 6(Cu K°c )

2.25 2.49 2.79 3.18 3.70 4.43 5.53 7.37 

"d" VALUES IN A UNITS

FIGURE 18. X-ray diffraction patterns of gibbsite, bayerite, and nordst-andite 

(Schoen and Roberson, 1970).

pseudo-boehmite by some workers, as boehmite has a major peak in 

this general range.

Electron micrographs of solids recovered from solution Q22B show 

small quantities of poorly formed gibbsite crystals. The other solutions 

which had OH:A1 ratios below 3 did not appear to contain any gibbsite 

crystals, but several had aggregates of layered material without 

well-defined crystal form, which might have produced the 6-8 A peaks 

just noted.

The presence of minute amounts of quercetin in acidic solutions 

evidently hinders the formation of gibbsite and seems to encourage the 

formation of significant quantities of poorly crystallized A1(OH)3 with 

a different structure. The material that is ultimately precipitated is 

relatively stable and is not amorphous to X-rays, but it has large, 

definite ranges of structural dimensions. Considering that the 

quercetin solutions were studied over a period of 310-312 days without 

producing well-formed gibbsite, they are a sharp contrast to those of 

Hem and Roberson (1967), who found microcrystalline gibbsite
G44 CHEMISTRY OF ALUMINUM IN NATURAL WATER

2.56 2.98 3.56 4.44 5.90 8.84 17.66 

"d" VALUES IN A UNITS

FIGURE 19. X-ray diffraction patterns for solid material recovered from solutions 

Q21D-F and Q22A-C aged 311 and 312 days. (The lower line on each graph is the 

pattern given by the empty glass slide. The darkened areas are where the pattern went 

off scale and the peak was diminished in size by one-half.)
EFFECTS OF ORGANIC SOLUTES G45

particles in solutions with rn's ((OH):(A1)) from 2.0 to 3.0 within 2-10 

days of aging at 25°C, and to those of Smith and Hem (1972), who 

achieved a solubility value of well-formed gibbsite for their solutions 

within 254-259 days.

BLOCKAGE OF CRYSTALLIZATION IN BASIC SOLUTIONS

The formation of bayerite which normally occurs when aluminum is 

precipitated in an alkaline solution is also retarded by the presence of 

quercetin. An inspection of the *KS ^ values, represented by the 

equation

A1(OH)3 (C) +H2O = A1(OH)4- + H+ ,

will illustrate this. Hem and Roberson (1967) found a solubility 

product of 10~ 12- 71 in their day-old solutions of pH 7.5-9.5, but the 

material was amorphous to X-rays. Within 10 days, the solubility 

product reached that of bayerite, 10~ 13 -95. The solid present at that 

time had an X-ray pattern of bayerite. In the aluminum-quercetin 

solutions, with the lesser amount of quercetin and with a pF greater 

than the minimum solubility of gibbsite and bayerite (well above pH of 

6.0), the *KS 4 values were higher than that of bayerite from age 3 to 311 

days. With the solution of rn of 4.42, the *KS 4 eventually closely 

approached that of bayerite; the period of most marked decrease in 

*KS 4 in this solution was between ages 40 and 126 days. In the series 

with more quercetin, the *KS 4 values were close to the amorphous 

material value even though aged 121 days, and by age 509 days the 

value for Q23G, the most basic solution in this series, was still only 

midway between that of the amorphous material and that of bayerite. 

(See table 8.)

DEVELOPMENT OF BAYERITE

Some stages in the development of bayerite crystals can be traced by 

means of the electron micrographs of precipitates from Q22E, rn of 

3.65. (See figs. 20-24.) Early in the aging period, a sheetlike mesh of 

interlaced crystals is formed. On further aging this mesh assumes a 

"jackstraw" pattern, and a few scattered bayerite crystals appear (fig. 

20). At this time a general view of this material indicates no 

widespread crystallinity. Upon additional aging, better ordered 

bayerite crystals appear (fig. 21). A general view of this material 

indicates an occasional well-formed crystal shape but meetly still 

poorly defined fragments (fig. 22). The 126-day-old solid still had a 

mixture of free and enmeshed crystals. The micrographs seem to 

indicate that where the interlaced mesh is absent the bayerite 

particles themselves once again have only the fine pebbly texture 

brought out by coating with platinum (fig. 23), but some of the bayerite 

crystals leave the impression that the mesh is part of the crystal itself.
G46 CHEMISTRY OF ALUMINUM IN NATURAL WATER

TABLE 8. Composition during aging of solutions having rn values greater than 2.75

Solution ((OH "(A!))

Q22C ____ 2.76

Q22D ____ 2.99

Q22E ___- 3.65

Q22F ___, 4.42

Age

(days)

0

4

19

40

126

311

0

4

19

40

126

311

514

0

4

19

40

126

311

0

4

19

40

126

311

PH

5.62

5.50

5.29

5.16

5.03

4.89

7.75

7.62

7.09

6.87

6.70

7.03

6.87

8.63

9.39

9.43

9.31

9.31

7.93

10.05

9.92

9.76

9.68

9.90

9.63

p(Al) 1

£UO

4.50

4.63

4.97

4.98

5~43

5.53

6^43

6.43

6.23

3767

3.77

3.81

4.20

5.43

3.37

3.43

3.39

3.78

4.13

log*#S4

___-

-Ts.io

-12.67

-l¥.18

-13.51

-13.15

-13~.ll

-13.25

-13.17

-13.56

-13.41

-13.34

-13.24

-13.12

-13.73

-13.81

r" Age Solution ((0H):(A1)) (days)

Q23C ____ 2.85 0

3

21

41

121

310

Q23D __ _2.93 0

3

21

41

121

310

509

Q23E -.__ 3.02 0

21

41

121

Q23F ____ 3.27 0

3

21

41

121

Q23G ____ 3.46 0

3

21

41

121

310

509

PH

5.31

5.19

5.07

5.01

4.96

4.86

5.79

5.71

5.44

5.34

5.20

5.12

5.07

6.11

6.51

6.62

6.75

8.00

7.94

7.63

7.32

7.29

8.85

8.69

8.57

8.82

9.02

8.97

8.45

p'Al) 1

4~36

4.42

4.47

4.68

4.70

5~09

5.09

5.59

5.65

6.43

6.23

613

4^51

525

5~59

3 74

368

360

381

406

4.74

log*#S4

___-

____

----

____

-12.93

-12.56

-12.93

-T2~93

-12~.48

-12.30

-12.47

-12.88

-13.08

-13.24

1 Aluminum determined by the regular ferron method.

Figure 24 shows that even at age 311 days there still were bayerite 

crystals forming. This micrograph also is a pronounced demonstration 

of the pyramid shape that the bayerite crystal assumes. Tve whole 

series of micrographs gives real visual evidence of the existence of an 

intermediate material for an extended period of time during the 

organizational development of aluminum trihydroxides.

A third aluminum trihydroxide, nordstrandite, has been synthesized previously also. It has the same free energy as bayerite (Hem 

and others, 1973). Schoen and Roberson (1970) suggested that 

nordstrandite forms in solutions of pH 5.8 and more and that the 

sequence of crystallization in mildly alkaline solutions is from early 

formed bayerite to nordstrandite. In an attempt to determine the 

importance of pH in the formation of bayerite and nordstrandite, they 

found that in solutions near neutral pH these two minerals 

precipitated simultaneously. Some of both may be present in figure 25. 

The X-ray pattern for material from this solution (fig. 26) has a peak 

near the "d" spacing of boehmite, but because of uncertainties in the
EFFECTS OF ORGANIC SOLUTES G47

FIGURE 20. Electron micrograph of solids recovered from solution Q22E aged 40 days. 

A bayerite crystal beginning to form.

literature regarding identification and solubility of this mineral, it 

cannot be conclusively assigned.

Solution Q22F, rn of 4.42, had well-defined crystal shapes 

characteristic of bayerite and nordstrandite by an age of 126 days. The 

X-ray diffraction pattern after aging 311 days had peaks characteristic 

of bayerite (fig. 26). Nordstrandite crystals identified by Schoen and 

Roberson (1970) were elongated rectangles or parallelograms. 

Electron micrographs of this sample at a shorter aging time gave no 

indication of crystallinity and contained filmy amorphous matter 

having no observable texture.

No crystallinity was observed in electron micrographs of material 

from any of the Q23 solutions with the exception of Q23G. When aged 

310 days, Q23G had the "jackstraw" pattern and, when aged 509 days, 

in addition showed faint dots that possibly indicated focal points for the 

initiation of some type of further crystallization. The micrographs of 

the solids from this series supplement the solubility data in helping to
G48 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 21 Electron micrograph of solids recovered from solutior Q22E aged 57 days. 

A more clearly denned bayerite crystal.

substantiate that the crystallization process is even more retarded in 

the presence of higher concentrations of quercetin.

INTERMEDIATE PRODUCTS

The X-ray diffraction patterns of the quercetin series fit the trend 

established by the other data, namely, that the formation of bayerite is 

favored more strongly as the pH becomes more basic and that 

intermediate products are more prominent in less basic solutions and 

in those with higher concentrations of quercetin. (See figs. 27, 28.)

By the time all the basic solutions (with the exception of Q22F 

which was not tested) were aged 121-126 days, they were shown to 

have indication of a broad peak as pictured in figure 26 for Q22E at age 

311 days. The point of this peak suggests a "d" spacing range very near 

that of natural boehmite. The electron micrographs of Q23E and Q23F, 

age 126 days, did not show any of the precursor crystal mesh, and 

unfortunately no further crystallization could be traced in these 

solutions, as they had been used up. The entanglement of the bayerite 

crystals in the mesh of poorly organized material in figure 29 may
EFFECTS OF ORGANIC SOLUTES G49

FIGURE 22. Electron micrograph of solids recovered from solution Q22E aged 57 days. 

A general view of the solids.

explain why the X-ray diffraction pattern showed only one tiny 

bayerite peak, representing a "d" spacing of 4.34 A, at this age for 

Q22E. Along with this peak there was the described broad peak which 

was already pronounced both for it and for Q23G by this time. These 

same broad high peaks persisted even at age 514 days for Q22D and at 

age 509 days for Q23D and Q23G, but there was no indication of the 

development of new peaks as in the case of Q22E.

TOTAL EFFECT

Quercetin definitely slows the precipitation of bayerite and 

nordstrandite. The result may be described as temporary stabilization 

of dissolved intermediate substances having some similar properties to 

those found after very short aging periods in supersaturated solutions 

containing only aluminum. Also, crystalline material having a 

solubility near or slightly less than microcrystalline trihydroxide is
G50 CHEMISTRY OF ALUMINUM IN NATURAL WATER

-if:

FIGURE 23. Electron micrograph of solids recovered from solution Q22E aged 126 days.

formed, with an X-ray diffraction peak very near that of natural 

boehmite. The structural dimensions and possibly the composition of 

the material evidently depend on the specifications of the test solution. 

A review of the X-ray diffraction patterns indicates that there is little 

or no change in the location of the peak points on aging but that the 

intensities of the peaks increased dramatically with time. This would 

suggest that the intermediate structures did not convert to new 

structures with time but became better established.

EFFECTS OF ORGANIC SOLUTES ON ALUMINOSILICATE 

PRECIPITATION

The presence of aqueous silicic acid (silica) in solutions that are 

moderately supersaturated with respect to gibbsite or bayerite can 

cause substantial changes in the nature of the precipitated product and 

in the chain of chemical events leading to a state of equilibrium. The 

precipitated material described by Hem, Roberson, Lind, and Polzer
EFFECTS OF ORGANIC SOLUTES G51

0.1 /urn

FIGURE 24. Electron micrograph of solids recovered from solution Q22E aged 311 days. 

A bayerite crystal and associated mesh.

(1973) commonly was a mixture of amorphous aluminum hydroxide 

and a 1:1 aluminosilicate with the composition and solubility of 

halloysite. The optimum conditions for the precipitation of this 

synthetic halloysite appear from the data published by Hem, 

Roberson, Lind, and Polzer (1973) to include a substantial excess of 

silica over that of aluminum in solution and a pH near or a little above 

neutrality. All solutions are relatively dilute, although perhaps more 

concentrated than some weathering solutions in well-watered environments. The temperature was near 25°C in all solutions. At pH 

levels below 5.0 the precipitates were chemically similar but required 

a very long aging period before they stabilized (several years in some 

instances).

Research on clay-mineral synthesis has been extensive, and many 

reports on the results of this work have been cited in our earlier 

publications. One aspect rather briefly touched upon, however, is the 

possible influence of organic material on low temperature synthesis of
G52 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 25. Electron micrograph of solids recovered from solution Q22E aged 311 days. 

Possible nordstrandite crystals along with bayerite crystals.

clay minerals. Some investigators have prepared a gel, for example, by 

mixing organically complexed aluminum and silica species and have 

shown that a 1:1 clay mineral evolves during aging at a higher 

temperature (DeKimpe, 1967, 1969). Linares and Huertas (1971)

FIGURE 26. X-ray diffraction patterns for solids recovered from solutions Q22D-F aged 

311 days. Darkened areas are where the pattern went off scale and the peak was 

diminished by one-half. The scratched areas are where the peak was diminished by 

one-tenth. The lower line of each graph is the pattern given by the empty glass slide. 

The boxed-in areas are small peaks found far beyond the "d" spacings shown for the 

other test solutions. These small peaks have been repeated to verify their existence.
EFFECTS OF ORGANIC SOLUTES G53

2.56 2.98 3.56 4.44 5.90 8.84 17.66

Q22F Continued

1.34 1.43 1.54 1.67 1.82

"d" VALUES IN A UNITS

2.01 2.25
G54 CHEMISTRY OF ALUMINUM IN NATURAL WATER

2.56 2.98 3.56 4.44 5.90 

"d" VALUES IN A UNITS

8.84 17.66

FIGURE 27. X-ray diffraction patterns for solids recovered from solutions Q23A-C 

aged 310 days. Darkened areas are where the pattern went off scale and the peak was 

diminished by one-half. The scratched areas are where the peak was diminished by 

one-tenth. The lower line on each graph is the pattern given by the empty glass slide. 

The boxed-in area contains a peak that has been repeated to verify its position and 

shape.

reported a synthesis of kaolinite at 25°C using what they considered to 

be a fulvic acid-aluminum complex as one starting material.

The rationale for expecting that organically complexed starting 

material might expedite the formation of clay may perhaps be related 

to the structure of aqueous aluminum and silicon hydroxides. An
EFFECTS OF ORGANIC SOLUTES G55

 w^ 

^^^^jt*r 

2.56

^v* -

1

2.98

i i

3.56 4.44

X.^V

i i

5.90 8.84

*****

17.66

"d" VALUES IN A UNITS

FIGURE 28. X-ray diffraction patterns for solids recovered from solutions Q23D and 

Q23G aged 310 days. Darkened areas are where the pattern went off scale and the 

peak was diminished by one-half. The lower line on each graph is the pattern given by 

the empty glass slide. The boxed-in area contains a peak that has been repeated to 

verify its position and shape.

electrically neutral sheet polymer composed of aluminum ions in 

six-fold coordination with close-packed hydroxide ions constitutes the 

gibbsite structure, and a somewhat similar but more open sheet 

composed of silicon ions in fourfold coordination with oxygen 

constitutes the silica layer of a 1:1 clay. Although such a clay is 

sometimes described as having gibbsite layers alternating with silica 

layers, this is an oversimplification that obscures certain important 

structural factors. Possibly a noncrystalline material could be 

synthesized by adsorption of a layer of uncharged Si(OH)4 tetrahedra 

on the surface of a sheet structure of A1(OH)3, followed by
G56 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 29. Electron micrograph of solids recovered from solution Q22E aged 126 days.

polymerization and dehydration of the silica layer into an (81205 2~)rt 

sheet. The initial formula for this material would be A1(OH)3 Si(OH)4 , 

or Al2Si2(OH)1 4. It could be converted to the composition of kaolinite or 

halloysite by dehydration:

The apparent simplicity of this process is deceiving, for actually each 

water molecule produced probably involves the transfer of a proton 

from one OH~ ion to another OH~, leaving an O2 ~ and converting the 

second OH~ to a water molecule. To form the complete kaolinite 

structure, this water molecule would need to be removed from the 

reaction site. The final product contains an extensive network of 

Si-O-Al bonding, and each aluminum ion participates in two such 

bonds. The third unit of charge on each aluminum ion remains shared 

in a double hydroxide bridge with an adjacent aluminum ion. Because 

this kind of mechanism is so complicated, it is likely to be slow and 

difficult, especially if the gibbsite structure has developed extensively.

Most of the more successful kaolinite syntheses have resorted to
EFFECTS OF ORGANIC SOLUTES G57

high temperature and pressure to attain crystallinity (Hemley, 1959). 

Kittrick (1970) was able to crystallize clay minerals from aqueous 

solutions by providing previously formed clay minerals as nucleation 

surfaces. The work of Hem, Roberson, Lind, and Polzer (1973) does 

indicate that a small yield of crystalline material can be obtained at 

25°C, along with a large amount of poorly organized 1:1 clay, during 

long aging.

A possible advantage in using organically complexed aluminum and 

(or) silica as starting material is that Al-O bonding may be present in 

the organic structure and could thus be present as the synthesis 

reaction begins. Then the crystallization of the clay could proceed, 

forming Al-O-Si bonds without the need for removing water molecules 

from their coordination positions in the lattice. Also, the organic 

material might inhibit the polymerization of aluminum hydroxide and 

thus provide a less established crystal structure at the start of the 

synthesis.

MODIFICATION OF THE PRECIPITATION OF ALUMINOSILICATES BY 

TRACES OF QUERCETIN

To further verify the organoaluminum interactions with clay 

minerals, comparison of reactions with and without organic matter 

under controlled laboratory conditions is necessary. Solutions that 

precipitated a metastable aluminosilicate in previous studies by Hem, 

Roberson, Lind, and Polzer (1973) were duplicated with the exception 

of the addition of quercetin as a representative model of organic 

material. Again the concentration of quercetin used was that to be 

found for organic matter in natural water. The ratios of aluminum tp 

quercetin, 100:1, were such that aluminum-quercetin complexes 

themselves were, at most, only small fractions of the reacting species.

Three series of solutions containing 5 x 10~6 mole/1 (0.9 mg/1 organic 

carbon) were prepared. The first series was labeled Q19 and Q20, the 

second Q24, and the third Q25 and Q26. The method described by Hem, 

Roberson, Lind, and Polzer (1973), where silicic acid reagent was 

added directly to the basic solution II, was used. The quercetin was 

added in the amount mentioned to solution I and twice that amount to 

solution III. The initial and subsequent concentrations of aluminum 

and silica are listed in tables 9 and 10, with IAPa (ion-activity product) 

defined by equation 8, table 3, and IAPb by equation 9, table 3. The 

solids from these solutions were compared with well-identified 

minerals and with a soil sample of natural streambed sediment. The 

minerals are described as follows:

Halloysite 13, Eureka, Utah, (percent impurities: limonite 2; 

orthoclase 1-1.5, quartz trace, and alunite trace). 1

"Occurrence and Micro Examination of Reference Clay Mineral Species, M.S. Main, API project 49, Clay Mineral 

Standards 1950, Preliminary Report No. 5.
G58 CHEMISTRY OF ALUMINUM IN NATURAL WATER

Kaolinite 9, Mesa Alta, N. Mex., (percent impurities: sericite

1-1.5, orthoclase 1.5, limonite 11.5, and leucoxene I).1 

Keokuk Geode Kaolinite, Keokuk, Iowa.2

'Occurrence and Micro Examination of Reference Clay Mineral Species, M.S. Main, API project 49, Clay Mineral 

Standards 1950, Preliminary Report No. 5.

2A sample of this kaolinite was furnished by Professor W. D. Keller, University of Missouri, and has been described 

by Keller, Pickett, and Reesman (1966).

TABLE 9. Compositions of acidic aluminum-silica-quercetin solutions before and after

aging

Initial contents Aged solutions

Solution

Q19D _ .

Q19E ____

Q20A __

Q20B __

Q20B __

Q20C

Q19G ___.

Q20D _ _

Q20D -__

Q24A __

Q24A

Al 

(mg/1)

12.7

12.7

12.7

12.7

12.7

12.7

12.7

12.7

12.7

12.2

12.2

Si02 

(mg/1)

57.8

59.4

38.2

39.7

39.7

41.3

64.4

43.9

43.9

44.8

44.8

pH

4.58

4.63

4.63

4.68

4.68

4.76

4.86

5.37

5.37

5.51

5.51

days

332

332

155

155

332

155

36

155

332

136

251

Al 

(mg/1)

8.94

7.94

10.8

9.47

5.52

8.08

7.6

4.78

2.91

.01

.00

SiO2 

(mg/1)

54.4

55.4

36.4

38.3

37.6

39.8

63.6

40.1

35.5

29.7

29.4

pH

4.14

4.18

4.20

4.25

4.17

4.34

4.44

4.67

4.62

5.44

5.44

j Precipitate from 

TAP solutions Si;Al 

a (mole/mole)

10.98

11.10

11.12

11.34

10.42

11.76

12.66

13.08

12.32

10.56

10.56

0.406

.377

.425

.190

.131

.146

.070

.215

.385

.556

.567

TABLE 10. Composition of neutral and basic

Initial contents Aged solutions

Solution

Al 

(mg/1) (mg/1) pH Al 

lmg/1 (mg/1) pH A,

Precipitate A1

Age 155 days

Q20E _ 

Q20H III

-12.7 

-12.7 

-12.7 

_12.7

45.2 

46.2 

46.8 

48.3

6.47 

7.91 

8.62 

8.96

0.01 

.09 

.19

.54

26.6 

26.0 

25.4 

23.7

6.53 

7.34 

7.54 

7.55

-32.71 

-32.45 

-32.21 

-31.39

0.658

.719 0.01 

.768

.908 .04

(mg/1) PH

Age

25.4 7.44

23.7 7.87

Q24B 

Q24C 

Q24D 

Q24E 

Q24F 

Q24G 

Q24H

Q25A 

Q25B 

Q25C 

Q25D

12.2

12.2

12.2

12.2

12.2

.12.2

.12.2

.12.2 

12.2

.12.2 

12.2

Age 136 days Age

46.4

47.0

47.6

48.2

48.8

49.2

51.4

49.1

48.2

49.2

48.2

6.76

7.58

7.95

8.36

8.64

8.77

9.33

8.79

8.36

8.81

0.02

.04

.05

.11

.26

.46

3.06

27.3

26.9

26.5

26.4

24.4

23.3

21.4

6.75

7.05

7.29

7.44

7.55

7.64

8.92

-32.53

-32.55

-32.83

-32.47

-32.01

-31.73

-32.65

0704

.742

.780

.810

.918

.995

1.47

0.00

.01

.03

.04

.08

.25

1.08

27.4

26.8

26.2

24.8

23.7

22.1

16.1

6.80

7.00

7.16

7.31

7.36

7.83

795

Age 114 days Age

1.36

.84

1.00

.48

25.8

27.4

25.4

26.2

7.79

7.74

7.95

8.06

-30.92

-31.18

-31.50

-32.44

0.965

.822

.954

.866

0.37

.49

.21

.10

22.9

25.0

23.7

24.4

7.53

7.46

7.54

7.40

Age 115 days

Q26A __ 12.2 49.6 8.95

Q26B __ 12.2 48.8 8.68

Q26C ._ 12.2 49.4 9.04

Q26D __12.2 48.6 8.91

2.16 26.2 8.02 -31.04 1.04

1.49 25.8 7.94 -31.22 .964

.81 25.0 7.72 -31.34 .962

.55 26.0 7.73 -31.66 .871

.66 21.7 7.54

.35 22.5 7.49

.14 23.7 7.38

.14 24.2 7.38
EFFECTS OF ORGANIC SOLUTES G59

The soil sample was the <200-mesh fraction of bottom sediment taken 

from Colma Creek at Colma, San Mateo County, Calif., by C. E. 

Roberson on October 2, 1972.

CRYSTALLINITY OF PRECIPITATES

Electron micrographs of the two standard kaolinites and one of the 

halloysites are shown in figures 30-32. The characteristic thin layers 

and sharp straight edges and angles of kaolinite crystals in both large 

and small units are shown in figures 30 and 31. Surface texture on 

some of the crystals is the effect of shadowing material used in 

preparing the sample. The electron micrograph of halloysite (fig. 32) 

shows some of the flat sheets similar in appearance to kaolinite and 

rolled up material more typical of halloysite. The thin sheets 

apparently roll up and sometimes the rolls appear to have been 

flattened after forming.

Analytical data in table 9 indicate smaller losses of aluminum and 

silica during aging but a somewhat greater decline in pH than do 

comparable acidic solutions without quercetin studied by Hem, 

Roberson, Lind, and Polzer (1973, p. 11). Also, the silica to aluminum

aluminum-silica-quercetin solutions

Aged solutions Continued

log ra Al Si02 

IAPb (mole'stmole,' ^ pH log 

IAP

Precipitate 

(moles/mole)

Slo2

pH

i nB. Precipitate

( ^4 , , (moles/mole)

332 days

-34.57

-34.29

0.736

.873

251 days Age 358 days Age 481 days

-33.23

-33.65

-33.03

-33.13

-32.68

-32.68

-31.90

0.699

.744

.790

.864

.930

1.02

1.40

0.01

.02

.05

.07

.07

.24

.57

26.1

25.9

25.9

24.5

22.9

21.8

16.2

6.28

6.76

7.31

7.39

7.48

7.63

7.77

-32.24

-32.60

-32.91

-32.82

-33.06

-32.33

-32.12

0.748

.778

.802

.877

.959

1.03

1.36

0.02

.05

.05

.10

.14

.17

.26

26.4

26.2

25.3

25.0

24.1

22.4

16.9

6.77

7.07

7.27

7.48

7.55

7.66

7.55

-32.51

-32.33

-32.76

-32.58

-32.46

-32.58

-32.23

0.737

.769

.824

.861

.920

1.00

1.30

230 days Age 328 days Age 451 days

-31.72

-31.26

-32.19

-32.50

0.994

.900

.955

.883

0.37

.26

.21

.12

22.9

24.3

23.3

24.4

7.49

7.48

7.48

7.49

-31.63

-31.87

-32.09

-32.56

0.994

.899

.970

.885

0.16

.12

.09

.09

23.3

24.7

23.9

25.2

7.66

7.65

7.66

7.65

-32.60

-32.78

-33.07

-33.01

0.962

.873

.938

.875

-31.28 1.09 .42 22.9 7.66 -31.86 1.02

-31.70 .997 .37 22.2 7.55 -31.78 1.01

-32.20 .957 .13 23.7 7.54 -32.61 .956

-32.22 .909 .13 24.0 7.51 -32.54 .915

15 21.9 7.4 -32.2 103

.12 23.0 7.4 -32.3 .959

.08 24.6 7.4 -32.6 .919

.08 25.3 7.4 -32.6 .863
G60 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 30. Electron micrograph of kaolinite from Mesa Alta, N. Mex.

mole ratio of the precipitates indicated in table 9 is lower (further from 

the desired 1:1 ratio) than was found in solutions without quercetin. 

Possibly these differences could be attributed to the formation of a 

smaller amount of better ordered material when the organic ligand is 

present. However, the net solubility as indicated by the ion activity 

products does not appear to be greatly different, perhaps because the 

major solubility control in both sets of experiments is the predominating poorly crystallized material.

Neutral and basic solutions that had precipitated aluminosilicate 

material in the presence of low concentrations of quercetin gave 

evidence of much better crystallized material than had been obtained 

by Hem, Roberson, Lind, and Polzer (1973) without organic matter 

being present. Data in table 10 indicate about the same amount of 

aluminum and silica precipitation, though, and a similar rather close 

approach to a 1:1 mole ratio of silica to aluminum. There seems, 

however, to be a tendency for a more nearly uniform and somewhat 

lower final pH when quercetin was present. This again may indicate
EFFECTS OF ORGANIC SOLUTES G61

0,1 /UPI

FIGURE 31. Electron micrograph of Keokuk geode kaolinite (Keller and others, 1966).

that the product formed is more stable when the organic ligand is 

available.

As noted in table 10, the ion activity products for Q20E-H were 

similar to other values for solutions that contained solids identified as 

synthetic halloysite. Their pH was between 6.53 and 7.55 after 155 

days of aging, and they contained numerous small but very sharply 

defined kaolinitic crystals. Typical electron micrographs are shown in 

figures 33 and 34. Not enough of the well-crystallized material could be 

isolated to obtain an X-ray diffraction pattern. However, after aging 

332 days the log ion activity products for Q20F and H were more 

negative than before (approaching that of kaolinite), but no further 

determinations could be made because the solutions had all been used 

up.

To verify the results obtained in the Q20 series, another set of 

experimental solutions was prepared, the Q24 series. All these 

solutions had pH and solute concentrations in the range covered by the 

Q20 samples in which the best crystallinity was observed. After aging 

251 days, some layered material could be seen, but well-defined
G62 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 32. Electron micrograph of halloysite from Eureka, Utah.

crystallinity was lacking (fig. 35). After aging 358 days, rolled 

structures similar to halloysite (fig. 36) and somewhat more 

crystalline layered material (fig. 37) were observed.

Finally, after aging 481 days well-crystallized kaolinitic material 

was found in the precipitates of all solutions in the Q24 series. Four 

typical electron micrographs are given in figures 38-40.

The X-ray diffractograms for the precipitates still showed no 

well-defined peaks, presumably because poorly crystallized material 

still constituted the bulk of the solid. However, there was a general 

increase in signal above that given by an empty slide covering the "d" 

spacings for the more prominent kaolinite and halloysite peaks. This 

increase in signal became more pronounced with aging. At age 481 

days sample Q24F (dried at 250°C) was scanned several times at
EFFECTS OF ORGANIC SOLUTES G63

»

,>0.1 Mm

FIGURE 33. Electron micrograph of solids recovered from solution Q20G aged 155 days 

(from Hem and Lind, 1974). Copyright 1974 by the American Association for the 

Advancement of Science.

settings giving an amplified and somewhat noisy signal. Halloysite 

and metahalloysite have strong peaks near 7.4, 4.4, and 3.6 A. In five 

scans the signal consistently was high around 7.4 A. In four scans 

made covering 3.6 A, three had a high signal at this spacing.

The product in the Q24 series is comparable in all respects with that 

synthesized in the Q20 series. The length of time required to reach 

crystallinity appears to have been different, but the organic 

intermediate does seem to offer a feasible path to kaolinite synthesis at

9 fVT1

^° *-* SOLUBILITY OF PRECIPITATE

As noted in previous work, the rate at which the basic solution is
G64 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 34. Electron micrograph of solids recovered from solution Q20G aged 155 days 

(another view) (from Hem and Lind, 1974). Copyright 1974 by the American 

Association for the Advancement of Science.

added to the other two has an effect on the amounts of unstable 

aluminum hydroxide species that are formed, at least in those 

solutions where the pH remains below neutrality. To test the possible 

importance of this factor, solutions Q25A-D and Q26A-D, having 

specifications similar to Q20G, were prepared in several different 

ways, as described in table 11. The stock solutions I, II, and III were 

prepared in the same manner as described for the earlier aluminosilicate series. These solutions were added in such ratios as to make mixed 

solutions of two concentrations of aluminum and silica for each mode of 

addition. (The contents of the mixed solutions initially and after aging 

are listed in table 10.) In the Q25 series the mixtures of aluminum 

perchlorate and quercetin were allowed to stand for 3 days, before 

adding silica and sodium hydroxide, to test the idea that time might 

promote the formation of the aluminum-quercetin complex. In the Q26 

series this preliminary aging step was omitted. Two of the solutions in
EFFECTS OF ORGANIC SOLUTES G65

FIGURE 35. Electron micrograph of solids recovered from solution Q24D aged 251 days. 

TABLE 11. Preparation of solutions Q25A-D and Q26A-D

Solution Initial 

pH

Aee of solutions I and

III at time of addition

of solution II

(days)

Time interval over

which solution II

(basic) was added

(minutes)

Q25A 

Q25B 

Q25C 

Q25D 

Q26A 

Q26B 

Q26C 

Q26D

8.79

8.36

8.98

8.81

8.95

8.68

9.04

8.91

4

4

60

60

4

4

60

60

each group were brought to their starting pH by rapid mixing of the 

stock solutions and two by very slow addition of the basic solution with 

stirring. 

During the preparation of these solutions, the pH passed through the
G66 CHEMISTRY OF ALUMINUM IN NATURAL WATER

OJl

FIGURE 36. Electron micrograph of solids recovered from solution Q24C aged 358 days.

minimum solubility region for synthetic halloysite, and it seems 

reasonable that when the pH was in this range (near 6.0) for a 

relatively long time the precipitation would be more nearly complete 

and the structure better organized than if the pH were more rapidly 

increased to its final level. Thus, the solutions with short mixing times 

would be expected to be further from equilibrium at the start of aging 

than the ones with long mixing times. That this expectation is realized 

is indicated by the data in table 10. The ion activity products for short 

mixing-time solutions are less negative than those for corresponding 

long mixing-time solutions in every instance. Even after 328 days of 

aging this difference is discernible. The effect of the additional aging 

time after preparing solutions I and III before adding solution II is not 

discernible in any of the later analyses of the solutions.

Crystallization in the Q25 and Q26 solutions appeared to follow the
EFFECTS OF ORGANIC SOLUTES G67

FIGURE 37. Electron micrograph of solids recovered from solution Q24G aged 358 days.

same path as in Q24 solutions. The final observations of these showed 

small amounts of well-crystallized kaolinite and halloysite were 

present in the mass of poorly organized material (fig. 41).

STREAM SEDIMENT

The sample of sediment from Colma Creek contained considerable 

amounts of quartz and apparently some halloysite, according to its 

X-ray diffraction pattern. It exhibited the main 3.34-A peak and the 

10.1-A peak of hydrated halloysite (endellite), along with suggestions 

of the 4.42-A and 2.56-A peaks. An electron micrograph of some of the 

particles is shown in figure 42.

The natural material is of interest because it shows that stream 

sediments may contain significant amounts of aluminosilicate 

minerals similar in appearance and probably in other properties to the 

synthesized halloysite. The chemical behavior of sediments will, 

therefore, have some traits in common with synthetic halloysite.
G68 CHEMISTRY OF ALUMINUM IN NATURAL WATER

0,1 ftm

FIGURE 38. Electron micrograph of solids recovered from solution Q24C aged 481 days.

SUMMARY

Small amounts of quercetin appear to expedite the processes of 

clay-mineral synthesis at neutral or mildly alkaline pH. At lower pH 

(<5.0) the quercetin seems to inhibit the process. Well-crystallized 

kaolinite was obtained in several solutions in one experimental series, 

and later efforts to repeat these results were successful, although they 

required a somewhat longer aging time. The weight of evidence, 

including solubility, stoichiometry, and electron micrography, all 

supports the formation of a kaolinitic clay mineral. The yield of 

well-crystallized material obviously is small, but probably the yield 

could be improved by further refinement of the technique.

Considering these facts and the close interrelationships of organic 

matter with aluminum and clays in natural systems and of the 

evidence of changes of composition and structure of minerals by
EFFECTS OF ORGANIC SOLUTES G69

FIGURE 39. Electron micrograph of solids recovered from solution Q24E aged 481 days.

organic matter, it may be concluded that there is no doubt that 

water-soluble organic matter can be a vital factor in the ultimate 

composition and structure of clay minerals in contact with water.

IMPLICATIONS OF ORGANIC COMPLEXING IN ALUMINUM

GEOCHEMISTRY

Some indication of the difficulties involved in evaluating natural 

aluminum solubility controls can be obtained from calculations and 

graphs given earlier in this report. Certain of the conditions that must 

be specified for a calculation would normally be known or measurable. 

These include the total ionic strength of the solution, its content of 

complexing ligands and other dissolved species of interest, and the 

temperature and pressure. Existence of specific solids and their purity,
G70 CHEMISTRY OF ALUMINUM IN NATURAL WATER

"i*

0.1

FIGURE 40. Electron micrograph of solids recovered from solution Q24G aged 481 days 

(from Hem and Lind, 1974). Copyright 1974 by the American Association for the 

Advancement of Science.

or stability, is more difficult to ascertain, and the assumption of 

equilibrium, implying a ready reversibility of the governing reactions, 

is likely to be unreliable. Even if all the uncertainty were reduced to a 

question of the stability or crystallinity of the controlling 1:1 clays, it is 

apparent from figure 1 that the aluminum concentration at equilibrium might have a range of more than two powers of 10.

Considerations of reversibility and form of solid tend to be 

interrelated. For example, the studies of Hem, Roberson, Lind, and 

Polzer (1973, p. 15) showed a reversible solubility for their freshly 

precipitated synthetic halloysite. Better crystallized material has 

been found to form so slowly in the pH range studied here that a 

directly reversible solubility for crystalline kaolinite, the least soluble 

solid considered in figure 1, seems very improbable under these 

conditions. A more logical mechanism would probably view this form
EFFECTS OF ORGANIC SOLUTES G71

0.1

FIGURE 41. Electron micrograph of solids recovered from solution Q26B aged 451 days.

of kaolinite to be an equilibrium species only at conditions other than 

25°C and 1 atmosphere pressure and in solutions differing from the 

general composition we have studied.

The relationships in Al-Si-HsO systems might be represented 

schematically as

Halloysite 

-H20

Kaolinite

A13+ + SiO2(aq)

+OH~

ft

Gibbsite 

or

Bayerite

Equilibrium can be established wnere reversibility is indicated. 

Formation of aluminum hydroxide in preference to clays, or alteration
G72 CHEMISTRY OF ALUMINUM IN NATURAL WATER

f; *

FIGURE 42. Electron micrograph of fine-grained material separated from Colnia Creek

bed sediment.

of the latter to aluminum hydroxide, is indicated by this relationship to 

depend on dissolved activity of silica and hydrogen ions.

Natural weathering regimes capable of direct precipitation of either 

clay minerals or aluminum hydroxide would require supersaturation 

with respect to the given mineral in terms of dissolved aluminum and 

silica, and this probably most commonly happens only very near the 

mineral surface from which the aluminum and aqueous silica have 

been derived. Complexing of aluminum or the action of a low pH water 

can increase aluminum concentrations, but normally one would expect 

an excess of aqueous silica over aluminum in such systems because 

silicon normally is present in larger amounts than aluminum in the 

source materials.
EFFECTS OF ORGANIC SOLUTES G73

Weathering solutions in which a low concentration of silica is 

maintained, as might occur in a terrane receiving large amounts of 

rainfall, could alter clays, previously formed at the surface of 

weathered primary rock minerals or transported to the are£ in some 

way, to aluminum hydroxide. The literature contains many papers 

suggesting this type of origin for bauxite deposits. The pH in such a 

system is likely to be in a range favoring gibbsite precipitation, as 

dilute runoff waters in humid regions commonly have pH's somewhat 

below neutrality. The critical level for silica concentration probably is 

not as low as some authors imply from consideration of the stability of 

kaolinite. In our experiments the less well crystallized precursor, 

synthetic halloysite, superseded the crystallization of gibbsite only 

when aqueous silica concentrations exceeded about 9 mg/1 as SiCV It 

would seem, therefore, that bauxite might be produced from clays even 

where the dissolved silica in the associated solutions sometimes 

reached this level.

The effect of organic complexing in most weathering environments 

probably would be to increase aluminum solubility and possibly to 

decrease the availability of reactive species. Such effects would tend to 

promote removal of aluminum in the effluent solutions and prevent 

accumulation of aluminum enriched solids. Also, the organic material 

might interfere in the precipitation of aluminum hydroxide.

For the above reasons, one probably should not expect to find 

gibbsite accumulations in environments that were rich in accumulated organic debris, such as lake bottom deposits, swamps, or peat 

bogs. In humid tropical environments much organic matter may be 

produced, but it tends to be decomposed rapidly, rather than 

accumulating in the soil.

ANOMALOUS ALUMINUM CONCENTRATIONS IN NATURAL

WATER

In the absence of complexing effects, the solubility of aluminum in a 

dilute river or ground water would normally be in the vicinity of 10~6 

molar (fig. 17) near pH 7. This is equivalent to about 27 /u,g/l and is near 

the limit of detection for the usual methods of determination used in 

water analysis laboratories. Lower concentrations could be expected 

for runoff waters that have not reached equilibrium with available 

solid sources of aluminum, such as the poorly organized clay minerals 

that commonly occur in soils and streambed sediments. Higher 

concentrations from 100 to 1,000 /u,g/l occasionally appear in reports of 

water analyses published in the literature. Most of these represent 

water with low pH, but some are near neutral and have a yellow or 

brown color. It is of interest to examine a few of the analyses to see if 

they can be readily explained
G74 CHEMISTRY OF ALUMINUM IN NATURAL WATER

Before doing this it should be noted that routinely collected and 

analyzed water samples may not furnish reliable data on minor 

constituents or on solution parameters that may be unstable. Reported 

aluminum values may include particulate material that was not 

separated from solution, and the pH reported may represent a 

laboratory determination that may not indicate the pH ir the water at 

the time of sampling.

Ground water that has a low pH can be expected to have a much 

higher concentration of aluminum than the 10 moler mentioned 

above. Certain types of thermal springs, analyses for which are cited 

by White, Hem, and Waring (1963, p. 44-47), show these properties. 

Generally the aluminum contents of these waters are below saturation 

with respect to aluminum hydroxide or silicate minerals, ^here may be 

other aluminum minerals, however, that are less soluble than these; in 

the low pH, high sulfate environment, for example, alunite may be the 

least soluble species (Roberson and Hem, 1969). Water with a low pH 

may also occur in mine workings as a result of the large scale 

imposition of an oxidizing regime on fresh sulfide mineral surfaces. In 

such environments an equilibrium with kaolinite may t e reached, if 

this mineral is readily available. Hem, Roberson, Lind. and Polzer 

(1973, p. 38-39) were able to demonstrate a near approach to 

equilibrium with kaolinite in water from flooded coal mines in eastern 

Pennsylvania.

Water from acid springs, drainage from mines, and various kinds of 

industrial waste effluent may impart local conditions of low pH and 

high aluminum concentration to surface streams. Upon mixing with 

the main flow of water, supersaturation with respect to synthetic 

halloysite or a similar clay mineral is very probable. However, 

persistence of equilibrium conditions after precipitatior of the solid 

had taken place would be unlikely because of the influence of dilution 

and the many other effects that might occur in such an open system.

A system in which greater-than-expected equilibrium concentrations of aluminum might appear to be present would be one having 

organic material capable of complexing aluminum in solution and a 

supply of aluminum to be complexed. Perhaps more commonly an 

aluminum complex is already present in living herbaceous material 

and becomes available for solution after death or destruction of the 

plant. One could, therefore, postulate the most likely places and times 

for finding high aluminum content in runoff. For example, favorable 

conditions might include low soil pH, abundant vegetatior, and a large 

supply of runoff occurring at a time when fresh organic debi is was 

abundant, as during a rainy period shortly after the annual autumnal 

leaf fall in a deciduous forest.

Some indication of the amount of organic material ir solution in
EFFECTS OF ORGANIC SOLUTES G75

natural water can be gained from the color number determined as a 

part of the routine chemical analysis. The number is only an index of 

the color intensity and is far from being a quantitative measurement of 

organic material, but a highly colored water generally may be 

considered to represent one that has a relatively large content of 

organic material. Runoff from swamps and water in lakes and streams 

of forested regions, especially where topography is relatively flat and 

streams flow sluggishly, and in regions at high latitudes wher-^ tundra 

constitutes the vegetation may be highly colored.

EFFECT OF FILTRATION

The concentration of aluminum reported may represent not only 

that part which is in true solution but also colloidal and particulate 

matter. Many published analyses represent water samples that were 

not filtered before analysis, although it has been usual practice in 

Geological Survey laboratories and in other laboratories which are 

geochemically oriented to perform analyses only on water free from 

visible turbidity. Clarification commonly was by settling during 

storage of samples prior to analysis. Some colloidal particles might 

remain in an apparently clear supernatant, however. The filter pore 

size necessary for a filtrate to represent a true solution is still under 

question (V. C. Kennedy and B. F. Jones, written commun., 1971). One 

can conclude from existing data that a 0.1 -/urn filter can remote from 0 

to 100 percent of the aluminum in a water sample that has passed a 

0.45-|iim filter.

A classification of material according to particle size will give an 

idea of the type of material removed by filtration and the composition 

of the filtrate. Lammers (1968), in his study of fractionation of 

suspended and colloidal substances in natural water, classified 

particles smaller than O.Oljum as dissolved, those from O.Oljiim to 

200)itm as insoluble suspended and colloidal particles, and tho^e larger 

than 200|iim as materials that would normally tend to settle out 

rapidly unless they were free swimming or buoyant organic particles 

or unless the water was very turbulent.

The removal of aluminum from solution by a 0.1- and 0.45-jiim filter 

was somewhat inadvertently tested on the quercetin solutions. After 

about 2 months of aging of these solutions, the total aluminum was 

measured in filtrates that had passed a 0.45-jiim filter. (The 0.45-jiim 

filter is the pore size most widely used for filtration prior tc a water 

analysis.) All the other aliquots representing other aging times were 

filtered through 0.10-jum filters. Only one of the aliquots seemed to 

contain a significant amount of particulate material betweer the 0.10 

and 0.45-jiim diameters.

Kennedy and Jones (written commun., 1971) found that a p^rt of the
G76 CHEMISTRY OF ALUMINUM IN NATURAL WATER

fine particulars (<0.1-1.5-ju,m grain size) in the stream samples they 

analyzed was composed of organic matter containing trace metals. 

Thus, the finest particles in a river water could be in part organic, and 

some of the aluminum reported in excess of that attributable to 

inorganic ions and their complexes could be bound to these organic 

particles besides being in the form of dissolved organic complexes.

EXAMINATION OF ANALYTICAL DATA

It is recognized that natural waters are not necessarily in a state of 

chemical equilibrium. However, it would seem logical that metastable 

solids of the types first observed in aging of laboratory solutions should 

be the first to appear in natural waters. These species probably would 

persist until conditions are favorable for a higher state of crystallization. As in other parts of this report, the mineral species considered are 

microcrystalline gibbsite, microcrystalline trihydroxide (the prebayerite form), and metastable halloysite.

Tables 12 and 13 contain data taken from published sources and 

compare calculated aluminum solubility, assuming no organic 

complexing, with actually observed concentrations fcv 31 water 

samples. The calculated solubilities assume either an equilibrium 

with respect to microcrystalline gibbsite or microcrystalline trihydroxide where silica content was less than 9.0 mg/1 or with respect to 

synthetic halloysite where silica concentration was higher. Inorganic 

complexes are taken into account, but organic ones are not. 

Thermodynamic data are taken from table 3, and the activity 

coefficients are taken from the results of a computer output of the 

program described by Kharaka and Barnes (1973).

The departure from the calculated equilibrium concentration of 

aluminum is substantial in most of the more highly colored waters in 

tables 12 and 13. Probably, however, most of the aluminum is 

associated with particulate matter in these waters and probably could 

have been removed, for the most part, by filtration. This statement is 

based partly on the observed tendency for quercetin-aluminum 

complexes to precipitate at higher concentration and pH and also on 

our extrapolation of the behavior of quercetin to that of naturally 

occurring aqueous organic compounds of high molecular weight. Some 

additional support is available in numerous studies of principles and 

techniques for removal of organic color in the conditioning and 

treatment of water supplies. This removal is commonly accomplished 

by addition of aluminum, pH adjustment to an alkaline condition, and 

filtration. Most investigators have considered this as a colloidal system 

(Stumm and O'Melia, 1968).
EFFECTS OF ORGANIC SOLUTES G77

5Tab £o inW0)

fe

CO

soC

soc8S

a ci3

"3

-c

150 10^ <^ »^a 2

v.

^»"33

t3

2!

-2"oo

-5s "i3§

ts

CO

cu

CO ^ac1s

c4T 1 S«H

c1

*s a

"S-2 2~

Si- C ^2 

0)^2 S W> 11 1- "« 5T <

o ^

h § O p2 o S

0 3CaP. o1

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si ii | 

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 Sc cla g^S'13 ^-^ 01 '3^0 '803 ^^^ 

5. jj t, t. e %*"£ % a ® fe fc S 53 fe-te -^ 

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(£(£ 

!!^^f!HlfiIIi^Bp M02 KPnO JjKOnlXPHMtBO O
TABLE 13. Analyses of naturally colored water with high aluminum 

Location

St. Louis River at Scanlon,

Do. _._ ______

Patuxent River at Benedict 

St. Louis River at Scanlon, 

Little Patuxent River near 

Little Patuxent River at Fo 

Partridge River near Auror 

St. Louis River near Aurort 

Do. ._ ___ _____ .

Do. _ _._ .____._

St. Louis River at Scanlon,

, Md____ . _ . _ ______ 

Minn _ _ _ _ _ _ _ _ _____ _

rt George G. Meade, Md 

a, Minn

Minn _ _. ____ ___ _

Reported

Date

4- 9-65 

&-30-65 

10- 4-63 

10-30-64 

6-27-63 

10- 2-63 

12- 7-64 

12- 4-64 

1- 7-65 

2-26-65 

2-26-65

Aluminum

(All

700 

1,300 

1,000 

700 

100 

400 

1,000 

900 

800 

600 

800

Silica

(SiOj) 

(mg/1)

9.9 

9.7 

2.6 

7.3 

15.0 

11.0 

130 

13.0 

12.0 

14.0 

100

pH

7.10 

7.40 

6.90 

650 

6.40 

6.40 

7.20 

7.10 

6.80 

720 

740

n
EFFECTS OF ORGANIC SOLUTES G79

CONCLUSIONS

Natural waters commonly contain organic material, although 

amounts and species to be expected cannot be accurately predicted 

from existing published information. Aluminum may form stable 

soluble complexes with low molecular weight organic ligands and may 

be associated with relatively insoluble or colloidal particulate forms of 

high molecular weight organic compounds. Association of aluminum 

with high molecular weight organic material present as colloidal 

particles in river water can increase the apparent dissolved aluminum 

content if such particles are not removed from the water before 

analysis.

Concentrations of soluble organic ligands in the vicinity of 10 ~4 

molar probably are near the upper limit to be expected in most stream 

water. The complexing power of such concentrations of organic 

material for aluminum is comparable to that of strongly cor^olexing 

inorganic ligands, such as fluoride. Under favorable conditions organic 

complexes might be capable of increasing aluminum solubility by a 

factor of 100 in part of the pH range of natural water.

The structure of quercetin is similar to some of the raturally 

occurring organic coloring material in natural waters. Laboratory 

experiments with this material show that it forms complexes with 

aluminum. Besides the simple enhancement of solubility and possible 

formation of aluminum-enriched colloidal particles, there are significant effects on the previously elucidated processes of polymerization of aluminum hydroxide macroions and the formation of 

crystalline aluminum hydroxide and aluminosilicates. Such effects 

can be observed even when the concentration of organic material is so 

small that it can complex only a minor fraction of the dissolved 

aluminum. The results of the quercetin experiments may be 

summarized as follows:

1. Quercetin forms a complex with a 1:2 aluminum to quercetin mole 

ratio. The stability constant fa for the reaction

2Q-+A13+=A1Q2 +

is near 10 12 . A complex with the composition Al3Q2n+ also seemed 

to be formed but disappeared from solution on aging.

2. Polymerization of dissolved aluminum hydroxide specie? to form 

gibbsite, observed in previous studies below pH 6.5, was strongly 

inhibited by the presence of small amounts of quercetin. In 

solutions near neutrality and mildly basic, the precipitate formed 

in the presence of quercetin gave a large X-ray diffraction peak 

that might indicate the presence of boehmite. The more alkaline 

solutions ultimately yielded crystalline bayerite. The process of 

formation of the latter material, however, also is inhibited by the 

presence of quercetin.
G80 CHEMISTRY OF ALUMINUM IN NATURAL WATEP.

3. Small amounts of quercetin in solutions duplicating those used 

earlier for synthesis of relatively amorphous clay minerals with 

1:1 aluminum to silica ratios ("synthetic halloysito") improved 

the yield of well-organized small crystals of kaolinite. The bulk of 

the material remained amorphous to X-rays, however, during 

aging periods of 6 months or more at 25°C. Probably in these 

solutions the quercetin inhibited aluminum hydroxide crystallization and by so doing made it possible for the formation of 

Al-O-Si bonding to keep pace with the other aspect of the 

crystallization process.

The aluminum concentrations in a selected group of water analyses 

that reported organic color were in excess of calculated equilibrium 

saturation, considering only inorganic species. These waters probably 

contained aluminum associated with organic matter.

REFERENCES CITED

Bate-Smith, E. C., 1969, Flavonoid patterns in the monocotyledons, in Harborne, J. B.

and Swain, Tony, eds., Perspectives in phytochemistry: Proc. Phytochemical Soc.

Symposium, Cambridge, April 1968, 235 p. 

Behr, Barbara, and Wendt, Harmut, 1962, Fast ion reactions in solutions (I) formation of

the aluminum sulfate complex: Zeitschr. Elektrochemie, v. 66, p. 223-228. 

Birge, E. A., and Juday, C., 1934, Particulate and dissolved organic matter in inland

lakes: Ecol. Mon., v. 4, p. 440-474. 

Black, A. P., and Christman, R. F., 1963, Chemical characteristics of fulvic acids: Am.

Water Works Assoc. Jour., v. 55, no. 7, p. 897-912. 

Bowen, H. J. M., 1966, Trace elements in biochemistry: New York, Academic Press, 241

PBruckert, Sylvain, 1970, Influence des composes organiques solubles sur la pedogenese

en milieu acide: Annales Agronomic, v. 21, no. 4, p. 421-452. 

Butler, J. N., 1964, Ionic equilibrium, a mathematical approach: Reading, Mass.,

Addison-Wesley Pub. Co., 547 p. 

Chenery, E. M., 1948, Aluminum in the plant world, Pt. 1, in General survey in

dycotyledons: Royal Bot. Gardens (Kew, England) Kew Bull. no. 2, p. 173-183. 

Christman, R. F., and Ghassemi, Masood, 1966, Chemical nature of organic color in

water: Am. Water Works Assoc. Jour., v. 58, no. 6, p. 723-741. 

Christman, R. F., and Minear, R. A., 1971, Organics in lakes, in Faust, S. J., and Hunter,

J. V., eds., Organic compounds in aquatic environments: New York, Marcel Dekker,

Inc., p. 119-143. 

Craigie, J. S., and McLachlan, Jack, 1964, Excretion of colored ultra-violet-absorbing

substances by marine algae: Canadian Jour. Botany, v. 42, p. 23-33. 

Das, Bamdev, and Aditya, Sukumar, 1959, Spectrophotometric study of the aluminumsalicylic acid complex: Jour. Indian Chem. Soc., v. 36, p. 473-478. 

Degens, E. T., 1965, Geochemistry of sediments, a brief survey: Englew>od Cliffs, New

Jersey, Prentice-Hall, Inc., 342 p. 

DeKimpe, C. R., 1967, Hydrothermal aging of synthetic aluminosilicate gels: Clay

Minerals, v. 7, p. 203-214. 

 1969, Crystallization of kaolinite at low temperature from aluminosilicic gel:

Clays and Clay Minerals, v. 17, p. 37-38. 

Dutt, N. K., and Sur, Bimares, 1957, Oxalate and malonato complexes of trivalent
EFFECTS OF ORGANIC SOLUTES G81

metals; II, Stabilities of trioxalato complexes of iron, chromium, aluminum, and

gallium: Zeitschr. Anorg. u. Allg. Chemie., v. 293, p. 195-203. 

Duursma, E. K., 1965, The dissolved organic constituents of sea water, in F : ley, J. P.,

and Skirrow, G., eds., Chem. Oceanography, v. 1: London, Academic Press, p.

433-475. 

Fiskell, J. G., 1969, Aluminum is active in Florida soils: Sunshine State Agr. Research

Rept., v. 14, no. 5, p. 3-5. 

Foye, W. O., Baum, M. D., and Williams, D. A., 1967, Stability of metal complexes of

salicylic acid derivatives and analogs III: Jour. Pharmaceutical Sci., v. 56, no. 3, p.

332-336. 

Friel, E. A., Wilmoth, B. M., Ward, P. E., Wark, J. W., 1967, Water resources of the

Monongahela River Basin, West Virgina: Virginia Dept. Nat. Resources, Div.

Water Resources, 118 p. 

Garrels, R. M., and Christ, C. L., 1965, Solutions, minerals, and equilibria: New York,

Harper and Row, 450 p. 

Gjessing, E. T., and Lee, F. G., 1967, Fractionation of organic matter in natural waters

on sephadex columns: Environmental Sci. and Technology, v. 1, no. 8, p 631-638. 

Goina, Teodor, Olariu, Maria, and Bocaniciu, Lucia, 1970, Formation constants of the

aluminum ion chelates with pyrocatechol and pyrocatechol violet: Rev. Roumaine

Chimie, v. 15, p. 1049-1056. 

Greenberg, S. A., and Price, E. W., 1957, The solubility of silica in solutions of

electrolytes: Jour Phys. Chemistry, v. 61, p. 1539-1541. 

Halbert, H. N., Bryant, C. T., and Hines, M. S., 1968, Water resources of Grent and Hot

Springs Counties, Arkansas: U.S. Geol. Survey Water-Supply Paper 1857, 64 p. 

Harborne, J. B., 1972, Evolution and function of flavonoids in plants, in Recent advances

in phytochemistry, v. 4: Proc. 9th Ann. Symposium Phytochemical Soc. North

America: New York, Appleton-Century-Crofts, p. 107-141. 

Harned, H. S. and Hamer, W. J., 1933, The ionization constant of water and the

dissociation of water in potassium chloride solutions from electromotive forces of

cells without liquid junction: Am. Chem. Soc. Jour., v. 55, p. 2194-2206. 

Hem, J. D., 1968, Graphical methods for studies of aqueous aluminum hydroxide,

fluoride, and sulfate complexes: U.S. Geol. Survey Water-Supply Paper 1827-B,33

P- 

 1970, Aluminum, in Wedepohl, K. H., ed., Handbook of Geochemistry, v. II/2: New

York, Springer-Verlag, p. 13-G-l to 13-L-3. 

Hem, J. D., and Lind, C. J., 1974, Kaolinite Synthesis at 25°C: Science, v. 184, no. 4142, p.

1172. [Copyright 1974 by the American Association for the Advancement of

Science]. 

Hem, J. D., and Roberson, C. E., 1967, Form and stability of aluminuir hydroxide

complexes in dilute solution: U.S. Geol. Survey Water-Supply Paper 1827-A, 55 p. 

Hem, J. D., Roberson, C. E., Lind, C. J., and Polzer, W. L., 1973, Chemical interactions of

aluminum with aqueous silica at 25°C: U.S. Geol. Survey Water-Supply Paper

1827-E, 57 p. 

Hemley, J. J., 1959, Some mineralogical equilibria in the system K,O-Al2 O3-SiO2-H0 O:

Am. Jour. Sci., v. 257, p. 241-270. 

Hoefs, Joachim, 1969, Organic geochemistry of carbon, in Wedepohl, K. H., ed.,

Handbook of Geochemistry, v. II/I: New York, Springer-Verlag, p. 6-L-l to 6-L-30. 

Huang, W. H., and Keller, W. D., 1972, Geochemical mechanics for the dissolution,

transport, and deposition of aluminum in the zone of weathering: Clays and Clay

Minerals, v. 20, p. 69-74. 

Huck, M. G., 1972, Impairment of sucrose utilization for cell wall formation in the roots

of aluminum-damaged cotton seedlings: Plant and Cell Physiology, v. 13, p. 7-14. 

Hutchinson, F. E., and Hunter, A. S., 1970, Exchangeable aluminum levels in two soils
G82 CHEMISTRY OF ALUMINUM IN NATURAL WATER

as related to lime treatment and growth of six crop species: Agronomy Jour., v. 62, p.

702-704. 

Job, P., 1928, Formation and stability of inorganic complexes in section: Annales

Chimie, v. 9, p. 113. 

Johnston, K. M., Stern, D. J., and Waiss, Jr., A. C., 1968, Separation of flavonoid

compounds on Sephadex LH-20: Jour. Chromatography, v. 33, p. 539-541. 

Kaurichev, I. S., Nozdrunova, Ye. M., and Yevseyeva, R. P., 1969, Content and forms of

water-soluble aluminum compounds in soil solutions (Translated from Russian):

Soviet Soil Sci., no. 5, p. 547-556. 

Keller, W. D., Pickett, E. E., and Reesman, A. L., 1966, Elevated dehydroxylation

temperature of the Keokuk geode kaolinite a possible reference mineral: Proc.

Internat. Clay Conf., Jerusalem, 1966, v. 1, p. 75-85. 

Kharaka, Y. K., and Barnes, Ivan, 1973, SOLMNEQ; Solution-mineral equilibrium

computations: Menlo Park, Calif., U.S. Geol. Survey Computer Contr., 82 p.,

available only from U.S. Dept. Commerce, Natl. Tech. Inf. Service, Springfield, Va.

22151 as report PB-215 899. 

Kittrick, J. A., 1970, Precipitation of kaolinite at 25°C and 1 atm.: Clays and Clay

Minerals, v. 18, p. 261-267. 

Kroes, H. W., 1972, Extracellular products from chlorococcum ell ;r»soideum and

chlamydomonas globosa: Archiv Mikrobiogie v. 84, p. 270-274. 

Lamar, W. L., and Goerlitz, D. F., 1963, Characterization of carb^xylic acids in

unpolluted streams by gas chromatography: Am. Water Works Asroc. Jour., v. 55,

no. 6, p. 797-802.

 1966, Organic acids in naturally colored surface waters: U.S. Geol. Survey

Water-Supply Paper 1817-A, 17 p. 

Lammers, W. T., 1968, Fractionation of suspended and colloidal particles in natural

water: Union Carbide Corporation Nuclear Div., report no. K-1749, Oak Ridge,

Tenn., 42 p. 

Lee, C. R., 197la, Influence of aluminum on plant growth and tuber yield of potatoes:

Agronomy Jour., v. 63, p. 363-364.

 1971b, Influence of aluminum on plant growth and mineral nutrition of potatoes:

Agronomy Jour., v. 63, p. 604-608. 

Linares, Jose, and Huertas, Francisco, 1971, Kaolinite synthesis at roo"n temperature:

Science, v. 171, p. 896-897. 

Lindholm, G. F., 1970, An appraisal of ground water for irrigation in the Wadena Area,

Central Minnesota: U.S. Geol. Survey Water-Supply Paper 1983, 56 p. 

McCarren, E. F., and Keighton, W. B., 1969, Water quality and discharge of streams in

the Lehigh River Basin, Pennsylvania: U.S. Geol. Survey Water-Supply Paper

1879-H, 48 p. 

Martell, A. E., 1971, Principles of complex formation, in Faust, S. J., and Hunter, J. V.,

eds., Organic compounds in aquatic environments: New York, Marcel Dekker, Inc.,

p. 239-263. 

Napoli, Aldo, 1972, Aluminum complexes with diglycolic and thiodiglycilic acids: Jour.

Inorganic Nuclear Chemistry, v. 34, p. 1225-1231. 

Nelson, K. H. and Lysyj, Ihor, 1968, Organic content of Southwest and Pacific Coast

municipal waters: Environmental Sci. and Technology, v. 2, no. 1, p. 61-62. 

Pauli, F. W., 1966, Some recent developments in biogeochemical research: Bull. Service

Carte Geol. Alsace-Lorraine, v. 19, no. 3-4, p. 221-240. 

Pecci, Joseph, and Foye, W. O., 1960, The avidity of salicylic, gentisic, and salicyluric

acids for many heavy metal cations: Jour. Am. Pharmaceutical Sci., v. 49, p. 411. 

Porter, L. J., and Markham, K. R., 1970a, The aluminum (III) complexes of

hydroxyflavones in absolute methanol; Part 1, Ligands containing only one

chelating site: Chem. Soc. [London] Jour., sec. C, no. 2, p. 344-34JT.

 1970b, The aluminum (III) complexes of hydroxyflavones in absolute methanol;
EFFECTS OF ORGANIC SOLUTES G83

Part II, Ligands containing more than one chelating site: Chem. Soc. [London] Jour., 

sec. C, no. 9, p. 1309-1313. 

-1970c, The unsuitability of ethanol as a solvent for the spectroscopic detection of

functional groups in hydroxyflavones with aluminum chloride: Phytochemistry,

v. 9, p. 1363-1365. 

Rashid, M. A., 1972, Role of quinone groups in solubility and complexing of metals in

sediments and soils: Chem. Geology, v. 9, no. 4, p. 241-247. 

Roberson, C. E., and Hem, J. D., 1969, Solubility of aluminum in the presence of

hydroxide, fluoride, and sulfate: U.S. Geol. Survey Water-Supply Paper 1827-C, 37 p. 

Robinson, Trevor, 1967, The organic constituents of higher plants [2d ed.]: Minneapolis,

Minn., Burgess, 306 p. 

Saxby, J. D., 1969, Metal-organic chemistry of the geochemical cycle: Rev. Pure and

Appl. Chemistry, v. 19, p. 131-150. 

Schnitzer, Morris, 1971, Metal-organic matter interactions in soils and waters, in Faust,

S. J., and Hunter, J. V., eds., Organic compounds in aquatic environirents: New

York, Marcel Dekker, Inc., p. 297-315. 

Schnitzer, Morris, and Skinner, S. I. M., 1969, Free radicals in soil humic compounds:

Soil Sci., v. 108, no. 6, p. 383-390. 

Schoen, Robert, and Roberson, C. E., 1970, Structures of aluminum hydroxide and

geochemical implications: Am. Mineralogist, v. 55, p. 43-77. 

Seikel, M. K., 1964, Identification of phenolic compounds, in Harbome, J. B.,

Biochemistry of phenolic compounds: New York, Academic Press, p. 618. 

Sillen, L. G., and Martell, A. E., 1964, Stability constants of metal-ion complexes: Chem.

Soc. [London] Spec. Pub. 17, 754 p. 

Skopintsev, V. A., and Bakulina, A. G., 1966, Organic matter in the Rybinsk Reservoir

in 1964, in Shtegman, B. K., ed., Production and circulation of organic matter in

inland waters: Inst. Biologii Vnutrennikh Vod., v. 16, no. 13, Moscow, U.S.S.R., p.

1-31. 

Smith, R. W., 1971, Relations among equilibrium and nonequilibrium aqueous species of

aluminum hydroxy complexes, in Nonequilibrium systems in natural water

chemistry: Am. Chem. Soc., Advances in Chemistry Ser., v. 106, p. 250-279. 

Smith, R. W., and Hem, J. D., 1972, Effect of aging on aluminum hydroxide complexes in

dilute aqueous solutions: U.S. Geol. Survey Water-Supply Paper 1827-D, 51 p. 

Stevenson, F. J., 1972, Role and function of herbicides and chelation of micronutrients:

BioScience, v. 22, no. 6, p. 643-650. 

Strecher, P. G., Finkel, M. J., Siegmund, O. H., Szafranski, B. M., I960, The Merck Index

of Chemicals and Drugs [7th ed.]: Railway, N. J., Merck & Co., Inc., 8?9 p. 

Stumm, Werner, and O'Melia, C. R., 1968, Stoichiometry of coagulation: Am. Water

Works Assoc. Jour., v. 60, p. 514-539. 

Summer, M. E., and Meyer, J. H., 1971, Incidence of toxic aluminum in sandy soils: Proc.

South African Sugar Technologists Assoc., June 1971, p. 1-5. 

U.S. Geological Survey, 1967, Quality of surface waters of the United States 1963: U.S.

Geol. Survey Water-Supply Paper 1947, 472 p. 

 1969, Quality of surface waters of the United States, 1964: U.S. Geol. Survey

Water-Supply Paper 1954, 606 p. 

-1970, Quality of surface waters of the United States, 1965: U.S. Geol. Survey

Water-Supply Paper 1962, 480 p. 

Vosburgh, W. C., and Cooper, G. R., 1941, Complex ions; I, The identification of complex

ions in solutions by spectrophotometric measurements: Am. Chem. Soc. Jour., v. 63,

p. 437-442. 

White, D. E., Hem, J. D., and Waring, G. A., 1963, Chemical composition of subsurface

waters, in Data of Geochemistry [6th ed.]: U.S. Geol. Survey Prof. Paper 440-F, 67 p. 

Whitehead, D. C., 1964, Identification of p-hydroxybenzoic, vanillic, p-covmaric, and

ferulic acids in soils: Nature, v. 202, p. 417-418.

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