Chemical Interactions of 

Aluminum with 

Aqueous Silica at 25°C

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1827-E
Chemical Interactions of 

Aluminum with 

Aqueous Silica at 25°C

By J. D. HEM, C. E. ROBERSON, C. J. LIND, and W. L. POLZER

CHEMISTRY OF ALUMINUM IN NATURAL WATER

GEOLOGICAL SURVEY WATER-SUPPLY PAPER 1827-E

An evaluation of the chemical properties 

of colloidal clay like material formed by 

mixing solutions of aluminum and silica

UNITED STATES GOVERNMENT PRINTING OFFICE, WASHINGTON : 1973
UNITED STATES DEPARTMENT OF THE INTERIOR 

ROGERS C. B. MORTON, Secretary

GEOLOGICAL SURVEY 

V. E. McKelvey, Director

Library of Congress catalog-card No. 72-600332

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

Washington, D.C. 20402 - Price 70 cents

Stock Number 2401-00229
CONTENTS

Page

Abstract----_--_--__-______-______________________________________ El

Scope and purpose,---__------__-__________________________________ 1

Previous work.-___________________________________________________ 2

Theoretical considerations,_________________________________________ 5

Preparation of aluminum hydroxide polymer._________________________ 8

Introduction of silica_ _____________________________________________ 9

Composition of aged solutions_______________________________________ 10

Tests for equilibrium_______________________________________________ 12

Effect of dissolved-silica activity..___________________________________ 16

Reactions in alkaline solutions.__ ___________________________________ 16

Stability fields and alu minum solubility graphs ________________________ 20

Effects of fluoride________________________________________________ 23

Application to natural-water composition.____________________________ 33

Chemical composition of precipitates- _______________________________ 40

Crystallinity of solids ______________________________________________ 40

Validity of equilibrium model_______________________________________ 47

References cited___________________________________________________ .55

ILLUSTRATIONS

Page 

FIGURES 1-4. Graphs showing:

1. Log of activity of uncomplexed aluminum versus pH

in aged solutions containing silica._____________ E13

2. Sums of logs of activities of uncomplexed aluminum

and aqueous silica versus pH in aged solutions--- 14

3. Log of activity of A1(OH) 4~ versus pH in aged solutions containing silica and solutions containing 

crystalline bayerite___________________________ 18

4. Sum of logs of activities of Al(OH),r and undissociated aqueous silica versus pH in aged solutions _ 19 

5-9. Diagrams showing stability fields of:

5. Synthetic halloysite and microcrystalline gibbsite and 

solubility of aluminum as functions of pH and 

activity of undissociated aqueous silica at 25°C 

and 1 atmosphere. Ionic strength 0.01; fluoride 

absent._____________________________________ 21

6. Synthetic halloysite and microcrystalline gibbsite 

and solubility of aluminum as functions of pH and 

activity of undissociated aqueous silica at 25° C 

and 1 atmosphere. Ionic strength 0.10; fluoride 

absent_____________________________________. 22

in
IV CONTENTS

FIGURES 5-9. Diagrams showing stability fields of Continued Page

7. Synthetic halloysite and microcrystalline gibbsite 

and solubility of aluminum as functions of pH and 

activity of undissociated aqueous silice, at 25° C 

and 1 atmosphere. Ionic strength 0.01; total 

fluoride concentration 0.19 mg/1 as F___________ E24

8. Synthetic halloysite and microcrystalline gibb&ite 

and solubility of aluminum as functions of pH and 

activity of undissociated aqueous silic?, at 25° C 

and 1 atmosphere. Ionic strength 0.01; total 

fluoride concentration 1.9 mg/1 as F____________ 25

9. Cryolite, synthetic halloysite, and microcrystalline 

gibbsite and solubility of aluminum as functions of 

pH and concentration of dissolved silica at 25°C 

and 1 atmosphere. Ionic strength 0.10; total 

fluoride concentration 190 mg/1 as F; sodium concentration 2,300 mg/1 as Na___________________ 27

10. Graph showing relationship of log [Al] or log [A1(OH) 4 ] + 

 log [H4 SiOJ to pH for selected water samples from 

ground- and surface-water sources ____________________ 38

11-23. Electron micrographs of :

11. Microcrystalline gibbsite synthesized at 25°C______ 42

12. Crystalline bayerite from solution 14E____________ 43

13. Fragments of Lewistown, Mont., standard kaolinite. 44

14. Montmorillonite (Arizona bentonite)_____________ 45

15. Solid recovered from solution 12E________________ 46

16. Solid recovered from solution l7D-._---_-________ 47

17. Solids recovered from solution 17A_______________ 48

18. Solids recovered from solution 17C_____________ 49

19. Solids recovered from solution 11J_ ______________ 50

20. Solid recovered from solution 4D_________________ 51

21. Solids recovered from solution 18A_______________ 52

22. Solids recovered from solution 18B_-_____________ 53

23. Solids passing 0.45/-m filter caught on O.lOjum filter, 54

TABLES

Page 

TABLE 1. Gibbs free energies of formation at 25°C and 1 atmosphere- _____ EG

2. Equations used for calculation of solubilities at 25° C and 1 atmosphere pressure_ _____________________________________ 6

3. Experimental solutions in which mole ratio OH :A1 wa,s less than

3.0._-____----_____-__________________-_-__-------_--- 11

4. Experimental solutions in which mole ratio OH :A1 was greater

than3.0 _-______-__-___-_---__-_----------_----------- 12

5. Composition of solutions A-F after 16 months aging- __________ 30

6. Ion-activity products for solutions A-F after 16 months aging _ _ _ 31

7. Analytical data from actual water samples. _____________--_--- 36
CHEMISTRY OF ALUMINUM IN NATURAL WATER

CHEMICAL INTERACTIONS OF ALUMINUM 

WITH AQUEOUS SILICA AT 25°C

By J. D. HEM, C. E. ROBERSON, C. J. LIND, and W. L. POLZER

ABSTRACT

Solutions containing from 10~5 to 10~2 moles per liter of aluminum and dissolved 

silica in various ratios were aged at pH levels between 4 and 10 at 25°C. A colloidal 

amorphous product having the composition of halloysite was produced in most 

solutions. It had a consistent and reversible equilibrium solubility equivalent to a 

standard free energy of formation of 897 ±1.0 kcal per mole for the for-nula 

Al2 Si2 O5 (OH)4. Some aging times were longer than 4 years, but most solutions 

gave consistent solubilities after only a few months of aging. Where silica concentrations were below about 10~4 molar, microcrystalline gibbsite was formed below 

pH 6.0 and crystalline bayerite above pH 7.0, but only after much longer s ging 

than was required for crystallization in silica-free solutions.

Electron micrographs and diffraction patterns of the synthesized material 

indicate some crystallinity in the aluminosilicate, but no X-ray diffraction patterns could be obtained even in the material aged 4 years.

Solubility relationships for solutions containing fluoride as well as silica and 

aluminum are explainable by using cryolite stabilities determined in previous 

work.

Aluminum contents of 51 samples of water analyzed for other purposes are in 

reasonable agreement with the assumption of equilibrium with amorphous clay 

mineral species similar to the material synthesized in this work. Solubility 

calculation? are summarized graphically for solutions of ionic strength of 0.01 

and 0.10.

SCOPE AND PURPOSE

Framework structures of various types in which oxygen ions are 

bonded to silicon and aluminum cations are present in most common 

rock minerals. When such structures are attacked by water in processes of weathering, the frameworks will be extensively rearranged, 

and both aluminum and silicon will become available for solution, 

along with other cations present in the structure. The aluminum and 

silicon contents of natural water must depend upon the rates at which 

these elements become available for solution and how they are recombined into clay minerals or hydroxides. This recombination seems 

to, be a particularly important natural chemical control of aluminum 

solubility, but it is not well understood.

El
E2 CHEMISTRY OF ALUMINUM IN NATURAL WATER

The studies described here were made to obtain a better understanding of the behavior of aqueous aluminum species in the presence 

of aqueous silica. The techniques used were similar to those used 

previously for studying the chemistry of dilute solutions of aluminum. 

Silica was added in various amounts to test solutions that contained 

aluminum, and the results were compared with observe-tions made in 

the absence of silica. The resulting hypotheses can be tentatively 

applied to analyses of natural water. The information available on 

aluminum contents of natural water is inadequate (1) in the number of 

reliable published concentrations and in specifying whether reported 

concentrations represent dissolved or suspended material and (2) in 

the form in which the aluminum solute occurs. One aim of this study 

has been to provide a theoretical basis which will help interpret existing information on aluminum content of natural water. The study also 

indicates some possible controls over the dissolved-silica concentrations in natural water.

Although the investigations shed some light on the composition and 

stability of certain synthetic aluminosilicate minerals which seem to 

be related to natural clay minerals and are probably similar to natural 

metastable weathering products, the synthesis of clays was not a 

primary purpose of the study. As many other investigators have 

reported, the synthesis of well-crystallized clay minerals is very 

difficult to accomplish at 25°C in laboratory-type systems.

The authors' interpretation of the experimental results summarized 

in this paper have benefited from extensive discussions with colleagues 

and from their reviews of this paper. In addition to assistance from 

members of the U.S. Geological Survey, reviews and suggestions were 

received from Professors J. O. Leckie and G. A. Parks, Stanford 

University. W. L. Polzer, U.S. Atomic Energy Commission, during his 

employment by the Geological Survey, performed a considerable part 

of the experimental work that is described in this pape1'.

PREVIOUS WORK

This report represents a continuation and uniting of two previous 

lines of research conducted by the authors and their associates. These 

studies concerned the chemistry of aluminum in dilute aqueous solutions and the solubility of the clay mineral kaolinite. Reports summarizing this work were written by Hem and Roberson (1967), Hem 

(1968), Roberson and Hem (1969), Smith and Hem (1972), Polzer 

and Hem (1965), and Schoen and Roberson (1970). Mechanisms of 

formation of gibbsite, bayerite, and nordstranclite (forms of A1(OH) 3 ) 

appear to be inhibited or altered and aluminum solubility appears 

substantially decreased by the presence of silica. Polzer and Hem
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E3

(1965) indicated that an apparent equilibrium solubilit}' could be 

determined for kaolinite at a pH near 3.5 but that more stud}' of 

mechanisms of solution and precipitation was needed. The kaol: nite 

used by Polzer and Hem was impure, and the results had to be interpreted to allow for that complicating factor.

The work described in this paper represents an approach to studies 

of hydrous aluminosilicate minerals through synthesis by precipite tion 

from supersaturated solutions. The work of Polzer and Hem (1965) 

essentially consisted in trying to attain saturation in initialh7 undersaturated solutions.

Literature that concerns the fundamental chemistries of aluminum 

and silicon, geochemical literature on weathering reactions, mireral 

syntheses and natural-water compositions, and some applied chemistry 

in such fields as colloidal behavior and aluminum Irydroxide coagulation 

as applied in water treatment may be of some interest but will not 

be fully reviewed here. An extensive literature exists relating to the 

synthesis of clay minerals in the laboratory. Only a few investigators 

of this subject, however, have paid much attention to reaction 

paths or rates, and most workers have relied on high temperature 

and pressure to bring about the desired results.

It is well known that kaolinite with well-defined crystalling can 

be produced by subjecting feldspar to aqueous solutions at high 

pressure and temperature. For example, Hemle}T (1959) subjected 

potassium feldspar and water to a temperature of 300°C and a pressure 

of 15,000 pounds per square inch to produce kaolinite and other 

alteration products. That a crystallized clay mineral can be formed 

from amorphous aluminosilicate gels by heat (200°-400°C) and pvessure has been known from the work of various investigators as far 

back as the 1930's. Kaolinite is a common accessory species in mineral 

veins, and relative^ high temperature and pressure environments 

obviously can produce this and other varieties of clay. The clay 

minerals are by no means restricted to such environments, but they 

commonly seem to be formed in place during the normal surficial 

weathering of many igneous rock minerals. Attempts to produce 

clays in the laboratorjr at 25°C have generally met with complete 

failure, although during the 1960's a few studies reported limited 

success in synthesizing solid species having clay mineral characteristics identifiable by electron microscopy or other sensitive methods. 

Yields generally were small and reactions very slow, and the crystallinity seemed to be poor unless higher temperatures were imposed 

on the system in some part of the synthesis.

Field observation commonly associates kaolinite, and similar clays 

having equimolar aluminum and silicon contents, with weathering 

regimes in which leaching is relative^ rapid, weathering solutions are
E4 CHEMISTRY OF ALUMINUM IN NATURAL WATER

dilute, and soil pH is low. Montmorillonite and other more siliceous 

clays appear to occur where drainage is poorer, water is less abundant, 

and dissolved solids and pH are relatively high. These relationships 

evidently have been considered in the design of some of the more 

recent experiments attempting synthesis of kaolinite and other clays 

at 25° C. A series of experiments in Germany (Harder, 1965, 1966, 

1967) in which solutions were generally between 10~3 p.nd 10~5 molal 

with respect to silica and aluminum were reported to yield aluminosilicates amorphous to X-rays. On aging at temperatures ranging from 

3° to 60° C, several of the precipitates developed crystalline quartz. 

Kaolinite, chlorite, and probably illite and feldspar vrere identified 

in material aged for 30 days at 60° C.

Much highly significant recent work has been done elsewhere in 

Europe, especially by investigators at the University of Louvain in 

Belgium. DeKimpe (1967, 1969), for example, described syntheses of 

clay at 175° C, starting with aluminosilicate gel. Related studies were 

described by DeKimpe, Gastuche, and Brindley (1961) and in many 

other published papers by members of the research groups at Louvain. 

In general the technique used by these workers has been to combine 

solutions of organically complexed aluminum (for example, aluminum 

isopropoxide) with an organically complexed form of orthosilicate (for 

example, ethyl orthosilicate). The organic portions of these species 

react with each other, and relatively high concentrator s of aluminum 

and silica can be brought together in this way without introducing 

large amounts of other inorganic ions. Gels of various compositions 

which can be made to crystallize on aging are produced.

Siffert (1962) reported syntheses of several species of magnesium 

clays and of kaolinite at room temperatures, using a technique somewhat similar to that just described. Other recent significant studies of 

low-temperature aluminosilicate synthesis and related solubilities 

have been made by Kittrick (1966, 1967) and Reesman, Pickett, and 

Keller (1969).

Kittrick (1970) precipitated kaolinite with a relatively well defined 

crystallinity at 25° C in solutions supersaturated with respect to 

kaolinite, in which montmorillonite was suspended at the start of the 

experiment. Reaction time in this experiment was 3-4 years.

Linares and Huertas (1971) reported a synthesis of kaolinite at 

25° C, starting with an aluminum-fulvic acid complex and silicic acid 

in dilute solutions.

A general discussion of coagulation processes by Stumm and O'Melia 

(1968) contains some data of interest relating to interactions between 

aluminum hydroxide and silica and gives many references to work of 

others.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E5

THEORETICAL CONSIDERATIONS

The simplest mineral structures which might be formed by raising 

the pH of solutions containing aluminum and silicic acid are aluminum 

hydroxide in its various forms and arrangements where both silicon 

and aluminum are coordinated with hydroxide, or possibly oxygen ions 

or water molecules. Because of differences in ionic radius and charge, 

aluminum is generally octahedrally coordinated with six of tl^se 

species, and silicon is tetrahedrally coordinated with four. Hence, eubstitution of silicon for aluminum within the gibbsite structural Iryer 

seems unlikely. The silicon may, however, form a tetrahedral structure 

with oxygen that in effect attaches an oxygen layer to the double Iryer 

of hydroxide in the structure of gibbsite. Each silicon ion associated 

with this third layer must displace a proton from an OH ion in the 

gibbsite structure in order to maintain a charge balance. Thus, twothirds of the ions in the middle layer are oxygens and the remaining 

one-third are hydroxides. Structural formulas for this type of clay 

mineral (kaolinite or halloysite are common examples) are given in 

textbooks and elsewhere and will not be repeated here. The compositional formula for kaolinite is Al2Si2O5 (OH) 4 . Halloysite has essentially 

the same structure, but its formula usually is written with some additional water. It can be represented by the kaolinite formula for stoichiometric or equilibrium computations.

A convenient starting point for the development of a conceptual 

model is the knowledge already obtained for the structure and synthesis of gibbsite. Introduction of aqueous silica at the stage of gibtsite 

synthesis where the ordered structure is forming should give the opportunity for synthesizing a 1:1 clay mineral, and some prelimiranr 

electron micrographs published by Polzer, Hem, and Gabe (1967) 

show what appear to be halloysite particles synthesized by aging such 

mixed solutions at 25°. Gibbsite might also form in these solutions. 

The ultimate equilibrium condition ought to represent a system in 

which equilibrium exists among the participating aluminum, silicon, 

and hydrogen ionic species and the solids that are formed.

Table 1 presents Gibbs free energies of species, and table 2 contains 

mass-law and solubilit}^ expressions that are needed for calculating 

the equilibrium relationships of interest. Most of the free energy and 

equilibrium constant values were taken from published literature. 

The free energy of the silicate ion was computed by using the dissociation constant given for silicic acid and the free energies for other 

species involved. Some of the equilibrium constants were determined 

experimentally by procedures that will be described in this paper.

476-920 O - 73 - 2
E6 CHEMISTRY OF ALUMINUM IN NATURAL WATER

TABLE 1. Gibbs free energies of formation at 25°C and 1 atmosphere

Species AG° (kcal per 

mole)

Source of data

Al+3 _______-_.__________________ 115.0 Latimer (1952).

AlOH+2________________________ -164.9 Raupach (1963).

Al(OH) 2+_______________________ -215. 1 Do.

Al(OH)j (c):

microcrystalline gibbsite______ 272. 3 Hem and Roberson (1967).

gibbsite-___________________ -273.9 Latimer (1952).

bayerite_-_-_________.______ -274.0 Hem and Roberson (1967).

nordstrandite.______________ 274. 0 This report.

Al(OH) 4 -_______________________ -311.7 Hem and Roberson (1967).

H4SiO4 (aq)---_-___________.____ -312.7 Polzer and Hem (1965).

H3 SiO4 -________________________ -299.4 Calculated. (See text.)

SiO2 (c), a-quartz_______________ -204.8 Wagman and others (1968).

SiO2 (c), amorphous glass. ________ 203. 3 Do.

Al2Si2 O5 (OH) 4 :

kaolinite...._____.__________ 902.9 Robie and Waldbaum (1968).

halloysite _________________ - 898. 4 Do.

synthetic halloysite__________ 897. 0 This report.

Na3AlF6 , cryolite__'______________ 745.4 Roberson and Hem (1969).

F-.____________________________ -66.08 Latimer (1952).

TABLE 2. Equations used for calculation of solubilities at 25°C o,nd 1 atmosphere

pressure

Mass action equations References for constant

1. [A1+3 ][OH-] 3 = 10-32 - 65 _________ ___ _ _________ Hem and Roberson

(1967), p. 46).

[A1OH+21[H+1

2. = 10~5 - 00 _ __________________________ Smith and Hem

LAJ J (1972). 

^=10-'-w_--_----__--------------- Do.

H L= iQ-22.07 _ _ _ ______ Calculated. IAITO J

o. [H+][OH-] = 10- 14 - 00 _________ Harred and Hamer

(1933, p. 2198).

6. [A1+3]|tH*SiO40] 2 = 1011 - 28 --------------------------- This report.

8- [Hrw^-n[o?+] = 10" 9 - 77 - -------------------------- Greenberg and Price [H4SiO4°] (1957).

9- r4§^,L = 103 - 17 _______________-________-------- Butler (1964, p. 113).

10. -At±^ ^^io?.oi __________________________ Hem (1968, p. 7).

11 [AIF-+] = 1Q12 .75 DO. [F-] 2 [A1+3 ] ----------------------------
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E7

TABLE 2. Equations used for calculation of solubilities at 25° C and 1 atmosphere

pressure Continued

Mass action equations References for constant

12' [F^^I=1° 17 - 02 - - - Hem (1968, p. 7).

13 [All 4 J _jfl1972 ,-j [F-] 4 [A1+3 ] U ------------------------------ i;o.

[A1F5 -] ^20.91 J-)

 [F-]«[A1+8] ----------------------------- uo.

15 [Albs ] 1 Q20 . 86 j-v [F-RA1+3 ] ------------------------------- uo.

16. [Na+]3 [Al+3 ][F-]6 = l(r33 - 84 _ ________________________ Roberson and Hem

(1969, p. 3). 

Mass balance equations

, [A1(OH) 2+] , [A1(OH) 4 -] [A1F+2 ]

it. OAI = ~~i ~T~ ~T~ ~T~ 

7A1+3 7A1OH+2 7A1(OH) 2+ TA1(OH) 4 - 7A1F+2

, [H4Si04] , [H3 SiQ4 -] 18. Cs 

[A1F2+] [A1F,°] [A1F4 -] [A1F5-2 ] [A1F6~3 ]

7A1F2+ 7A1F3 ° 7A1F4~

7H4 8i0 4 7H3 8i04 ~

1Q [F-] [HF] [A1F+2 ] 2[A1F2+] 3[A1F3 ] 4[A1F4~] 5[A1F5 -2 ] 6[A1F5~3 ]

iy. O F= | -| - h -| h I ; I 

TF~ THF TAIF+- 7AiF2+ TAiF3

Relationships among the various species at equilibrium can be 

conveniently represented b}T stabilit}^-field diagrams, based on thermodynamic calculations. In the diagrams presented here the activit}^ of 

aqueous silica, considered as Si(OH) 4 or H4SiO4 , is used as ordinate, 

and hydrogen-ion activit}^, pH, as abscissa. Aluminum activity in 

solution at equilibrium is shown by contours plotted on the stability 

fields of the solids.

Stability-field diagrams such as those used by Feth, Roberson, and 

Polzer (1964) have been rather widely used for representing the 

domains of gibbsite, clays, and other minerals in terms of dissolved 

silica and the ratio of certain cations to hydrogen-ion activity. The 

calculations on which the diagrams are based assume that aluminum 

is conserved and that none of the chemical equilibria treated contain dissolved aluminum as a variable. It seems obvious, however, that soHite 

species of aluminum do occur; otherwise, it is difficult to see how 

changes of one form of solid to another could take place in an aqueous 

system.
E8 CHEMISTRY OF ALUMINUM IN NATURAL WATER 

If the

reaction

,8 CHEMISTRY OF ALUMINUM IN NATURAL WATER

If the synthesis of a clay mineral is represented by the summarizing 

;action

2Al(OH) a (c) + 2H4SiO4 (aq) = Al2Si2Os (OH) 4 +5H2O,

one may calculate from data in table 1 the value of Kicr the resulting 

equilibrium:

#=[H4SiO4] 2 .

The other species in the original reaction do not appear here because 

their activities are unity by definition. The boundary between the 

aluminum hydroxide field and the clay mineral field thus is only a 

function of aqueous silica concentration. Along this line the system 

is in simultaneous equilibrium with two solids and has only one degree 

of freedom.

This relationship provides a simple starting point for preparation 

of an equilibrium solubility diagram for aluminum in the presence of 

silica. The equilibrium constant to be used, however, poses somewhat 

of a problem because it depends on the free-energy values selected for 

the solid species. Normally in depicting a final equilibrium situation 

one might select the most negative free energies of solids involved, 

assuming these values are believed reasonably free from errors caused 

by impurity and disorder in the solids. One might reason that the less 

negative free energies would represent metastable material that would 

alter on aging to a final lower equilibrium value. In this instance, 

however, both the previous laboratory work with aluminum hydroxides 

and field observations of clay mineral characteristics suggest that the 

most stable species are not readily synthesized at ordinary temperature 

and pressure, and therefore a diagram that depicts natural systems 

realistically should be based on free energies for less stable forms. 

In the absence of solute silica, the precipitated alumirum hydroxide 

forms micro crystalline gibbsite at first, with a free energy of 272.3 

kcal per mole. An aluminosilicate species less stable than kaolinite 

seems likely as well; therefore, some laboratory experiments were conducted to try to evaluate the behavior of synthetic aluminosilicates and 

give a logical basis for assignment of a free-energy value.

PREPARATION OF ALUMINUM HYDROXIDE POLYMER

In previous work (Hem and Roberson, 1967) experimental solutions 

containing aluminum hydroxide species were prepared by a technique 

of mixing measured volumes of three stock solutions. The same procedure was generally followed here. The first of the three solutions 

contained aluminum, sodium, hydrogen, and perchlorate ions, with an 

aluminum concentration 4.5X 1CT4 molar, a pH near 2.5, and a perchlo-
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E9

rate concentration of 10~2 molar. The second stock solution was made 

up of sodium, hydroxide, and perchlorate, with a pH near 11.0 and the 

same perchlorate concentration as solution 1. The third solution was 

the same as solution 1 except that it had 9.0X 10~4 moles per liter of 

aluminum and was less acid (pH near 3.5). Aliquots of solution 1 v^ere 

put in polyethylene containers, and selected volumes of solution 2 vere 

added to each to provide different quantities of OH ions. At the same 

time a volume of solution 3 equal to the amount of solution 2 used was 

added so as to maintain a constant total aluminum concentration. This 

technique makes possible the preparation of a series of test solutions 

in which all are alike except for concentrations of free and bound OH 

ions. Successive series of such solutions can be prepared to investigate 

the effect of other factors where the starting quantity of OH is the 

same. The ionic strength and total content of aluminum are essentially 

the same throughout, which simplifies some of the necessary calculations. The amounts of free and bound OH can be calculated from 

measurements of pH, and the composition and proportions of solutions 

1,2, and 3 can be used.

Some solutions also were made in which the ionic strength was 

higher or lower (range 0.1-0.001) by using larger or smaller amounts 

of sodium and perchlorate. In addition, some experiments were irade 

in which aluminum and all other concentrations were increased by 

a factor of 10 over the levels used by Hem and Roberson (19^7). 

The calculations of bound OH were checked for some solutions by 

blank titrations in which all conditions were the same except that 

aluminum in the solutions was replaced by an equivalent concentration 

of sodium. The calculated value for the ratio of moles of bound OH 

to total moles of aluminum per liter, termed r in our earlier work, is a 

convenient parameter for characterizing these solutions. Data for 

solutions in which r<3.0 are listed separately from those in which 

r>3.0 for convenience in discussing the results.

The value of r remains nearly constant during the aging of the 

mixed solutions, but the pH tends to shift downward somewhat in 

most solutions. These changes are related to alteration of the OH 

ions from nonbridgiiig to bridging positions in the polymer and to 

changes in distribution of ionic hydroxy-aluminum complexes (Hem 

and Roberson, 1967; Smith and Hem, 1972).

INTRODUCTION OF SILICA

Silica was introduced into the test solutions by one of two different 

techniques. For some experiments a fourth stock solution was used 

to add silica. This stock solution was prepared from sodium silicate 

reagent (Na2SiO3 -9H2O) which was dissolved in demineralized, distilled 

water and passed through a cation-exchange column (IR 120) to
E10 CHEMISTRY OF ALUMINUM IN NATURAL WATER

replace sodium with hydrogen. The ionic strength of this solution was 

brought to the same level as the other three stock solutions with 

sodium perchlorate, and it had a final pH near 5.0. In other experiments 

the silica was introduced simply by dissolving the desired weight of 

sodium silicate or silicic acid reagent in the basic solution 2. This 

technique avoided dilution of the aluminum concentration in the 

mixed solutions but introduced the silica as dissociated silicate 

(SiO3 ~2 or related species) ions rather than as H4SiO4 . The first procedure was used in the first part of the study and onhT for solutions 

whose final pH was below neutrality.

The molar proportion of silicon to aluminum in the experimental 

aging solutions ranged from about 1:10 to near 5:1. Concentrations 

of silica ranged from near 10~5 to about 3X 10~3 molar, or from a 

little less than 2 mg/1 to about 200 mg/1.

COMPOSITION OF AGED SOLUTIONf

After preparation the solutions were stored in polyethylene containers at or near 25° C in a controlled atmosphere box from which 

carbon dioxide was excluded and were agitated occasionally to insure 

adequate mixing. From time to time a portion was removed from 

each solution, and its pH determined. Part of the aliquot was then 

filtered through a 0.10/mi porosity filter, and the silica and aluminum 

contents were determined on the filtered and sometimes also on the 

unfiltered fractions. Material caught by the filter wa? examined by 

X-ray diffraction and by the electron microscope or electron diffraction; some residues were analyzed quantitatively by wet chemical 

procedures. The aging times were as much as 4 years for a few of the 

solutions, but an equilibrium or steady state appeared to be closely 

approached in many of them after as short a time as 3 or 4 months; 

few significant changes occurred after 1 year in solutions where 

silica was added as silicate ions.

Previous studies have shown that it is important to measure as 

many properties of systems of this kind as possible, and thereby 

decrease the number of assumptions and uncontrolled or unknown 

variables. An important factor is the filtration to remove participate 

material. The 0.10/mi porosit}^ filters generally passed solutions very 

slowly, for a filter cake built up soon after the start of the filtration 

process. This cake probably lowered the effective pore size sufficiently 

to decrease the particulate material in the filtrate to an insignificant 

concentration. Smith and Hem (1972) showed that polynuclear 

aluminum hydroxide species react slowly with ferron, and colloidal 

gibbsite is almost completely unreactive. Species having these characteristics could not be detected in the filtrates.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA Ell

Initial and final aluminum and silica concentrations in the filtrates 

for 56 solutions are given in tables 3 and 4. These tables also give 

some information on apparent composition of material removed by 

filtration. This was calculated from the difference between the initial 

solution composition and the final composition observed in the filtrate 

and probably is not very accurate. These calculations indicate, 

nevertheless, the approximate aluminum to silica ratio in the precipitates and show which of the precipitates were the most enriched in 

silica. Some of the precipitates were probably mixtures of amorphous 

aluminum hydroxide and aluminosilicate and are generally those in 

which the amount of silica removed by filtration was relatively small.

Wet chemical analyses were performed on many of the precipitates 

that contained substantial amounts of silica, and these showed a mole 

ratio of silica to aluminum near or somewhat less than 1.0. No evidence 

of more siliceous mineral species was obtained from these analyses. 

For this reason and because the final aluminum and silica concentrations of aged solutions fit calculations based on a 1:1 aluminum to

TABLE 3. Experimental solutions in which mole ratio OH:A1 was less than 3.0

[IAP=ion-activity product]

After aging (mg/1)

Solution

IV2._____

113...---.

IV3._____

114.......

IV4......

116.......

IV6......

117.......

IV7.__...

1C.......

ID ....

2C_ ......

2D. ......

119. .......

IV9. ......

3D.......

2C2......

4D. ......

6D___.__.

7D.......

8D_......

10D2.....

10D4_..__

11A......

11B..____

11C......

12A......

12B_...._

15A......

15B......

15C......

15D......

17A......

17B..._._

17B1.....

17B2..._.

Meaii..

Age 

(days)

.. 1,479

761

. . 1, 456

.. 1,439

.. 1,439

.. 1,403

.. 1,403

.. 1,373

.. 1,373

.. 1,362

669

667

667

634

634

528

.. 1,195

490

379

.. 1,045

330

897

897

64

594

594

563

583

119

119

390

118

51

51

.. 86+55

. 41+55

Initial contents 

(nig/1)

Al

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

8.4 

16.9 

16.3 

8.4 

8.4 

7.9 

8.4 

8.3 

9.4 

4.4 

4.1 

12 

12 

12 

12 

12

12 

12 

12 

123 

123 

123 

123

SiO>

21 

100 

100 

100 

100 

100 

100 

100 

100 

18 

96 

18 

98 

94 

94 

100 

26 

98 

98 

92 

98 

50 

49 

34 

37 

40 

17 

17 

42 

44 

46 

47 

2511 

138 

13S 

138

PH

5.24 

4. 69 

4.91 

4.66 

4.85 

4.68 

4.91 

4. 69 

4.96 

4.85 

4.67 

5.12 

4.93 

4.55 

4.80 

4.57 

5.03 

4.60 

4.51 

4.57 

4.83 

4.60 

4.64 

4.48 

4.59 

4.82 

4.90 

5.15 

4.52 

4.64 

4.84 

5.52 

4.46 

5.09 

2 29 

2.35

Passing 0.10 

M m filter

Al

0.20 

1.15 

.34 

1.60 

.35 

1.32 

.10 

1.24 

1.30 

3.20 

2.24 

2.13 

.56 

5.6 

1.34 

3.48 

1.82 

3.0 

2.5 

2.3 

1.7 

1.9 

1.2 

6.4 

5.3 

2.0 

6.5 

4.0 

5.2 

5.5 

1.1 

.04

~. 10

96 

99

SiOj

9.3 

82 

68 

75 

76 

76 

75 

78 

68 

14 

76 

9.1 

70 

70 

63 

88 

14 

86 

88 

77 

84 

3il 

35 

31 

32 

34 

15 

16 

38 

40 

35 

26 

181 

16 

128 

132

Removed Ionic 

by filter strength

Al

8.2 

7.3 

8.1 

6.8 

8.0 

7.1 

8.3 

7.2 

7.1 

5.2 

6.4 

6.3 

7.8 

11.3 

15 

4.9 

6.6 

4.9 

5.9 

6.0 

7.7 

2.5 

2.9 

5.6 

6.7 

10 

5.5 

8.0 

5.8 

6.5 

10.9 

12 

101 

123 

27 

24

SiO:

12 

18 

32 

25 

24 

24 

25 

22 

32 

4 

20 

8.9 

28 

24 

29 

12 

12 

12 

10 

15 

14 

11 

14 

3 

5 

6

1 

4 

4 

9 

21 

78 

122 

10 

6

0.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.005 

.005 

.01 

.10 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.01 

.10 

.10 

. 10 

.10

 log activities in 

aged solutions

Al+3 H4 Si04°

5.94 

4.73 

5.46 

4.70 

5.46 

4.80 

6.07 

4.82 

4.84 

4.41 

4.51 

4.68 

5.16 

4.11 

4.78 

4.30 

4.65 

4.44 

4.39 

4.44 

5.08 

4.62 

4.84 

3.98 

4. 19 

4.67 

4.11 

4.41 

4.10 

4.16 

4.89 

8.09 

3.87 

6.65 

3.21 

3.22

3.81 

2.86 

2. 95 

2. 90 

2.90 

2.90 

2.90 

2. 89 

2. 95 

3.63 

2.85 

3.82 

2.93 

2. 93 

2.98 

2.83 

3.33 

2.84 

2.83 

2. 90 

2.85 

3.19 

3.23 

3.29 

3.27 

3.25 

3.60 

3.57 

3.19 

3.18 

3.23 

3.36 

2.52 

3.58 

2.67 

2. 68

H>

5.20 

4.35 

4.68 

4.47 

4.73 

4. 49 

4. 92 

4.40 

4.59 

4.50 

4.23 

4.74 

4.45 

4.15 

4.45 

4.11 

4.65 

4.26 

4.27 

4.35 

4.53 

4.37 

4.51 

4.30 

4.33 

4.66 

4.49 

4.77 

4.33 

4.41 

4.58 

5.78 

4.02 

5.33 

3.84 

3.85

log

IAPa

11.70 

10. 92 

11.26 

11.62 

11.66 

11.54 

11.58 

10. 98 

11.96 

10. 92 

10.66 

11.44 

10.52 

10.82 

11.18 

10.40 

10.14 

11.00 

11.18 

11.42 

11.32 

10.60 

10. 92 

11. 26 

11.08 

12. 12 

11.52 

12. 66 

11.40 

11.78 

11.24 

11.78 

11.34 

11.52 

11. 28 

11.34

.11.28
E12 CHEMISTRY OF ALUMINUM IN NATURAL W^TER

TABLE 4. Experimental solutions in which mole ratio OH:A1 was greater than 3.0

[IAP =ion-activity product]

After aging (mg/1) log activities in aged 

Initial contents solutions

Solution Age (mg/1) Passing 0.10 Removed by Ionic log 

(days) ton filter filter strength IAP b

_____________ ______________ A 1 JJ. JJ+

Al SiO2 pH Al SiCh Al SiO2 (OH).~ SiCM

11D-. _ _ 64

11F_______ 64

11H. 659

12C.______ 199

12D 199 

12E--.!.!_ 199

12E2. 659

15E.______ 118

15F_.____ 118

15G__ _ 390

15H ______ 390

16A.__.___ 361

17C_______ 106

17D__.__ 95

19

12

12

12

19

12

12 

12

12

12

12

12

12

12

12

12

12

12

123

123

43

45

56

CO

18

19 

30

32

28

30

49

50

51

53

54

11

48

130

218

7.60

8.67

9.08

10.20

9.79

6.50

8.65 

9.85

8.90

9.15

9.22

6.65

7.67

8.14

8.49

8.90

9.75

9.25

8.68

7.02

0.01

1.4

5.2

11.6

8.0

.03

.06 

2.3

.18

.40

.46

.05

.36

.83

.62

.90

.50

.24

1.4

.03

25

23

38

41

6.6

6.9 

16

11

11

9.4

25

23

23

22

23

4.7

15

2.2

24

12

10.6

6.8

.4

4.0

12

12

9.7

12

12

12

12

12

11

11

11

12

121

123

18

22

9

15

13

11.4

12 

14

21

17

21

24

27

28

31

31

6.3

33

128

194

0.01

.01

.01

.01

.01

.01

.01 

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.01

.10

.10

6.48

4.33

3.76

3.41

4.04

6.03

5.70 

4.11

5.22

4.87

4.81

5.80

4.92

4.56

4.68

4.52

3.73

5.10

4.40

6.11

3.38

3.42

3.31

3.42

3.65

3.96

3.94 

3.58

3.74

3.75

3.82

3.37

3.42

3.49

3.44

3.42

4.18

3.60

4.43

3.40

7.12

7.72

9.13

9.54

8.90

6.55

7.24 

8.16

7.92

7.87

8.19

7.27

8.46

8.94

8.08

8.07

8.95

7.54

8.17

6.65

-33. 96

-30.94

-32.40

-32.74

-33. 18

-33.08

-33.76 

-31.70

-33. 76

-32. 98

-33.64

-32. 88

-33.60

-33.98

-32.40

-32.02

 33.72

-32. 48

-34.00

-32. 32

-32.98

silicon ratio better than other ratios, significant amounts of 2:1 l&yer 

silicate structures probably were not formed.

Additional data for solutions given in tables 3 and 4 include values 

for the negative log of the molar activity of uncomplexed A1+3 in acid 

solutions (or of A1(OH) 4~ in basic solutions), negative log activity of 

H4SiO4 (aq) and pH, and the computed ion-activity product for each 

aged solution. Appropriate equations from table 2 and the DebyeHiickel equation were used to compute activities of sclutes and ion 

activity produced.

Nearly all the solids present at the end of the aging period in the 

solutions given in tables 3 and 4 were amorphous to X-rays. Precipitates from some solutions in which silica concentrations were very low 

did not contain significant proportions of silica. These results are not 

included here, but these precipitates did give X-ray patterns for 

gibbsite or bayerite, depending on pH. Additional data on composition 

of precipitates are given later in this paper.

TESTS FOR EQUILIBRIUM

Although the solids produced from solutions given in tables 3 and 4 

did not show any definite evidence of containing aluminum hydroxide 

in crystalline form, it is of interest to see if the aluminum concentrations observed might obey predictions based on solubilities of aluminum hydroxide species alone. Figure 1 is a plot of log [A1+3 ] in 

moles per liter against pH for the samples in table 3, all of which had 

a final pH below 6.0. Solutions IV2-10D4 were prepared \)y adding a
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA 

3.0

E13

9.0

FIGURE 1. Log of activity of uncomplexed aluminum versus pH in aged 

solutions containing silica. Dashed line is equilibrium line for microcrystalline gibbsite. O = nrst 23 solutions, table 3; X =last 13 solutions, 

table 3.

separate solution of silicic acid (method 1), and the remaining solutions 

were prepared by adding silicate that was included in the basic stock 

solution.

It is evident that the microcrystalline gibbsite solubility relationship does not satisfactorily explain the results observed in these solutions because the points scatter and tend to show too low a solubility 

between pH 4.0 and 5.0.

If the synthesis reaction is postulated to be the reverse of 

Al2Si2O5 (OH) 4 +6H+-2Al+3 +2H4SiO4(aq) + H2O,

the solubility relationship at equilibrium would be

[Al+3 ] 2 [H4Si04 ]2

[H+]8 ~ AS0 .
E14 CHEMISTRY OF ALUMINUM IN NATURAL WATER

Figure 2 shows log activities of Al+3 and H4SiO4 added together and 

plotted against pH. This gives a relatively well defined relationship. 

Aline with slope 3 drawn through the points is a good fit and suggests the reaction postulated shows correctly the proportions of 

reactants and products. The equilibrium constant estimated from the 

graph is 1011 3 . A more exact value calculated from the mean of the 

36 sets of data is 1011 - 28 and is given in table 3.

By using 1011 28 as the equilibrium constant and the standard freeenergy data in table 1, the standard free energy of the aluminosilicate

4.0

5.0

6.0

9 7.0

CO

8.0

o 9.0

10.0

11.0 -

\

\

\

\

\

\

o°\

00 X

0 \

oo 

o

\

\

V

\

\

12.0 

3.0

\ \ \

\x _|

\

___:

4.0 5.0 6.0

PH

FIGURE 2. Sum of logs of activities of uncomplexed aluminum and aqueous 

silica versus pH in aged solutions. Dashed line is equilibrium line for synthetic halloysite. Q = first 23 solutions, table 3; X = last 13 solutions, 

table 3.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA El5

solid is calculated to be 896.7 kcal per mole. This calculation assumes 

the formula for this material is equivalent to kaolinite or halloysite 

without considering added water of hydration.

The spread in values of ion-activity products in table 3 suggests 

that an uncertainty of about ± 1 kcal is a reasonable estimate. The 

material produced in this set of solutions is somewhat less stable than 

halloysite that has a readily recognized crystalline character, for 

which Robie and Waldbaum (1968) gave a value of 898.4 kcal per 

mole.

Several of the sets of data in table 3 are of special significance in 

evaluating the validity of the equilibrium assumption. A special 

experiment was made by using solution 17B and a duplicate of that 

solution. The analytical data for 17B after 51 days of aging are shewn 

in table 3. The duplicate sample had essentially the same composition 

and behavior. The solutions initially contained 123 mg/1 Al and 138 

mg/1 SiO2 (aq), with an ionic strength near 0.1. During aging the pU of 

17B increased slowly from an initial value near 5 to a value of 5.33 

after 51 days. At that time nearly 90 percent of the silica and almost all 

the aluminum could be removed from this solution by filtration through 

a O.lOjttm porosity filter. Although solution 17B was opalescent when 

originally prepared, it appeared to be practical^ clear after a few 

weeks of aging. The point in figure 2 representing the composition of 

the solution after 51 days falls essentially on the equilibrium line. The 

duplicate solution prepared December 3 had a very similar composition to 17B after a similar period of aging and thus is not included in 

table 3.

After 3 months of aging, aliquots of 17B and the duplicate solution 

were put into separate polyethylene bottles, and 1.0 ml of 0.863 

molar perchloric acid was added for each 100 ml of solution. This 

amount of H+ was sufficient to react with about two-thirds of the calculated bound OH in the aluminosilicate. The bottles were allowed 

to stand at 25°C, and the pH of the solutions was measured from 

time to time until it had stabilized. One hour after the acid was added 

the pH values were near 2.3. A gradual increase in pH occurred for 

several weeks, but after about a month no further significant change 

took place. After 55 days of reaction the solutions were filtered 

through 0.10 jttm porosity filters, and the concentrations of aluminum 

and silica in the filtrate were determined. Calculating the activities 

of A1+3 and H4SiO4 in these solutions gave the closely agreeing values 

reported for solutions 17B1 and 17B2 in table 3. These values plot 

almost exactly on the equilibrium line in figure 2. This demonstration 

of the reversibility of the precipitation reaction to reattain equilib-
E16 CHEMISTRY OF ALUMINUM IN NATURAL WATER

rium at a much lower pH greatly strengthens the interpretation that 

the material formed in the test solutions has a well-defined equilibrium solubility, even though its crystal structure is not well defined.

EFFECT OF DISSOLVED-SILICA ACTIVITY

The initial amounts of silica in the aging study solutions so far 

considered were generally rather substantial. Replicates of most of 

the samples in the first part of table 3 (IV2-10D4), however, were 

made by using lower silica concentrations. Where the initial silica 

content was 2.0 mg/1 or less as SiO2 and the aging was a few weeks or 

months, the precipitates all contained gibbsite identif able by X-ray 

diffraction. Solutions that contained 20 mg/1 of SiO2 did not produce 

gibbsite and generally gave equilibrium results after long aging that 

agreed well with the data shown in figure 2. Solutions containing 

silica concentrations between 2 and 10 did not all show crystalline 

gibbsite. The highest silica concentration observed in solutions that 

did produce gibbsite was about 9 mg/1. Solutions in this range may 

have tended to form mixed solids consisting of aluminum hydroxide 

and aluminosilicate structures. The relationship of the boundary 

between the aluminum hydroxide and aluminosilicate stability fields 

to dissolved-silica concentration indicates there should be a silica 

concentration at which both solids are stable. These experiments 

suggest this boundary probably lies near a silica concentration of 9 

mg/1 SiO2 . Solubility calculations for the solutions lov^ in silica suggest that they are generally close to equilibrium with respect to 

microcrystalline gibbsite and near the equilibrium line in figure 1.

REACTIONS IN ALKALINE SOLUTIONS

A considerable number of solutions were prepared to explore the 

behavior of aluminum and silica in systems where the pH remained 

above neutrality during aging. These solutions were prepared by using 

the techniques already described for the pH 4.5-6.0 rarge. Solution II 

contained silicate ions, and the silicate concentration was varied in 

the different groups to give a range of silica to aluminum mole ratios 

from as much as 3 in some solutions to as little as 0.1 in others.

The alkaline solutions all appeared to decrease substantially in 

concentration of dissolved aluminum and silica during aging. Initial 

and final concentrations are given in table 4. The activities of silica 

and A1(OH) 4 ~ for these at the end of the aging period were determined 

by analysis and corrected for ionic strength. Aging times ranged from 

about 2 months to more than 2 years. As with solutiors at lower pH, 

the material reported in solution was determined on filtrates passing 

0.1 nm porosity filter membranes. Calculated activities of A1(OH) 4 ~
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA El7

were obtained by solving appropriate solute species equilibria from 

table 2. This form of aluminum is strongly predominant in solutions 

with pH>7.0. The silicic acid activities were calculated by correcting 

total dissolved silica for the effects of dissociation.

Solutions in the alkaline pH ranges that were low in dissolved silica 

(SiO2 <^10 mg/1) yielded bayerite on aging. Crystallinity was slow to 

develop but was detectable by X-ray diffraction after two months in 

a few solutions and in a total of six solutions after a year of aging. 

Data for these solutions are not included in table 4.

Although there may be some question as to the applicability of the 

aluminosilicate solubility data for low pH solutions in table 3 to these 

more alkaline solutions, the same general approach was tested to see 

what results would be obtained. As a first approximation, figure 3 

was prepared to show the relation of pH and dissolved aluminum 

where bayerite was identified in the precipitate and where no crystalline aluminum hydroxides were identified. The dashed line is the 

solubility at equilibrium of bayerite synthesized in the absence of 

silica (Hem and Roberson, 1967). Five of the six experimental points 

for the solutions where bayerite was present plot very near the dashed 

line. Deviations from the equilibrium value could be related to 

impurities in the precipitate or to slowness of reactions as equilibrium 

is approached. The aluminum solubility in the other solutions is 

obviously not controlled by bayerite precipitation.

Solutions which showed substantial losses in silica concentration 

during aging were then tested for equilibrium with 1:1 aluminosilicate 

by plotting the sum of log [A1(OH) 4 ~] + log [H4Si04] against pH 

in figure 4. The dashed line here represents equilibrium with a solid 

having the same composition and standard free energy of formation 

as the halloysite synthesized in the solutions below pH 6.

Although the experimentally determined points in figure 4 are 

somewhat more scattered than those in figure 2, they do follow the 

trend of the dashed regression line, which represents the solubility of 

halloysite determined from data for samples at lower pH. In fact, 

this is about as close a fit to the points as can be drawn visually. The 

possibility of a 2:1 silicon to aluminum mole ratio in the precipitated 

product was explored by plotting the relationship log [A1(OH) 4 ~] 

+ 2 log [H4Si04 ] versus pH for solutions in table 4. That the resulting 

points (not shown) were substantially more widely scattered than the 

ones in figure 4 indicates a 1:1 composition.

An average value for the ion-activity product for the relationship

[Al(OH)r] 2 [H4Si04] 2 [H] 2 =IAP,
E18 CHEMISTRY OF ALUMINUM IN NATURAL WATER

1 1 1 1 1 1 1 1 1

\

\

\\

\

\

\

0 X

\ o o x

\

\

\

x° \

\

00° \

\x

X \ x

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\

\

\

\

\

V

\

\

\

\

\ -

\

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\

1 1 1 1 1 1 1 1 1

o -g'3"S

o

0

cc

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co

1^_^

0 Q,

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"m «

bD M

S S

o3 o

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cc c3 

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00 D. ^ o

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w ^

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c3

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h?

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INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E19

°. X 

co a

wa

£ £

C3 CO

co o 

1 g

w
E20 CHEMISTRY OF ALUMINUM IN NATURAL WATER

was calculated from data in table 4 as 10~32- 98 . This value gives for 

the free energy of formation of the aluminosilicate precipitate 897. 0 

kcal per mole. The uncertainty in the number is probably at least 1 

kcal. This value is nearly the same as the 896.7 kcal obtained for 

material produced between pH 4.5 and 6.0 and suggests the same kind 

of material is formed in all sets of solutions. Because of the experimental uncertainty the value should be rounded to the nearest 

kilocalorie, or 897 ±1 kcal per mole.

STABILITY FIELDS AND ALUMINUM SOLUBILITY GRAPHS

The determined value for halloysite free energy can be used with 

the value for microcrystalline gibbsite to prepare stability-field 

diagrams. For the boundary between solids, the reaction

2Al(OH) 8 (c) +2H4SiO4 (aq) - Al2Si2O5 (OH) 4 (c) +5H2O

reduces to

[H4SiO4] 2 =10-7 - 66 

or

This activity is equivalent to a concentration of 8.9 mg/1 dissolved 

SiO2 . Because the dissolved silicic acid carries no ionic charge, the 

activity coefficient will be near 1.0. When the pH is high enough, 

however a significant fraction of the silicic acid will be dissociated 

into silicate anions, and the total silica concentration will not be 

equivalent to the calculated boundary based on [H4SiO4]. Data obtained by Hem and Roberson (1967, p. 48) showed that above neutral 

pH the freshly precipitated form of aluminum hydroxide had the same 

solubility as microcrystalline gibbsite.

Figure 5 is a stability-field diagram for microcrystalline gibbsite 

and synthetic halloysite with aluminum solubility contours in both 

fields. This diagram specifies an ionic strength of 0.01. The differences 

between this diagram and one for ionic strength 0.00 are very small, 

and this figure should be applicable to most natural waters whose 

dissolved-solids concentration is not greater than 1,000 mg/1. Figure 

6 shows the same system at an ionic strength of 0.10. The effect of 

increased ionic strength is to increase aluminum solubility in the pH 

range where multivalent aluminum species predominate. The solubility of aluminum is calculated from equations 1-5 and 17 in table 2 

for gibbsite and from equations 2-8 and 17 in table 2 for halloysite.

As noted earlier the laboratory results showed qualitatively that 

crystalline gibbsite did not appear in solutions whose dissolved-silica
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E21

IN MILLIGRAMS PER LITER

p

b

o 

o 

o

1000

§ o' o'o 

500 -

9.0 100

FIGURE 5. Stability fields of synthetic halloysite and microcrystalline 

gibbsite and solubility of aluminum as functions of pH and activity of 

undissociated aqueous silica at 25°C and 1 atmosphere. Ionic strength 

0.01; fluoride absent.

476-920 O - 73 - 4
3

O G CR O

^ g, 5s a 0 2r o- »

- co co W

g.p S g-

[H 4Si04 ] as Si02 , IN MILLIGRA

II II

<* o cs

' ? ,? i-h co *<d ^

O>

2. o a H "-H M

2 SB

tf. ^

g^

r

&t-l

a P 2.

CO O

co y. ^ 

g- <! M
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E23

concentration remained above about 9 mg/1 as SiO2 . The solubility 

data provide a more quantitative verification of this.

The position of the aluminum hydroxide-aluminosilicate boundary 

is shifted if other values are selected for free energies of the solids. 

The ones used for figures 5 and 6 are appropriate for systems in which 

freshly precipitated material is present. If the clay miner?,! is 

considered to be a well-organized kaolinite, the boundary between 

hydroxide and silicate fields is shifted to a considerably lower eilica 

concentration. The value of 902.9 kcal per mole for kaolinite, for 

example, would give an equilibrium silica concentration at the twosolid boundary between kaolinite and well-crystallized gibbsite of 

about 1.0 mg/1 SiO2 . If the aluminum hydroxide species were microcrystalline gibbsite in equilibrium with this kaolinite, the value for 

silica would be less than 0.1 mg/1 SiO2 .

Reversible equilibria in systems involving crystalline kaolinite 

probably are unrealistic for predicting conditions in the experimental 

solutions used in this study because our material did not display good 

crystallinity.

The silica concentrations shown in figures 5 and 6 extend to values 

higher than the usually accepted solubility of amorphous forms of 

silica at 25°C (100-200 mg/1 as SiO2). The highest silica activities 

shown are unlikely to be encountered in natural water with ionic 

strength less than 0.1.

Once the spacing of solubility lines on the aluminum solubility 

grid has been calculated for a given ionic strength, the grid can be 

used to estimate equilibrium conditions involving any pair of aluminum 

hydroxide and 1:1 clay mineral species to which the kaolinite formula 

applies. It is only necessary to calculate the correct equilibrium 

position of the horizontal boundary between the solids on the silica 

concentration axis and aline the vertical edge of the grid with the 

proper pH. The grid then is moved up or down so the horizontal phase 

boundary lies on the calculated equilibrium value for silica.

The solubility lines in figures 5 and 6 take into account the hydroxy 

complexes of aluminum but ignore other solute complex species. 

For a satisfactory fit to natural water, it may be necessary to consider 

the effect of other aluminum complexes.

EFFECTS OF FLUORIDE

As noted elsewhere (Hem, 1968; Roberson and Hem, 1969) the 

complexes formed in solution between aluminum and fluoride are 

particularly strong and can exert an important influence on aluminum 

solubility. Calculations adapted from Roberson and Hem (1969) were 

grafted to the aluminum-silica system calculations to give the additional solubility diagrams, figures 7 and 8, which show the influence of
E24 CHEMISTRY OF ALUMINUM IN NATURAL WA^ER

CA | IN MILLIGRAMS PER LITER

1000

4.0 5.0 6.0 7.0 8.0 9.0 10.0

PH

FIGURE 7. Stability fields of synthetic halloysite and microcrystalline 

gibbsite and solubility of aluminum as functions of pH and activity of 

undissociated aqueous silica at 25°C and 1 atmosphere. Ionic strength 

0.01; total fluoride concentration 0.19 mg/1 as F.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E25

IN MILLIGRAMS PER LITER

4.0 5.0 6.0 7.0 8.0 9.0 10.0

FIGURE 8. Stability fields of synthetic halloysite and microcrystall: ne 

gibbsite and solubility of aluminum as functions of pH and activity of 

undissociated aqueous silica at 25°C and 1 atmosphere. Ionic strength 

0.01; total fluoride concentration 1.9 mg/1 as F.
E26 CHEMISTRY OF ALUMINUM IN NATURAL WATER

fluoride at the level of 0.19 and 1.9 mg/1 total fluoride, respectively. 

These concentrations are equivalent to 10~5 and 10~4 moles per liter 

and include not only the free fluoride ion F~ but also an^ fluoride that 

is present in the form of aluminum fluoride complexes.

Figures 7 and 8 are for an ionic strength of 0.01. Some shifts will 

occur in the solubility lines if other ionic strengths are assumed. 

The general pattern is for increasing ionic strength to increase aluminum solubility if other factors remain constant. The importance 

of this effect is indicated by differences among solubility lines in figures 5 and 6.

The aluminum solubilities calculated for figures 7 and 8 include all 

fluoride and hydroxide complexes and require simultaneous solution 

of 10 or more of the equations in table 2. In the interest of brevity 

these calculations are not given here; details can be obtained in 

Roberson and Hem (1969).

In the presence of sodium at levels of concentration attained in 

some of the experimental solutions, it is possible for aluminum and 

fluoride to combine to form cryolite, Na3AlF6 . Conditions favoring 

precipitation of cryolite and the related solid ralstonite vrere described 

by Roberson and Hem (1969). Ralstonite obtained in that study 

appeared to be a sodium aluminum hydroxy-fluoride, but its exact 

composition could not be ascertained.

Cryolite solubility graphs published by Roberson and Hem (1969, 

p. 28) show that where sodium concentrations are near 0.01 or 0.10 

molar cryolite will not precipitate unless fluoride concentrations are 

substantially above 19 mg/1, regardless of the aluminum concentration. The importance of cryolite precipitation as a solubility control 

over aluminum in natural systems presumably is not great. If the 

concentration of fluoride were 190 mg/1 (about 10~2 molar) and the 

sodium concentration were 2,300 mg/1 (about 10" 1 molar), however, a 

stability-field graph for 25° and 1 atmosphere would show a large 

cryolite stability region. The cryolite-gibbsite and cryolite-halloysite 

boundaries are shown in figure 9 for a fixed concentration of 10" 1 moles 

per liter of sodium and a fixed concentration of 10~2 moles per liter of 

fluoride. An ionic strength of 0.10 is specified in the calculations which 

are based on the equilibria

2Na3AlF6 (c) +2H4SiO4 (aq) +H2O(1) = Al2Si2O5 (OH) 4 (c) +6Na+

+ 12F-+6H+

and

Al(OH) 3 (c)+3Na+ +6F-=Na3AlF6 (c)+3OH'-.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E27

 Al

1000

500

300

S 200 IIIM MILLIGRAMS PER LITER

100

50

20 k

- 30

C4 

O

V) 

LL 

O

2

g

i10

3h

Cryolite

3.0 4.0 5.0 6.0 7.0 8.0 9.0 10.0

pH

FIGURE 9. Stability fields of cryolite, synthetic halloysite, and microcrystalline gibbsite and solubility of aluminum as functions of pH and 

concentration of dissolved silica at 25°C and 1 atmosphere. Ionic strength 

0.10; total fluoride concentration 190 mg/1 as F; sodium concentration 

2,300 mg/1 as Na.
E28 CHEMISTRY OF ALUMINUM IN NATURAL WATER

Mass law expressions for these can be written as

I J.1 u> I L-J- I l-L-L I - rt _ no 07 

± : r__^ J___iii o». 6t

[H4Si04 ] U 

and

ITkTT-13

.^ini.19 ft J-" [Na+]3 [FEquilibrium constants were calculated by using free energy values in 

table 1. Because aluminum is not given as a solute species in either 

equation, the solubility of aluminum cannot be shown by these 

equations.

To take the ionic strength effect into account, so that stoichiometric concentrations rather than activities can be used for Na+ and 

F~, appropriate values for TNa+ and JF - were obtained from Butler 

(1964, p. 435) and were used to calculate equilibrium constants for 

7=0.10:

(Na+) 3 (F-) 6 [H+] 3 n_ 3 , 32

(H4Si04 ) 

and

 =io-

__ 14 \R iu (Na+) 3 (F-)

Parentheses indicate concentrations, and square brackets activities. 

It is assumed the activity coefficient for dissolved silicic acid is 1.0. 

At high pH, equation 8 (table 2) for the first dissociation constant of 

silicic acid must be used to determine the effect of dissociation. The 

activity coefficient of the H3SiO4~ ion at 0.10 ionic strength is estimated 

to be 0.77, assuming it behaves like comparable monovalent ions 

listed by Butler (1964).

The stability boundaries for the three solids are shown in figure 9 

by solid lines. The curvature in the boundary between gibbsite and 

synthetic halloysite is the result of the effect of dissociation of dissolved 

silica at high pH. The ordinate in figure 9 is total dissolved silica as 

SiO2 , the usual analytical value. In figures 5-8 the ordinate was the 

actual activity of the undissociated material.

The dashed lines in figure 9 show solubility of aluminum, which at 

concentrations above 0.01 mg/1 is essentially all present as A1(OH) 4~. 

Below pH 8.3 the solute species are aluminum fluoride complexesTheir total concentration is not a function of pH, nor ir it a function 

of dissolved silica, within the remaining area of the diagram. Solubility 

of aluminum in all this area is approximately 0.01 mg/1. The specification needs to be made, however, that aluminum is present throughout in small quantity compared with the amount of fluoride. Thus
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E29

the total fluoride dissolved is near the value for the uncomp^xed 

fhioride ion.

At very low pH (near 1.0) nearly all the fluoride will be tied up as 

HF°, and thereby the solubility of aluminum can be greatly increased. 

The pH grid in figure 10 does not extend into this region.

The minimum fluoride concentration that will give a cryolite stability region in the presence of 10~J moles per liter of sodium is near 

10-2.8 moiar (about 30 mg/1 as F~), and the aluminum solubility in 

the cryolite region will be near 1.0 mg/1. It appears evident that 

precipitation of cryolite from a natural solution at 25°C requires 

unusually high concentrations of fluoride, and the process would 

definitely not be expected to be a common one where waters of low 

dissolved-solids content are prevalent.

The position of the phase boundary between cryolite and the 

aluminum hydroxide solid depends on both the sodium and fluoride 

concentrations in the system. A decrease of one log unit in sodium 

concentration shifts the boundary one pH unit to the left, and a 

decrease of one log unit in fluoride shifts the boundary two pH units 

to the left. More highly concentrated solutions of sodium than tenth 

molar (about 2,500 mg/1) are not uncommon, however, and in these 

cryolite precipitation might occur when fluoride content was in a more 

commonly encountered range. As shown by the data of Roberson and 

Hem (1969), the mutual influence of complex formation between 

aluminum and fluoride results in an optimum value of the mole ratio 

of dissolved fluoride to dissolved aluminum for cryolite precipitation. 

The precipitation ought to occur most readily when the ratio CF :CA i 

is near 10, where C is the total concentration of all species of each ion.

Influence of silica on the cryolite system would appear to be relatively minor at 25°C, at least when viewed through the equilibrium 

calculations made for figure 9. If the reactions between silica and 

aluminum were faster than those involved in cryolite precipitation, 

there might be some tendency toward formation of metastable silicates 

within the cryolite stability region. Actually, however, laboratory 

studies show that at 25°C cryolite synthesis is rapid, and aluminosilicate syntheses are generally slower.

Some experimental data on the behavior of aluminum in the presence of silica, fluoride, and hydroxide were obtained by using established techniques (Roberson and Hem, 1969, p. 21). Six solutions 

(A-F) were prepared, each having fluoride and sodium concentrations 

near 10" 1 moles per liter and initial aluminum near 1,000 mg/~ (between 10~3 and 10~2 molar). In four of these (A-D) silica was also 

present, ranging in concentration from about 45 to near 300 mg/1 SiO2 , 

and in the remaining two there was no silica. The pH observed when 

the solutions were prepared ranged from 4.58 in solution A to 8.85 in
E30 CHEMISTRY OF ALUMINUM IN NATURAL WATER

solution D. A finely divided precipitate appeared almost immediately 

in all six solutions. Most of the silica disappeared from solutions C 

and D. Essentially no loss of silica occurred in solutions A and B. 

Solubility calculations showed all these solutions to be strongly supersaturated with respect to cryolite at the time they were prepared. 

Composition data are given in table 5 for solutions A-F after 16 

months of aging at 25°C. Aluminum and silica were determined by 

atomic absorption spectrophotometry and represent total concentrations of all aluminum and silicon species in solution. Activities of 

uncomplexed Na+, F~, and H+ were measured with specific ion electrodes. The solids which had precipitated were examined by X-ray 

diffraction. The solids were removed from the solution before analysis 

by filtering the aliquots through membranes with 0.10/mi maximum 

diameter pores.

As indicated in table 5 cryolite was formed in all ^x solutions. 

Two solutions having the lowest pH precipitated ralstonite, and the 

three most alkaline solutions had a form of aluminum hydroxide, 

identified by X-ray diffraction as nordstrandite. No silicate species 

was identified by X-ray, although the initial silica concentrations near 

10~2 - 5 molar in solutions C and D had nearly all disappeared from 

solution when the first analyses were made six days after the solutions 

were prepared. The quantities of cryolite precipitated in these two 

solutions was nearly 10 times as great as the maximum possible 

quantity of silicate precipitate, and it is possible that the X-ray 

diffraction technique could not have detected so small a proportion, 

even if the aluminosilicate had been well crystallized.

The silica present after 16 months in solutions E and F apparently 

came from impurities introduced into the reagents or by other means, 

for none had been added originally.

TABLK 5. Composition of solutions A-F after 16 months aging 

[C is total concentration of all species of ion)

Moles per liter

A. ___________

B.__ _

C. .....___.__

D......

E

F._

tin

______ 4. 41

______ 4. 32

______ 6.72

______ 9. 40

______ 4.19

______ 8.29

PIT J

4. 06

4. 02

3. 19

1. 82

4. 01

2. 27

PU> aj

1. 06

1. 04

1. 08

1. 01

1. 04

1. 06

CAI

3. 34X10-2

1. 57X10-2

1. 26X10-"

. 74X10-5

2. 07 X 1C-2

1. 1X10-5

Csio

2

7. 17X10-"

1. 57X10-3

4. 50X10- 5

1. 12X10-"

1. 3X10- 5

. 5X10- 5

NOTE. Solid species identified by X-ray diffraction: 

Cryolite A-P; 

Ralstonite B and E; 

Aluminum hydroxide C, D, and P.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E31

In table 6 are results of three different calculations using data from 

table 5 for solutions A-F. The first calculation tests for adherence to a 

two-solid equilibrium involving cryolite and a clay mineral with the 

formula and stability of the synthetic aluminosilicate material 

precipitated in earlier experiments. The equilibrium relationship is 

that used for establishing the cryolite-halloysite (c+h) boundary in figure 

9. Although the calculated ion-activity products are about the same 

for all four solutions that initially contained silica, they are consistently 

less negative than the theoretical value.

The second calculation tests for a two-solid equilibrium involving 

cryolite and microcrystalline gibbsite (c+g). The three lowest activity 

products, which come the closest to the theoretical value, are for the 

three solutions from which precipitates did contain aluminum 

hydroxide as identified by X-ray diffraction.

The third calculation tests for agreement with the cryolite solubility 

(cs) product reported by Roberson and Hem (1969). The average of 

the six values is 10~33 - 78 , which is in good agreement with their 

theoretical value of 10~33 - 84 . Cryolite was identified in all these 

precipitates. Considering the large exponents on some of the measured 

activities, the spread of the values is probably within experimental 

error.

One can also test these results for possible adherence to hydroxide 

or aluminosilicate equilibria by plugging appropriate values into 

figures 1-4. Solutions C, D, and F plot reasonably close to the equilibrium line in figure 3, but none of the other pertinent sets of data come 

close to any equilibrium lines in these four figures. It would be possible to interpret the results for solutions C and D by postulating the 

precipitation of a somewhat more stable form of halloysite than the 

material synthesized in solutions that did not contain fluoride. The

TABLE 6. Ion-activity prodiicts for solutions A-F, after 16 months aging 

[Text explains abbreviations in boxheads]

Solution

A. ..____.___.__._

B. ...__.._.._.___

C..._

D ____ _

E. ..__._____.___.

F

IAPc+h

10-37.63

ID"37 M

1Q-38.19

._.____ ID-3?- 9 *

IAPc+e

101 -23

1Q1.80

lQ-0.54

1Q-0.15

1Q2 .25

1Q0.33

IAPc,

1Q-34.16

1Q-34.38

1Q-33.92

10-34.75

10-34.11

10-33.17

[Al]

1Q-6.61

10-7.03

1 0~n -54

10~20.81

10-6 .93

10-16 .37

Theoretical ___________ 10-39.37 1Q-1 - 19 10-33.84

Equations:

IAPc+h=[Na+p [F-]« [H+p [H4Si04]-i 

lAPc+B=[Na+P [F-]« [OH-]-3 

IAPcs=[Na+]
E32 CHEMISTRY OF ALUMINUM IN NATURAL WAT^R

standard free energy of formation of such a species would need to be 

near 900 kcal per mole to correct for the deviations of the activity 

products for these solutions from the theoretical value. The evidence 

for existence of aluminosilicate species in solutions A-F is inconclusive, and no quantitative interpretation seems justifiable.

Somewhat more conclusive evidence is available for the aluminum 

hydroxide, which was detected by X-ray diffraction in the precipitate 

from solutions C, D, and F. The strongest hydroxide peak observed is 

at the spacing of nordstrandite (Robert Schoen, oral commun., 1970). 

This form of aluminum hydroxide was produced in an earlier experiment through alteration of a previously formed precipitate of bayerite 

during aging in solution at pH near 9 (Schoen and Roberson, 1970). 

The earlier experiments had all produced mixtures of nordstrandite 

and bayerite, however, and did not seem to give a basis for calculating 

the standard free energy of formation of nordstrandite. As pointed 

out by Schoen and Roberson (1970), the crystal form of the aluminum 

hydroxide species seems to be related to the pH of the solution from 

which they are precipitated. Thus, gibbsite does not seem to be 

reversibty altered to bayerite without substantial changes in the 

solution. Baj^erite to nordstrandite alteration does appear to occur, 

however, in mildly alkaline systems, and the two species would be 

expected to have similar free energies of formation. The value given 

by Hem and Roberson (1967) for bayerite is 274.0 kcal per mole.

Solubility data from solutions C, D, and F offer a basis for calculating a free energjr of formation for nordstrandite. The average of the 

three ion-activitjr products for the aluminum and hydroxide in these 

solutions gives a value of 273.6 kcal per mole. This must be considered as not significantly different from the value for bayerite, 

because the uncertainty in this value is near ±1.0 kcal.

The experimental data for solutions containing both fluoride and 

silica show that the dissolved silica has no important inhibiting 

effect on the formation of crystalline cryolite in systems that are 

supersaturated with respect to cryolite. Reservations were expressed 

by Roberson and Hem (1969, p. 36) that the relative rarity of cryolite 

as a rock mineral might be due to inhibition of cryolite precipitation 

by silica. Such an effect was not observed in these experiments, and it 

probably is not important in most natural systems, Tie solutions 

that precipitate cryolite, however, must be higher in sodium and 

fluoride than can ordinarily be expected in ordinary circulating waters 

near the land surface. It seems likely the rarity of such solutions, 

especially with respect to fluoride concentration, is the principal 

reason that cryolite is a rather rare mineral.

Solutions A-F have about the sodium activity and ionic strength 

assumed for figure 9 but differ in other respects from the system
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E33

conditions used in preparing the diagram. Solutions D and F, however, 

are fairly close to the assumed fluoride activity. The shifts of the 

cryolite-gibbsite boundary required bring the aluminum solubility 

grid into close agreement with the CM values reported in table 5 f c r the 

two solutions.

Eugster and Jones (1968) described sodium-aluminum silicate gels 

that formed at Lake Magadi, Kenya. These apparently formed by 

water from hot springs that was alkaline and rich in sodium, fluoride, 

and silica coming in contact with extrusive igneous rocks. The gel 

could be crystallized in the laboratory to form analcite (a zeolite). 

Cryolite evidently was not produced, but the composition of the spring 

water indicated by analyses in the paper of Eugster and Jones would 

not be within the cryolite stability region in figure 9.

APPLICATION TO NATURAL-WATER COMPOSITION

The practical value of experiments and calculations of the type 

described here lies mainly in their applicability to natural systems. 

Hopefully the solubilities of the synthetic aluminosilicate might approximate some of the natural products of rock-weathering reactions 

at ambient earth-surface temperatures.

There are many references in the literature to the occurrence of 

X-ray amorphous clay-type minerals in soils and stream sednrents, 

and they probably are also present in most ground water systems .

Most of the chemical analyses of natural water that can be obtained 

in published literature are not satisfactory for testing aluminosilicate solubility calculations. Generally aluminum concentrations 

are not determined in the water analyses, and even when they are, 

the results may be unreliable. Metastable colloidal aluminum hydroxide or aluminosilicate with some of the properties observed in our 

synthetic halloysite can be expected to occur in river water. Removal 

of such material by filtration may be incomplete unless a very low 

porosity filter is used.

Samples of high flow in the Mattole River near Petrolia, Calif. , 

collected November 8, 1969, as part of geochemical investigations by 

Kennedy (1971) were filtered at the time of collection through plastic 

membrane filters with 0.45jum maximum diameter pores, and aluminum and iron concentrations were determined in the filtrates. The 

filtered samples were passed through another plastic membrane 

filter, one having O.lOjum maximum diameter pores, and the aluminum and iron concentrations were determined on these filtrates also.

Sample Filter, O.tfum (nig/l) Filter, O.lO^m (mg/l)

Al Fe Al Fe

2092... ________ 0.11 0.10 0.08 0.03

2093___________ .14 .08 .09 .04
E34 CHEMISTRY OF ALUMINUM IN NATURAL WATER

The apparent aluminum concentration was decreased by about a 

third by the second filtration. Iron concentration was decreased by 

one-half or more. This behavior of iron has been observed before 

(Hem and Cropper, 1959, p. 25), but the effect of filtration on apparent 

aluminum content of river water has not been documented as effectively.

Investigators who have studied the chemistry of river water have 

not generally been particularly concerned with colloidal material 

contained in water samples. Samples are usually filtered or allowed 

to settle, and the clear solution is analyzed. Particles ir the colloidal 

size range may still be present, although the usual filtration, through 

membranes having 0.45Mm maximum diameter pores, will remove 

particles capable of scattering light. A few milligrams of colloidal 

material per liter could remain in the filtrate but will rot be visible. 

This material is generally an insignificant factor compared with concentrations of the major dissolved constituents. The colloidal fraction 

may, however, be a very large factor in the determination of minor 

element concentrations. The filter porosities used here were 0.45jum, 

the ones most commonly used for filtration of river water samples, 

and O.lOjum, used in previous studies of aluminum chemistry. Although 

filters with smaller pores are available, previous work (Hem and Roberson, 1967, p. 24) showed that they did not remove a significantly 

greater amount of aluminum hydroxide from laboratory solutions 

than the O.lOjum filters.

It is obviously impossible to be certain that any filtration technique 

would remove all colloidal particles. Electron micrographs of synthetic material made for this study showed many particles with 

diameters near O.lOMm, and the results of analyses of filtrates suggest 

that the O.lOAtm filter is capable of removing particulate material 

considerably smaller in size than O-lO^m.

The analyses of waters of major rivers of the world published by 

Durum, Heidel, and Tison (1960) show high and erratic aluminum 

concentrations for some streams. The samples for the determination of 

aluminum and other minor metals were filtered througH a relatively 

porous medium and then acidified for preservation before the spectrographic determinations were made. This treatment resulted in reporting the aluminum content of colloidal particles as dissolved 

aluminum in some of the samples.

Another large body of data containing many spectrographically 

determined aluminum values are the analyses of the water supplies of 

the 100 largest United States cities published by Durfor and Becker 

(1964). The analyses for raw water samples taken from surface-water 

supplies also show effects of colloidal material. The ground-water
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E35

samples should have been more reliable in this respect, and some 

analyses appear to give reasonable results. Treatment of a public 

water supply before putting it into the distribution system tends to 

upset any aluminosilicate equilibria that might have existed before 

treatment. The process of coagulation with alum seems almost invariably to cause a measurable increase in dissolved aluminum content, probably because colloidal or polymeric aluminum hydroxide 

species are formed that can pass through the filters. Adjustments of pH 

are commonly made also during treatment processes.

Presumably spectrographic aluminum determinations include all 

species present. Regular wet chemical analyses for aluminum are 

relatively insensitive and require very careful work to obtain results 

at concentrations below 0.1 or 0.2 mg/1. This low concentration range 

is the one to be expected in most natural water near neutral pH, if the 

solubility calculations given in this paper are applicable. Most of the 

available analytical data report aluminum determined by wet chemical 

methods and cannot be relied upon closely enough to give a fair test 

of our solubility calculations.

A further shortcoming of most published analyses is that the pH 

reported was usually determined after the sample had been brought to 

the laboratory and commonly after days, weeks, or even months of 

storage. Such pH values may not apply to the water in its original 

environment.

After careful consideration of their reliability and the likelihood of 

their representing systems where aluminosilicate equilibria might occur 

about 50 published analyses were selected from three principal sources, 

and pertinent data from these supplemented with data from a few 

analyses especially made for this study are given in table 7. They provide what is believed to be a valid test of the solubility calculations. 

Most of the analyses were taken from the study by Feth, Roberson, 

and Polzer (1964) of water that had been in contact with the granitic 

rocks of the Sierra Nevada in California and Nevada. The sampling 

and analytical work done for this study were very carefully conducted 

to obtain the best possible values for aluminum, and pH determinations 

were made at the sampling site. The pH range covered by these samples 

was from near 8.0 to about 5.5. Observations made by these investigators indicated that most of these waters were associated with kaolinitic 

or amorphous clay mineral species. A few analyses of waters with 

higher and lower pH values were obtained from the compilation of 

White, Hem, and Waring (1963) of analyses of ground waters from 

various geologic environments in which clay minerals are considered 

likely to be present. A few analyses of acid waters from various depths 

in three mines in the Pennsylvania anthracite region were published
E36 CHEMISTRY OF ALUMINUM IN NATURAL WATER

TABLE 7. Analytical data from actual water samples

[Source of data: 1-39: Feth, Roberson, and Polzer (1964); 40-45: White, Hem, and Waring (1963); 46-48: 

Barnes, Stuart, and Fisher (1964); 49-51: unpub. data, U.S. Geological Survey files]

No.

1

9

3

4

5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

20

21

99

23

24

25

26

27

^8

29

30

31

32

33

34

35

36

37

38

39

40

41

42

43

44

45

4fi

47

48

49

5(1

51

Source

Ephemeral springs:

1,218 

1,222...---.----.---.-.-.- . ... .

1,226......_ -.--....---.............

1,233 

1,256. . 

Quartz monzonite springs:

1,138 

1,217----.---.--..-----..-.-.......

1,250.. ......................_._....

1,255.. .--..---_-------_...._...-...

1,257 -- -- .. -----------

1,224 .

1,529-..----------------.---..--.-.

1,561 -- 

1,565.-----... --..-...._.--._.-_.__.

1,569 

Graiiodiorite springs:

1,170...---------------.-...--......

1,172 _

1,572 (resampled),-. ........ ....___-

1,286--.---.-----------.-.-...-...-.

1,287--.---------.-----------------.

1,293_- - -- - --------- ------------

1,294...............................

1,295 . 

1,566.-. .....-.-----....._.-.--_-...

1,567----... ---------------- -----

1,573 ----------------------------

1,574.. ....... -.-..-.....---------

Granitic rock springs:

811-- - 

1,147--..-. --------- ------------

1,246.--------------.-----....-----.

1,247.-----.---.--------.-----.-.--.

1,248 ------------ --. ... ..-.--

1,249. 

1,251 . 

1,252.----....------------....----.

1,259.........-------------.-----.

1,541.----------------.--.--..-----.

1,542.----.-.-.-.-.-----.-------.--.

1,543..------ -----------------------

Well, Park Lake, Ky-_. .............._

Well, Montieello, Ark_____. _______._

Well, Lake Roland, Md . ............

Spring Yellowstone Park upper basin.

Sisjodandi Spring, Iceland............

Well, Mexican Water, Ariz. ............

Mine water, Storrs No. 1 shaft- -.--.-.

Mine water, South Wilkes-Barre No. 5- . Mine water, Loree No. 2 shaft. .........

Mattole River, Calif., filtered 0.10/um. _ .

 do . _ ... . .-

Conklin Creek, Calif., filtered 0.10/um--

Date of

sampling

5- 2-60

5- 3-60

5- 4-60

5- 5-60

6- 3-60

10-21-59

5- 2-60

6- 2-60

6- 2-60

6- 3-60

5- 4-60

8-22-61

9-26-61

9-27-61

10-17-61

12-11-59

12-12-59

10-17 61 8- 9-60

8- 9-60

8-11-60

8-11-60

9-27-61

9-27-61

10-18-61

10-18-61

8-29-58

11- 8-59

6- 1-60

6- 1-60

6- 1-60

6- 2-60

6- 2-60

6- 2-60

6- 3-60

9-14-61

9-14-61

9-14-61

1- 5-55

12-13-55

3-19-54

10-16-57

8-31-58

3-11-55

11- 8-69

11- 8-69

11- 8-69

I

0.0002

.0003

.001

.0004

.0003

.001

.002

.002

.0003

.0002

.002

.003

.003

.0005

.0005

.001

.003

.003

.0004

.0002

.005

004

.0007

.001

.0009

.0004

.0006

.002

.001

.002

.001

.0007

.001

.001

.002

.0002

.002

.002

.002

.004

.003

.007

.02

.015

.007

.05

.1

.1+

.001

.001

.001

CAI

(mg/1)

0.04

.12

.04

.08

.02

.02

.02

.02

.02

.04

.02

.06

.03

.02

.03

.01

.02

.02

.02

.01

.02

.03

.01

.03

.02

.03

.02

.05

.01

.03

.02

.02

.02

.03

.04

.02

.19

.04

.02

1.5

28

2

.9

.64

.2

24

20

143

.10

.08

.09

Csi0

(mg/1)

24

13

37

17

15

23

35

22

20

11

24

57

51

26

29

27

20

20

24

11

39

43

28

16

13

15

23

26

19

19

15

14

24

17

19

13

25

24

20

22

98

31

363

359

11

25

49

85

7.8

8.7

13

(P[A1] or

[A1(OH) 4])

9.23

i 10. 79

9.18

i 10. 08

9.75

9.56

9.38

9.59

i 9.88

9.60

9.56

8.69

9.07

9.57

9.29

9.78

9.63

9.63

9.54

10.81

9.46

9.30

11.21

9.54

i 10. 09

9.57

9.57

9.11

9.93

9.46

9.75

9.78

9.54

9.50

9.35

10.99

8.55

9.26

9.62

8.43

6.55

8.47

6.01

6.21

7.92

7.11

6.99

5.95

9.32

9.37

9.14

I PH

7.5

5.6

2 6.9

5.9

6.3

2 6.9

6.2

7.9

6.2

6.9

7.0

7.45

7.55

6.5

6.7

= 7.1

26.8

7. 7

27.0

5.9

6.4

6.3

5.. 5

7.6

6.2

7.3

7.25

27.6

27.6

7.9

7.5

7.4

7.7

7.3

7.4

5.6

7.90

6.65

6.60

2 4.9

24.0

- 8. 3

29.6

29.7

29.2

3.95

4.20

3.92

7.42

7.60

7.80

i Effects of A1F complexes included. 

- pH measured in laboratory.

by Barnes, Stuart, and Fisher (1964), who postulated a pH control 

in these systems exerted by reversible solution of kaolinitic clays. 

Three analyses representing the deepest waters from three mines were 

also included.

The activity of [H4SiO4] was calculated directly from the analytical 

data, assuming an activity coefficient of 1.0 and correcting for the 

effect of dissociation at high pH if necessary by means of equation 8
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E37

in table 2. The reported pH was considered to represent the activity 

of hydrogen ions. Most of the measurements were made at temperatures between 10° and ]5°C. The solubility calculations refer to 25°. 

Although this discrepancy can be expected to cause some disagreement between observed and calculated solubilities, the effect should 

not be great enough to be of major concern at this point.

Figure 10 is a plot of the sum of activities of aluminum ([Al+3 ] or 

[A1(OH) 4 ~] only) and silica against pll. Dashed regression lines shown 

are the same as the ones established from laboratory data, plotted in 

figures 2 and 4.

In genera] the aluminum and silica activities for the Sierra Nevada 

springs follow the predicted pattern. Some of the individual analyses 

plot as much as one log unit away from the line, but most are much 

closer. Notes in the table of analyses from which the data came (Feth, 

Roberson, and Polxer, 1964, p. 4-15) show that kaolinite-type as well 

as X-ray amorphous clay minerals were present in association with 

the water at most sites. The experimental solubility data reported 

here for the synthetic halloysite thus appear to have some value as 

indicating the aluminum concentrations likely to be found in natural 

weathering solutions, where an equilibrium involving amorphous claytype minerals could be expected.

The activities of uncomplexed aluminum, [Al+3] for acid waters and 

[A1(OH) 4 ~] for all waters whose pH was 5.6 or over, were calculated 

from the analytically determined A1(2A1) by means of the equation

, [A10H+2] [A1(OH) 2 -1 [Al(OH)r],

~T~ 0.9 i , ~r _ " TAKOHV 7Ai(OH) 4

in combination with equations 2, 3, and 4 in table 2. Because pH and 

SA1 are known, for a given ionic strength there are four equations and 

four unknown variables. The final working expressions for ionic 

strength 0.001, for example, are

and

[Al(OH)r]=SAl/(1.04+1012 - 33 [H+] 2 +1017 - 13 [H+]3+102220[H+]').

For most of the natural waters, fluoride complexing effects are insignificant compared with the hydroxide complexing (Hem, 1968, p. 29), 

but if necessary the expressions can be expanded by adding ter-ns for 

fluoride complexes and using equations 10-15 (table 2) to evaluate 

them.
5.0

g

 a7 7.0

T

+

I? 8'°

i

0

^ 9>0

T 

" 10-°

en

-2 11.0

'

12 0

 i i i i 

\

\ A

\

*\ 

 A \ 

\ 

\ 

- \ 

\ 

\ 

, \ _^ 

\ ^o °° ^ 

\ °° 0 %8 oo 

\ -o- ^^ o 

\ -o S/^

Ox^" o

_ ^6- 

o

: i ill 

's.O 4.0 5.0 6.0 7.0 

PH

FIGURK 10. Relationship of log [Al] or log [A1(OH) 4]+ log [H4SiO4 ] to pH 

surface-water sources: Sierra Nevada perennial springs, circle; Sierra Nevada e

ground water, dot; filtered Mattole River water, x; filtered mine water, triangl
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E39

For reasons already cited, the solubility graphs should not be expected to apply to all natural water. There usually is no way of knowing if such systems are at equilibrium and if the necessary solid products are present. Thus, many analyses will yield points that are off 

the equilibrium line, either above or below saturation for the measured 

pH.

Because most analyses are riot reliable indicators of actual dissolved 

aluminum and the pH reported also may be unreliable for equilibrium 

calculations, no extensive further study of existing analytical data 

seems justified. However, at extreme values of pH, either substantially higher or lower than the 5.5-8.0 range represented by the Sierra 

granitic water, aluminum is more soluble, and existing analytical data 

may be more significant. There are a few analyses of ground waters in 

a compilation by White, Hem, and Waring (1963) which are shown by 

solid dots in figure 10. The solutions more alkaline than pH 8.5 seem 

generally to contain less aluminum than would be predicted at 

equilibrium. It seems likely the pH of such waters might be controlled 

by some reaction other than aluminosilicate dissolution. Dissolution 

of clay minerals in mildly alkaline solutions is likely to be extremely 

slow. Hence their failure to attain equilibrium is not surprising. The 

two points above pH 9.5 represent water from thermal springs, and 

solubilities at these high temperatures probably are greater than 

would be expected at 25°.

The solutions at low pH ought to attain equilibrium more readily. 

Analyses for two ground waters from fine sediments reported by White, 

Hem, and Waring (1963) give points very near the theoretical solubility line for pH 4.0 and 4.9.

Another type of acid solution which might be expected to equilibrate 

with clay minerals is water affected by pyrite oxidation. Waters of 

this kind occur in abundance in coal mines, and analyses for such 

solutions are relatively abundant. Three analyses of this type of water 

were taken from a paper by Barnes, Stuart, and Fisher (1964). 

They represent water from three mine shafts in eastern Pennsylvania. 

The observed pH ranged from 3.9 to 4.2, and substantial amounts of 

aluminum were present. Barnes, Stuart, and Fisher (1964, p. 5) 

ascribed a buffering effect in these water bodies to reactions of H+ with 

silicates and clays.

Although the deepest water sampled in each of the three shafts gave 

points that plot near the theoretical solubility line, the significance 

of the results is open to some question. The mine waters were high in 

sulfate, iron, calcium, and magnesium. As Hem (1968) pointed out, the 

complexing effect of sulfate may influence aluminum solubility. The 

calculations of aluminum activity in the mine waters, however, were 

made without considering sulfate complexes. The availability of
E40 CHEMISTRY OF ALUMINUM IN NATURAL WAT^R

sulfate for entering into aluminum complexes in these solutions cannot 

be readily calculated, as there are other sulfate complexes with Ca+2 , 

Fe+2 , and Mg+2 , and these will compete with aluminum for the 

available sulfate.

If the complexing effect of sulfate is considered, the points calculated 

for the mine-water samples will lie near the theoretical solubility for 

a kaolinite having a standard free energy of formation of 903 kcal 

per mole. The laboratory studies of Polzer and Hem (1965) in which a 

natural kaolin was placed in distilled water acidified to a pH between 

3.5 and 4.0 suggest that a reversible equilibrium with this more stable 

kaolinite is possible at low pH. In any event the data from their 

experiments also lie near the solubility referred to previously for the 

mine water. This solubility line is not entered in figure 10 but would 

parallel the left-hand segment of the dashed line, displaced downward 

by two log units.

CHEMICAL COMPOSITION OF PRECIPITATES

The amounts of aluminum and silica which disappeared from the test 

solutions were determined from the analytical data and are indicated in 

tables 3 and 4. The composition of particulate material is not very 

closely indicated by this type of information, however, because each 

of the aging solutions was analyzed many times during the aging 

period and very little solution was left at the time of the f nal observation. Nevertheless, the average of mole ratios of silica to aluminum in 

the solid materials was near 1.0 for all the observations in table 3.

Analyses of recovered solid material were made by wet chemical 

procedures for many of the aging solutions. These also usually showed 

a mole ratio near 1.0. Some of the material filtered out contained an 

excess of aluminum that is, the aluminum to silica ratio was greater 

than 1. These solids probably were mixtures of aluminum hydroxide 

and 1:1 aluminosilicate.

CRYSTALLINITY OF SOLIDS

Crystal structure of the solids synthesized in this studT^ was investigated by X-ray diffraction and electron microscopy. As noted earlier 

in the discussion, the materials recovered from solutions given in 

tables 3 and 4 did not give identifiable peaks in their X-ray diffraction scans, with the single exception of solution 12E. A peak characteristic of quartz appeared in the diffractogram for the solid recovered 

from solution 12E after three weeks of aging and was present in all 

later scans of the precipitate up to the termination of the experiment 

after 6 months. The occurrence of crystalline silica in this particular 

solution is anomalous, and several attempts to duplicate the result
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E41

by preparing similar solutions were not successful. At the high pH 

(initially 9.85) of this solution, the amount of silica retained in solution probably is near the equilibrium solubility of quartz, but no 

particular significance is attributable to this single occurrence.

Where the initial silica concentrations were below about 9.0 mg/1, 

the diffractograms of solids produced ultimately showed characteristic peaks for aluminum hydroxide species. Gibbsite patterns appeared 

in material recovered from solutions whose pH was below neutrality, 

and bayerite and nordstrandite could be identified in material aged 

in alkaline solutions. These observations are similar to results described 

earlier (Schoen and Roberson, 1970) for solutions where silica was 

absent. The presence of small amounts of silica, however, does slow 

down the crystallization process. Where silica was absent, material 

giving characteristic X-ray diffraction peaks was formed in a few 

days or weeks of aging. Where silica was present, even in concentrations as low as 2.0 mg/1, the development of crystals that would give 

X-ray diffraction peaks required 2 months or more of aging. Data 

for solutions that precipitated crystalline aluminum hydroxide are 

not included in tables 3 and 4.

Failure of the precipitates of aluminosilicate material to give X-ray 

diffraction peaks does not mean complete absence of crystallinity. 

The individual crystals may have been too small, or perhaps the 

better organized particles were too few in number to be detected by 

that technique. A considerable amount of information on the degree 

of organization of the precipitates was obtained by means of the 

electron microscope. Particulate material from all the solutions was 

examined this way. The results showed that many of the precipitates 

contained aggregates of small crystals, and some forms showed layering and other general resemblances to clay minerals when viewed 

by the electron microscope. Examples of the micrographs obtained 

are given in figures 11-22.

Figures 11-14 represent materials of known composition which also 

could be identified by X-ray diffraction. Figure 11 represents gibbsite 

obtained from a solution that contained no silica. Figure 12 represents 

bayerite obtained from a solution that contained a small amount of 

silica. The bayerite crystal habit, as noted by Schoen and Roberson 

(1970), displays a tendency for stacking of the A1(OH)3 sheets in the c 

direction rather than lateral growth to form the thin plates of gibbsite. 

The bayerite crystals in figure-12 are of a shape generally similar to 

those portrayed by Schoen and Roberson (1970, p. 59) and are similarly oriented.

Figure 13 represents two particles of a standard Montana kaoJinite. 

Typical micrographs of kaolinite that is well crystallized show abundant layered crystals with rather irregular, generally hexagonal
E42 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 11. Electron micrograph of microcrystalline gibbsite synthesized at 

25°C. Silica absent; pH of solution near 5.00; aged 18 months.

outlines (Beutelspacher and Van der Marel, 1968, p. 53-61). The 

particles in figure 13 are relatively thick, and part of their somewhat 

rounded appearance results from the carbon shadowing given the 

specimen before it in the electron beam. There was also a tendency for 

some of the specimens to deteriorate while being subjected to the 

electron beam.

Figure 14 is an electron micrograph of montmorillonite (Arizona 

bentonite) showing the very thin plates characteristic of this mineral. 

Nothing resembling this structure was seen in any of the synthesized 

aluminosilicates, which is considered to be further evidence that 

these materials have 1:1 rather than 2:1 silicon to aluminum ratios.

Figure 15 represents synthesized aluminosilicate from solution 12E 

(table 4). The layered nature of the material and its generally hexagonal structure and appearance resemble kaolinite. Very few units 

this well developed were seen.

Polzer, Hem, and Gabe (1967) published electron micrographs of 

aluminosilicates synthesized early in this study. Among the materials 

observed were some lathlike or tubular particles similar to halloysite 

portrayed by Beutelspacher and Van der Marel (1968), p. 62-68. 

It is characteristic of this mineral that the thin sheets roll up into tubes.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E43

FIGURE 12. Electron micrograph of crystalline bayerite from solution 14E. 

Aged 420 days; pH, 9.15; Al, 0.52 mg/1; SiO2 , 1.4 mg/1.

Although tubular material was observed from time to time in the 

synthesized solids, it was not generally abundant. Figure 16 shows 

typical material of this type, obtained from solution 17D.

More commonly the solids recovered have a more amorphous 

appearance. Figure 17 represents material recovered from solution 17A 

(table 3) that was aged only 51 days. There is some indication of 

developing layered crystals in this material, but most of the particles have rather irregular outlines. Figure 18 shows material from 

solution 17C, an alkaline solution. It too has material that shows 

layering and some degree of organization. A still more strongly alkaline solution produced the material shown in figure 19.

A few irregularly shaped crystals are present in figure 20. This 

solution (4D, table 3) was aged at a relatively low pH, comparable 

with that of 17A. If these had been some form of aluminum hydroxide,
E44 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 13. Electron micrograph of fragments of Lewistown, Mont., standard

kaolinite.

it should have been detected by X-ray diffraction. The precipitate 

from 4D was amorphous to X-rays.

One final experiment was conducted to help correlate crystallinity 

observations and solubility results more closely. In the various studies 

conducted by Belgian investigators and their associates, aluminosilicate gel was first prepared by reacting aluminum isopropoxide with 

ethyl orthosilicate. The gel was then aged under various conditions 

to obtain finally a kaolinite that gave characteristic X-ray diffraction 

peaks (for example, DeKimpe, 1969). The fact that the aluminosilicate gel prepared in this way will ultimately be converted to a 

crystalline clay mineral seems to be well established by the many 

studies that have been made by using this approach. It is of interest 

to examine aluminosilicate gel that was prepared this way after fairly 

short aging periods in solutions like the ones we have used and to analyze the solutions to compare them with the ones summarized in 

tables 3 and 4. If the solids appear similar to the ones we have prepared (figs. 15-20) and they appear to approach a similar equilibrium 

solubility, this will constitute further evidence that our synthetic 

materials are the same, essentially, as aluminosilicates prepared by 

mixing aluminum isopropoxide and ethyl orthosilicate and that they
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E45

FIGURE 14. Electron micrograph, montmorillonite (Arizona bentonite).

can be expected to evolve, given favorable conditions, to well-defined 

kaolinitic clay minerals.

Aluminosilicate gel was prepared by mixing equimolar quantities 

of the two organic compounds in an ethyl ether solution. After being 

freed from most of the organic solvents by gentle warming, 2-gram 

portions of the gel were placed in two polyethylene bottles. To one of 

these was added 2 liters of a 0.10 molar solution of sodium perchlorate 

which had been lowered in pH by addition of 4 milliliters of 0.11 molar 

HC1O4. To the other bottle a 0.10 molar solution of sodium perchlorate was added that had been made mildly alkaline by adding sodium 

hydroxide. The pH's of the two solutions were 4.67 and 8.78, 

respectively.

After 2 months the analyses of filtered solutions were as follows:

ISA pH, 4.76; Al, 0.58 mg/1; SiO2, 48 mg/1; 

18B pH, 6.13; Al, 0.03 mg/1; SiO2 , 25 mg/1.

Ion-activity products were 1010' 72 for ISA and 1012-34 for 18B. Both 

numbers are within the range given for similar solutions in table 3.
E46 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 15. Electron micrograph, solid recovered from solution 12E (table 4) 

after 199 days aging. pH, 8.16; Al, 2.3 mg/1; SiO2 , 16 mg/1.

Thus one may assign a solubility similar to that of our synthetic ' 

halloysite to the solids.

Figure 21 is an electron micrograph of material from solution ISA. 

This material appears to be the same as the tubular halloysite obtained 

in some of the other solutions. The electron micrograph of material i 

from solution 18B, figure 22, shows tubular material and a few hexagonal flakes of apparently layered material that resembles synthetic 

kaolinite.

Figure 23 is an electron micrograph of naturally occurring colloidal 

material from a floodflow in the Mattole River, Calif. A sample collected November 8, 1969 (also represented by Nos. 2, 49, and 50 in 

table 7) was filtered through 0.45jum and then through O.lOjum average 

pore-diameter filters. Material photographed was that caught on the 

final filter. Resemblance of the poorly crystalline residue to synthetic 

products in figures 17-19 is obvious.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E47

FIGURE 16. Electron micrograph, solid recovered from solution 17D (table 4) 

after 265 days aging. pH, SiO2, and Al contents essentially the same as after 

95 days aging.

Some further indication of crystallinity was obtained for some of 

the residues by means of electron diffraction. This technique permits 

focusing on what may appear to be promising centers of crystallization 

seen through the electron microscope and thus is a very sensitive means 

of detecting crystal organization. Electron diffraction patterns could 

be obtained from many of the aluminosilicate residues, and these 

patterns resembled those given in the literature for kaolinite (Siffert, 

1962). Similar patterns seem to be obtained, however, for various 

different related minerals, and we have not been able as yet to use 

these patterns for positive identification of species.

VALIDITY OF EQUILIBRIUM MODEL

A synthetic aluminosilicate amorphous in crystal form, at least 

when examined by X-ray diffraction techniques, has been prepared. 

Its equilibrium solubility has been determined in laboratory solutions 

at 25°C, and it appears to be consistent with the assumption that the
CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 17. Electron micrograph of solids recovered from solution 17A (table 3) 

after 51 days aging. pH, 4.02; Al, 22 mg /I; SiO2 , 181 mg /I.

material is stoichiometrically equivalent to halloysite. We have 

applied these data to a set of carefully analyzed samples of water from 

ephemeral and perennial springs in the Sierra Nevada of California and 

Nevada and have demonstrated a reasonably close fit of the aluminum, 

silica, and pH to the values predicted by the model. We have also 

predicted the effect of fluoride on solubility of aluminum in laboratory 

systems with some success.

There are some fundamental factors that need to be considered, 

however, before hopes become too high for broad application of these 

equilibrium solubility calculations. Some of these factors have already 

been mentioned. But even if the analysis for aluminum is a reliable 

indicator of the truly dissolved forms and the pH value obtained 

represents the original system accurately, the kind of equilibrium 

obtained here should not always be expected.

The synthetic halloysite prepared here may not be identical to the 

amorphous clays that are produced in weathering of granitic rock 

minerals, but it is not unreasonable to expect that their behavior
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E49

FIGURE 18. Electron micrograph of solids recovered from solution 17C (table 4), 

after 106 days aging. pH, 8.17; Al, 1.4 mg/1; SiO2 , 2.2 mg/1.

would be similar. In his study of weathering of a granitic rock in 

Maryland, Wolff (1967) did identify halloysite, both by X-ray 

diffraction and electron microscopy of the materials produced. The 

processes that occur in weathering in the Sierra may resemble somewhat the ones which occurred in the laboratory when the aging study 

solutions were prepared. An aggressive dilute runoff water perhaps 

fortified in dissolved CO2 , and thus in available H+ by the effects of 

vegetation, attacks the rock minerals and releases silica and aluminum, 

and at or near the mineral surfaces, conditions in the solution may be 

highly favorable for these solutes to combine into colloidal aluminosilicate. Most igneous rock minerals contain more silicon than aluminum, and surplus silica remains in solution or perhaps is temporarily 

precipitated as a relatively soluble amorphous form of silica. There is
E50 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 19. Electron micrograph of solids recovered from solution 11J (table 4) 

after 587 days aging. pH, 8.90; Al, 8.0 mg/1; SiO2 , 39 mg/1.

good evidence from studies by Kennedy (1971, p. 118) and by others, 

of the chemical composition of runoff, that a stable silica concentration is reached after a rather brief time of contact of the water with 

soil and rocks.

Although only indirect evidence is available for postulating the 

behavior of aluminum in runoff water, it appears likely that the 

aluminum released in the breakdown of rock minerals such as feldspar 

is rather quickly recombined into clay-type structures. The sixcoordinated structure of aluminum in the gibbsite or clay mineral 

sheet is probably established as soon as the aluminum ion is freed from 

its position in the feldspar, as this is the required configuration of 

the aquoaluminum ion and its cationic hydroxide complexes. Our 

previous studies (Hem and Roberson, 1967) showed that these ions 

polymerize rapidly enough at pH 5 or 6 so that a colloidal form of 

gibbsite could develop in the weathering solution very near the 

interface between the fresh mineral and the water.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E51

FIGURE 20. Electron micrograph, solid recovered from solution 4D (table 3) 

after 18 months aging. pH, 4.26; Al, 3.0 mg/1; SiO2 , 86 mg/1.

For a clay mineral structure to form, the tetrahedral silica layer 

must become bonded to the aluminum hydroxide layer. This bonding 

is through oxygen rather than hydroxide ions that is, the hydroxide 

ion in the gibbsite layer that is most closely approached by a silicon 

ion must become a divalent oxygen ion. The other ions coordinated 

with silicon must also be oxygens rather than hydroxides.

One obvious possible cause of the difficulty of synthesizing clay 

minerals from aqueous solutions of Al+3 and Si+4 is that dissolved forms 

of aluminum have OH~ or water molecules in all coordination positions and that silicon will occur similarly as Si(OH) 4 (aq). Incorporation of aluminum-oxygen-silicon bonds in any structure produced 

from encounters among these solutes requires removal of a proton and 

an OH~ from the site of this rearrangement of bonding. Similarly the 

establishment of silicon-oxygen-silicon bonding in the silica layer requires elimination of an OH~ and an H+ at each site. Some mechanisms 

for accomplishing these changes can be visualized in natural systems 

for example, the conditions near the water-solid mineral interface may
E52 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGURE 21. Electron micrograph of solids recovered from solution ISA after

7 days aging.

favor deprotonation and dehydration, or the silicon may be incorporated into the clay mineral structure with some of its original bonding to oxygen and aluminum relatively intact. In a dilute solution of 

aluminum and silicon species at 25°C, however, it is easy to see why 

the formation of a clay mineral structure is difficult to accomplish. 

Successful kaolinite syntheses at low temperatures reported in the 

literature, as has been noted, commonly have utilized as the starting 

materials aluminum propoxide, A1[OCH(CH3) 2]3, and ethyl orthosilicates, Si(OCH2CH3) 4 . In both these compounds there is a cationoxygen-carbon bonding structure rather than a cation-OH bonding. 

The encounter between the complexes provides a good chance for 

setting up aluminum-oxygen-silicon bonding.
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E53

FIGURE 22. Electron micrograph of solids recovered from solution 18B after

7 days aging.

Linares and Huertas (1971) reported a kaolinite synthesis using a 

"fulvic acid" complex of aluminum and suggested the organic complexing of aluminum in natural systems may be important in the 

natural processes by which clay minerals form. Such a reaction path 

does not seem to be followed, however, in all natural systems where 

clays are formed. Further attention to, and study of, the details of 

aluminum and silicon associations should shed considerable light on 

this poorly understood but very important weathering reaction.

As noted, the analytical data for samples of Sierra Nevada springs 

fit the postulate of an equilibrium with respect to synthetic halloysite. 

It may be supposed that reactions occurring between snowmelt and 

feldspar surfaces could produce conditions somewhat like those in 

our laboratory solutions and that once established the equilibrium 

would be maintained in the solution so long as the water does not move
E54 CHEMISTRY OF ALUMINUM IN NATURAL WATER

FIGUKE 23. Electron micrograph of solids passing 0.45/im filter caught on 0.10/um 

filter. Water from Mattole River near Petrolia, Calif., November 8, 1969.

into a substantially different chemical environment. Of all the solutions represented in figure 10, some of the best fits were among the 

sources listed by Feth, Roberson, and Polzer (1964) as "ephemeral" 

springs. These are characterized by short contact times between water 

and rock and presumably would represent systems where there is little 

chance for the solution to enter a different environment that could 

upset the equilibrium and where the early stages of rock-water reactions 

characterized by Carrels and MacKenzie (1967) are still predominant. 

The "perennial" springs, where contact times and flow paths are 

longer, include some waters that do not fit the equilibrium model as 

closely.

Another point of significance in considering aluminosilicate equilibria of this kind is the fact that in most natural water this type of 

reaction will be subordinate to others that involve larger stoichiometric 

amounts of reactants. For example, as noted by Wollast (1967), 

the rate and intensity of weathering of a feldspar are likely to depend
INTERACTIONS OF ALUMINUM WITH AQUEOUS SILICA E55

on the hydrogen-ion supply, derived from dissolved CO2 and maintained by a carbonate buffer system. The aluminum and silica concentrations in the solution may well be expected to arrange themselves 

at values predicted by the equilibrium model so long as the H+ 

activity remains high enough that aluminum and silica are extracted 

from the original rock minerals. A decrease of H+ activity (rising pH) 

may tend to slow the attack. At relatively alkaline pH, even though 

one might calculate a rather high solubility for synthetic aluminosilicate, the reaction path whereby the previously formed material 

can be dissolved is likely to contain a high-energy barrier. The mechanisms for disruption either of aluminum-oxygen-silicon bonds or 

of double OH bridges between aluminum ions in a solution of high 

pH must be substantially different from those producing such disruptions at low pH.

The introduction of fluoride into a system containing synthetic 

halloysite would be predicted to increase aluminum solubility. Ag^in, 

however, the mechanisms for substituting F~ for OH~ in these structures probably involve slow steps and energy inputs.

It thus seems possible the aluminum, silica, and pH prevailing in 

a solution near the original site of synthesis of a clay mineral would 

be near predicted equilibrium values. An increase in pH brought 

about by some change in the bicarbonate buffer system, however, 

might not be balanced by increases in dissolved aluminum and silica, 

either because of slowness in reactions or lack of sufficient aluminosilicate reactant.

Thus the aluminum concentration in a ground water may tend to 

reflect the condition in the system at some other pH than the one 

observed at the sampling point. Examination of published water 

analyses from many sources suggests most strongly alkaline or high 

fluoride ground waters contain much less dissolved aluminum than the 

equilibrium solubility limit.

In spite of the many reservations that must be made in using 

equilibrium approaches to low temperature aqueous aluminosilicate 

chemistry, a beginning toward a better understanding of such systems 

in weathering solutions can be made in this way. With the aid of these 

models it is possible to deduce explicit areas in which research can 

be directed to gain more knowledge. For example, more study of mechanisms that are involved in establishing inter] ayer bonding in clay 

minerals is definitely necessary.

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