Please do not destroy or throw away this publication. If you have no further use 

for it, write to the Geological Survey at Washington and ask for a frank to return it 

UNITED STATES DEPARTMENT OF THE INTERIOR 

Harold L. Ickes, Secretary 

GEOLOGICAL SURVEY 

W. C. Mendenhall, Director 

Professional Paper 185-A 

STUDIES ON THE ALK_ALINITY OF SOM.E 

SILICATE MINERALS 

BY 

R. E. STEVENS 

Snorter contributions to general geology, 1934-35 

(Pages 1-13) 

UNITED STATES 

GOVERNMENT PRINTING OFFICE 

WASHINGTON : 1934 

For sale by the Superintendent of Documents, Washington. D. C. Price 5 rents 
CONTENTS 

Pagr 

Abstract---- ------------------------~------ ----------- ------- ------------ ---------------- ----- ---- ------- -- 1 

Introduction___ ____________________________________________________________ _____ _______ __ _________ ______ __ ___ 1 

Usefulness of measures of alkalinity in geologic studies_______ ________________ ______ _____ ____ ____ ______________ 1 

Acknowledgments__ _________________________ ____ _________ _________________ ____ _____ ______________________ 2 

Review of the literature___ __________ __ _____ ___ _______ ____________________ __ _________ _________ __ __ _________ 2 

Chemical theory of pH measurements______________________ ___ __ _________________________ ________ ___ ___ __ __ _____ 3 

Colorimetric determination of pH____ __ __ ________________ _______________________ _____ ___ ___ _________ __ _____ ____ - 4 

Method of preparing the mineral solution__________ ____________ ___________________ ____ ______ ____ __________ __ _ 4 

Neutralizing action of mineraL _______________ ________ ____ _________________ :... _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 4 

Details of colorimetric method____ ____________________ ____________________ ___ __ ___ ____ __ _____ ______ ____ ___ _ 4 

Method of color comparison_ _______________________ _____ ________ ________ ___________ __ ______ __ _____ ___ __ ___ 4 

Selection of indica tors _________ _______ ________ ____________ :. _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 4 

Selectionofbuffers____ ____ ________________ ________ ___________ _____________ ____________ _____ ______________ 5 

Results of colorimetric determinations_______ __ _______ _______________________________________________________ 5 

Electrometric determination of pH _________ _____ _____ _____________________________________________ .______ _ _ _ _ _ _ _ _ 5 

Kind of electrode used ___________________________ ____ ___ ___ ___________ _______ __ ~_ _________________________ 5 

Source and purification of the hydrogen ________________ __ __________ ____ ___ ___ _____ ________ ___ _______________ 6 

Assembly and manipulation of apparatus____ __ _________ __ _________ __________________ ________________________ 6 

Effect of carbon dioxide_____ _____ _____ _________ ___________ ___ ______________ ________ _______________________ 7 

Time of grinding _____ :... _______ ______ _________________ ____________________ __ ____ .- _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 7 

Effect of agate mortar ____ ____________________________________________ ___ __________ _____________________ -: _ 8 

Effect of iron in mineraL ____ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 8 

Effectoftemperature___ _________________ _____ _______ _____ ________________ ________________________________ 8 

Results of electrometric determinations_____ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ 8 

Comparison of colorimetric and electrometric methods and factors-involved in testing__ _____________ __________________ 9 

General applications _____________________________ .., ___ ____ _____ ____ ____ _______________ _________________________ 10 

ILLUSTRATION 

Page 

FIGURE 1. A, Microvessel and electrode; B, Assembly of apparatus for pH determination_ ______ ____ ______ ___ __ ________ • 6 

II 
SHORTER CONTRIBUTIONS TO GENERAL GEOLOGY, 1934-35 

STUDIES ON THE ALKALINITY OF SOME SILICATE MINERALS 

By R. E. STEVENS 

ABSTRACT 

By grinding minerals under water it has been found that they 

yield relative and reproducible measurements of the hydrogenion concentration resulting from their hydrolysis. Many silicate 

minerals and two glasses have been studied · in this way by a 

colorimetric method of determining hydrogen-ion concentration, 

and some of t:p.em have been studied more quantitatively by 

means of the hydrogen electrode. 

The results of these tests are a rough index of the weathering 

qualities of different silicate minerals. The effect of silicate 

minerals on underground waters and upon geochemical changes 

is discussed. 

INTRODUCTION 

USEFULNESS OF MEASURES OF ALKALINITY IN GEOLOGIC 

STUDIES 

The action of pure and carl?,onated waters in decomposing silicate minerals has long been observed in 

the field and confirmed by laboratory experiments. 

Field observation has told much as to the relative 

stability of n:J.inerals toward water. The micas, especially muscovite, are highly resistant, being present 

as end products of rock decomposition, and the stability 

of orthoclase has also been observed. Plagioclase feldspars are somewhat more easily attacked, and the more 

basic mineral constituents of rocks, such as olivine, 

pyroxene, and the amphiboles, in , comparison, yield 

readily. 1 

This action of water on rocks has been studied by 

numerous methods in the laboratory, and in a general 

way the mechanism of the process is understood. The 

mineral does not dissolve as a unit but, by hydrolysis, 

yields colloidal complexes of · silica and alumina, or 

under some conditions intermediate products, such as 

calcite and magnesite, and free alkali. The solutions 

are alkaline in reaction. 

A measurement of the hydrogen-ion concentration 

of these solutions, therefore, is an index to the extent 

of decomposition taking place and to the relative 

stability of the mineral in water. It also indicates the 

potential powers of the minerals in neutralizing ground 

waters containing carbonic, humic, sulphuric, and other 

acids. 

1 Clarke, F. W., The data of geochemistry, 5th ed.: U.S. Geol. Survey Bull. 770, 

p. 476,1924. Merrill, G. P., Rocks, rock weathering, and soils, New York, Macmillan Co., 1898. 

54471-34 

The importance of silicates as a source of alkali in 

geochemical processes seems to ~e underestimated. A 

great ·part of them are capable of yielding more highly 

alkaline solutions than the carbonates calcite and dolomite. .. Furthermore, they are present in great abundance. Clarke 2 says: 

In the solid crust of the earth the silicates are by far the most 

important constituents. They form at least nine-tenths of the 

entire known mass and practically all the rocks except the sandstones, quartzites, and carbonates. 

Numerous geochemical changes take place in an environment of silicate minerals, and their effect on the 

alkalinity of solutions is of primary i~portance. 

Many-of the chen1ical reactions can take place only · 

within a certain range of alkalinity or acidity. Thus 

Wells 3 has shown that copper is precipitated from 

cuprous chloride solution by ferrous chloride only in a 

weakly acid or an alkaline solution. Copper, in the 

Michigan deposits, is associated with prehnite, a 

highly 'basic silicate, but not commonly with laumontite a mineral that is similar but contains a smallPr ' . ' 

percentage of calcium. Freeman 4 'has shown that 

sulph]de ores might be carried in alkaline solution as 

douhl~ 7 sulphides of alkali metals.· Iri the process of 

enrichment certain o,res near the surface are oxidized 

and e~~~ied downward in acid solutions. . The solutions a;re gradually neutralized or ma~e alkaline by the 

rocks with which they come into co-q.tact, causing pre- • I 

cipita'tion of the ore. · 

A statement merely that the solutions are alkaliw~. 

or even alkaline to a certain indicator, is not sufficiently definite to be useful in many chemical studies. 

The degree of alkalinity or acidity may be precisely 

stated in tenns of hydrogen-ion concentration, expressed as pH, as explained more fully in the following 

pages. Suffice it to say at this point that nearly all 

chemical reactions in solutio_ns are affected by the pH 

number, and in many of them it is a predominant 

2 Clarke, F . W ., The constitution of the natural silicates: : qeot Srirvey Bull. 

588, p. 5, 1914. 

a Wells, R. C., Chemistry of deposition of native copper from ascending solutions: 

U.S. Geol. Survey Bull. 778, p. 65, 1925. . · 

• Freeman, Horace, Genesis of sulfide ores: Eng. and Min. Jour., vol. 120, pp. 973-

975, 1925. See also Foreman, Fred, Hydrothermal experiments on solubility and 

oxidation of iron and copper sulphides: Econ. Geology, vol. 24, p, 811, 1929. 

1 
2 SHORTER CONTRIBU'l'ION TO GENERAL GEOLOGY, 1934- 35 

factor. For this reason the role of hydrogen-ion concentrations in geochemistry is an important one. 

Solutions cannot come immediately into equilibrium 

with the rocks through which they pass. As alkali is 

dissolved and removed, the action is ret~rded by a protective film of the colloidal complexes of silica and 

alumina forme~ at the same time. Only when the 

mineral is finely divided and presents a large surface 

can equilibrium be established. In view of the abundance of silicates, however, there seems little doubt 

that they affec_t greatly the pH of underground waters. 

ACKNOWLEDGMENTS 

The writer wishes to acknowledge the helpful advice 

of R. C. Wells, of the United States Geological Survey, 

throughout the investigation and in the preparation ·of 

the manuscript. The suggestions of Mr. B. D. Van 

Evera, of the faculty of George Washington University, are gratefully acknowledged. Samples of laumontite and prehnite were obtained through the courtesy of Mr. E. P. Henderson, of the United States N ationa! Museum. Two samples of glass, analyzed by 

Mr. F. W. Glaze, were received from Mr. Conrad 

Faick, of the Bureau of Standards. 

REVIEW OF THE LITERATURE 

The decomposition of minerals by pure and carbonated water has long been known and studied. As 

early as 1848 W. B. and R. E. Rogers 5 tested many 

minerals and found them all at least slightly attacked. 

Daubree 6 showed that orthoclase is attacked by 

water, and more actively when carbon dioxide is 

present. Kenngott 7 found that many silicates give 

alkaline reactions to moistened test paper. Muller 8 

showed the effectiveness of carbonated waters in their 

attack on minerals and analyzed the dissolved material. His analyses show that the alkaline constituents are predominantly dissolved. 

Hoffman 9 determined the relative solubility of 

silicates by ·their loss in weight, but the . accuracy of 

this procedure is questionable-. The works of Johnstone 10 and of Dolter 11 show the action of carbonated 

water. 

~Rogers, W. B. and R. E., On the decomposition and partial solution of minerals, 

rocks, etc., by pure water, and water charged with carbonic acid: Am. Jour. Sci., 

· 2d ser., vol. 5, p. 401, 1848. 6 Daubree, Auguste, Etudes synthetiques de geologie experimentale, pp. 271-275, 

Paris, Dunod, 1879. 7 Kenngott, A., Uber einige Erscheinungen beobachtet an Natrololith: Neues 

Jahrb., 1867, p. 77. 

s Muller, Richard, Untersuchungen iiber die Einwirkung des kohlensaurebaltigen 

vVassers auf einige Mineralien und Gesteine: K.-k. geol.' Reichsanstalt, Jahrb., 

vol. 27, Min. Mitt., p. 25, 1877. 9 Hoffman, E. W., Untersuchungen iiber den Einfiuss von gewohnlicben Wasser 

auf silicate [inang. diss.], Leipzig, 1882. 

IO Johnstone, Alexand.er, On the action of pure water and of water saturated with 

carbonic acid gas on the minerals of the mica family: Geol. Soc. London Quart. Jour., 

vol. 45, p. 363, 1889. 

u Dolter, C., Einige Versuche iiber die Loslichkeit der Mineralien: Min. pet. 

Mitt., vol. 11, p. 319, 1890. 

Clarke 12 treated a number_ of powdered, silicate 

minerals with water containing phenolphthalein, 

noting the intensity of color produced, and · thus 

determined roughly the extent to ~hich they were 

attacked. Analyses by Steiger 13 show a discrepancy 

between the intensity of the color and the quantity of 

alkali found in solution by titration with methyl 

orange. No cause for this discrepancy is suggested. 

Cushman 14 has inve$tigated the mechanism of the 

decmnposition. He has shown that the alkalies dissolve, liberating colloids of silica and alumina. These 

colloids retain a portion of the separated alkali but 

give it up to electrolytic solvents. Thus, when ammonium chloride was present, 10 to 20 times as much 

alkali was freed as when pure water was used alone. 

This might have been merely a case of base exchange. 

Similar results have been obtained by AndreY 

Several papers confirm the investigations thus far 

mentioned. 16 Thugutt,17 however, believed that orthoclase was attacked little if any by water. 

Bouyoucos 18 determined the solubility of several 

minerals by the freezing-point method. . He used four 

parts of finely ground mineral (100 mesh) to three 

parts of water, and the silicates tested were amphibole, 

apophyllite, chrysolite, epidote, kaolinite, oligoclase, 

orthoclase, prochlorite, pyroxene, quartz, serpentine, 

stilbite, talc, and zircon. Solution takes place almost 

immediately, and very little change shows after 80 

days. Bouyoucos found that the extent of solution 

was somewhat dependent upon the ratio of mineral to 

water and that the solubility increased with increases 

in temperature or pressure, and also tha.t leached 

minerals were less soluble than fresh ones. Similar 

conclusions have been drawn from tests on the durability of glass by the powder method. 19 

Smyth 20 calculated from the average composition of 

igneous and sedimentary rocks the relative quantities 

of the rock constituents carried away by water. He 

12 Clarke, F. W., The alkaline reaction of some natural silicates: U.S. Geol. Survey Bull. 167, pp. 156-158, 1900. 

t3 Steiger, George, The solubility in water of certain natural silicates: U.S. Geol. 

Survey Bull. 167, pp. 159-160, 1900. 

u Cushman, A. S., The effect of water on rock powders: U.S. Dept. Agr. Bur. 

Chern. and Soils Bull. 92, 1905. 

1s Andre, G., Displacement of potassium of certain feldspars by substances used 

as fertilizers: Compt. Rend., vol. 157, p. 856, 1913. 

16 Levison, W. G., On the origin and sequence of the minerals of the Newark 

(Triassic) igneous rocks of New Jersey: New York Min. Club Bull. 2, 1909. Dumont, J ., On the chemical decomposition of rocks: Com pt. Rend., vol. 149, p. 1390, 

1909. Henrich, F., The action of water bearing carbonic acid on rocks: Zeitschr. 

prakt. Geologie, vol. 18, p. 85, 1910. Sicha, F., Untersuchungen iiber die Wirkungen 

des beim hohen Druche mit Kohlensaure ges!l.ttingten W assers auf einige Mineralien 

[inaug. diss.], Leipzig, 1891. 

11 Thugutt, S. J., The solubility of certain silicates in water: Soc. sci. Warsaw 

Conipt. rend., vol. 6, p. 629, 1914. 

1s Bouyoucos,, G. J., Rate and extent of solubility of minerals and rocks under 

different treatments and conditions: Michigan Agr. Exper. Sta. Tech. Bull. 50, 

1921. 

19 Thomas, M ., Determination of the dmability of glass by the powder method: 

Soc. Glass Tech. Jour., vol. 12, p. 87, 1929. 

20 Smyth, C. H., The relative solubilities of the chemical constituents of rocks: 

Jour. Geology, vol. 21, pp. 105-120, 1913. 
ALKALINITY OF SOME SILICATE MINERALS 3 

shows that alkalies are dissolved to the greatest extentsodium and calcium the most completely and potassium the least. 

But little work has been done on the hydrogen-ion 

concentrations of solutions from mineral silicates. 

The work of Clarke 21 was a rough measuren1ent of their 

relative alkalinity. More recently Atkins 22 has discussed the application of pH measurements to geology 

and made several measurements on rocks and minerals. 

Britton 23 has determined the electrometric precipitation curves of silicic acid, using various alkalies, and 

gives certain points that are of interest and use in the 

present article. 

CHEMICAL THEORY OF pH MEASUREMENTS 

Because the discussion of hydrogen-ion concentra~ 

tions involves principles not yet broadly applied to 

geology, it seems wise to explain briefly the fund a:.. 

mental ideas involved. 

Water ionizes to a moderate extent, yielding hydrogen and hydroxyl ions, thus: 

The extent to which this reaction occurs is slight but 

definite, and in pure water the number of hydrogen 

ions would be the same as the nurnber of hydroxyl ions. 

When other substances are present in the water, however, these two numbers are not, in general, equal. 

Under such circumstances use is made of the fact, 

based on the law of mass action, that the product of the 

concentration of H+ and OH- is a constant, Kw. 

That is, K w = [H+] [QH-] where the bracketed quantities represent concentrations or activities.· By adding 

alkali to the water the concentration of OH- is increased. In order to maintain the constancy of Kw 

a corresponding quantity of H+ must disappear from 

solution, recombining to form un-ionized water. In 

a similar manner acid increases the H+ concentration 

and removes OH- from solution. 

The value of Kw has been determined by numerous 

investigators. At 25° C. it is approxiinately 1 X 10-14 

when concentrations are in gram ion per liter. In 

pure water or a neutral solution each ion (H+ and 

OH-) is present in a concentration of 1 X 10-7 gram 

ion per liter. As these t\vo concentrations are so 

simply related it is sufficient for most purposes to 

consider the hydrogen-ion concentration alone. Such 

concentrations will be expressed as a negative power 

of 10. The figure following the negative sign is 

characteristic for each concentration and has been 

called for short the P!l number-for example, 7 is the 

21 Clarke, F . W., op. cit. (Bull. 167), pp. 156-158. 

22 Atkins, W. R. G., Some geochemical applications of measurements of hydrogenion concentration: Roy. Dublin Soc. Sci. Proc., vol. 19, pp. 455-460, 1930. 

23 Britton, H. T. S., Electrometric study of the precipitation of silicates: London 

Chern. Soc. Jour., 1927, p. 425. ' 

pH number for pure water or a neutral solution, as 

noted above. Expressed mathematically, , 

' ' 1 · log10 10-7 = -7· or 7 = -loglO 10-7 or 7 =log 10_7 , 

a~d, in general, 

In the same way 

1 

pH=log [H+] 

. 1 

pOH =log [QH-] 

At neutrality pH and pOH are both 7. A hydrogenion concentration of 10-5 is equivalent to pH= 5 and 

to p0H=9; [H+]=lo-n is the same as pH=ll and 

p0H=3, etc., in each case the sum of pH and pOH 

being 14. Thus the concentration of either ion in 

solution is given simply by pH, it being understood 

that pH= 7 signifies neutrality, that alkaline solutions are at pH greater than 7 and acid solutions less 

than 7, and that an increase in pH signifies a decrease 

in hydrogen-ion concentration or, in general, in acidity. 

It becomes obvious, from what has been said of the 

ionization of water, that the strength of an acid or a 

base is dependent on the extent to which it ionizes 

into H or OH ions in solution. Thus hydrochloric 

acid, nitric acid, and sulphuric acid ionize readily and 

are strong acids, whereas acetic acid, silicic acid, and 

others that ionize feebly are classed among the weak 

acids. Alkali and alkaline earth hydroxides ionize 

readily, giving OH ions, whereas ammonium hydroxide is more weakly ionized .. 

A salt, on hydrolysis, reacts with the ions of water 

to form an acid and a base. · A good exa.mple is the 

hydrolysis of an alkali silicate, thus: 

Na2Si03 + 2HOH --72N aOH + H2Si03 

"'-

The products are a strongly ionized base, sodium 

hydroxide (N aOH), and the weak metasilicic acid 

(H2Si03). For this reason solutions of sodium silicate 

are strongly alkaline. Silicate minerals hydrolyze in 

the same way and give alkaline solutions. 

Two methods of measuring hydrogen-ion concentrations are in common use. The direct measurement 

is with the hydrogen elec.trode. An electrode of 

platinum, covered with platinum black, dissolves a 

small quantity of hydrogen which is capable of ionizing reversibly when an electric current is passed. 

The potential of the electrode depends on the concentration of hydrogen ions in the solution being measured. 

This resulting potential is determined against that of 

a standard cal01nel half cell. The pH of the solution 

is then given by the equation 

I

J _ e.m.f. + Ebar- Ecal 

p =t- kT 
4 SHORTER CONTRIBUTION TO GENERAL GEOLOGY, 1934-35 

in which e.m.f. is the electromotive force of the cell as 

measured, Ebar is a srnall correction for the vapor 

pressure of water, Ecat is the potential of the stap.dard 

half cell, k is a constant, and T the absolute temperature. 

The colorimetric method of measuring pH is an 

indirect method used where the demands of accuracy 

are not great and as a check upon the hydrogenelectrode measurement. Organic dyestuffs are: used 

as indicators, their . color depending upon the pH of 

the solution. Phenolphthalein, for example, is colorless below pH = 8.3, shows increasing depths of red 

from pH = 8.3 to pH= 10.0, and develops the full red 

color above pH= 10.0. The determination is made by 

matching the color of the solution being examined with 

that of one of a set of standard buffer solutions 24 of 

known pH containing an equal quantity of indicator. 

(See also pp. 5, 10, and 11.) 

COLORIMETRIC DETERMINATION OF pH 

METHOD OF PREPARING THE MINERAL SOLUTION 

In order to obtain results that would be comparable 

for different minerals it was necessary to get the solution as near equilibrium with the mineral as possible 

before making measurements. The solution of silieates depends upon the surface exposed. This is 

probably the result of the protective action of colloids 

of silica and alumina. The use of high ten1perature 

and pressure is to be avoided because alkali liberated 

under such conditions reeombines but slowly on cooling with the colloidal siliea and alumina set free at the 

higher ten1perature and pressure. Continued churning of the powdered mineral for 15 n1inutes to 2 .hours 

in a Clark hydrogen-electrode vessel with water, 

followed by electrometric pH measurements, . gave 

results that were obviously low and that could not be 

duplicated in successive measurements. Repeated 

exposure of fresh surfaces to the action of water and 

the removal of proteetive films of silica and alumina 

was neeessary. 

By grinding eontinuously in an agate mortar with 

one drop of water and one of indicator solution many 

of the minerals caused phenolphthalein to turn full 

red (at least pH= 10.0). With orthoelase the depth 

of red reached a maxin1um after about 1% minutes 

of grinding. Other rninerals incapable of changing 

phenolphthalein to full red were found to cause a 

maximurn of change in the same period of time. The 

effeet of the agate mortar and pestle on the aeeur~ cy 

of the pH determination was thought negligible, 

owing to their extreme hardness and the relatively 

small surface exposed in comparison to that of the 

mineral. (See also p. 8.) 

24 Buffer solutions "have the ability to resist changes in pH through the addition 

or loss of acid or alkali." For a more comprehensive treatment see Clark, W. M., 

The determination of hydrogen ions, Williams & Wilkins 1928, 

NEUTRALIZING ACTION OF MINERAL 

11ost silicate minerals have a strong neutralizing 

capacity . toward acid solution when there is sufficient 

surface exposed to solution. Powdered orthoclase 

was ground in the agate mortar with one drop of 

phthalate buffer at pH= 5.0 and one drop of cresol-red 

indicator, which is yellow below pH= 7.2 and red 

above, with its full red color at pH = 8.8. After 3 

minutes of grinding the solution was bright red, showing that the acidity of the buffer had been neutralized 

bv alkali from the mineral. However, when orthoclase 

as ground with an alkaline buffer solution at pH= 10 

no change in pH was noted. 

DETAILS OF COLORIMETRIC METHOD 

The colorimetric determinations in table 2 were 

made by grinding the mineral with the appropriate indicator and matching the color with that of a standard 

buffer solution of known pH containing an equal quantity of the indicator. The quantity of mineral used 

had no measurable effect on the results, provided there 

was sufficient to give a heavy suspension. To the 

mineral was added one drop of carbon-dioxide-free 

distilled water and one drop of indicator solution, and 

grinding was continued for 2 minutes. One drop of indicator was then added to one drop of a standard buffer 

and the color comparison n1ade. 

. METHOD OF COLOR COMPARISON 

The turbidity of the mineral solutions made cmnparison with the clear buffer solutions difficult. The 

spot-plate method did not seem to give good results. 

Comparison in capillary Pyrex tubes, of about 1 millirneter bore, was more satisfactory. 

A column of the solution was made to rise in the 

tube by capillarity, and the tube was held upright for 

about a minute to allow mineral particles to settle as 

much as possible. Occasionally nearly clear solutions 

were obtained thus, but in most tests the solutions remained turbid. Any error due to absorption of dye by 

the mineral could be noted at the bottom of the tube, 

but none of the minerals tested seemed to cause such 

an effect. 

The tubes were laid on the edge of a table with tubes 

of standard buffers, a white background was placed 

about an inch below the tubes, and the comparison 

was then made in indirect sunlight. The buffers differed by 0.4 pH in the series, but in all cases interpolation to 0.1 pH was attempted. For certainty, colors 

were duplicated several times, and whenever possible 

checks were run with different indicators. 

SELECTION OF INDICATORS 

The indicators used were selected from those described by Clark. 25 

26 Clark, W. M., op. cit., pp. 92-94. 
ALKALINITY OF SOME SILICATE MINERALS 5 

TABLE I.-Indicators used in colorimetric determinations of pH 

Clark and Lubs 

Concen- Color change 

Indicator tration Range of pH (percent) Acid Alkaline 

Bromphenol-blue ___ 0. 02 3. 0-4. 6 Yellow ___ Blue. 

Bromcresol-purple __ . 01 5. 2-6. 8 ___ do _____ Purple. Bromthymol-blue ___ . 01 6. 0-7. 6 ___ do _____ Blut·. Phenol-red _________ . 01 6. 8-8. 4 ___ do _____ Red. Cresol-red _________ . 01 7. 2-8. 8 ___ do _____ Do. 

Thymol-blue _______ . 01 { 1. 2-2. 8 Red _____ Yellow. 

8. 0-9. 6 Yellow ___ Blue. 

Sorensen 

Methyl-orange _____ 0. 01 3. 1- 4. 4 Red _____ Yellow. 

Methyl-red ________ a. 02 4. 2- 6. 3 ___ do _____ Do. 

Phenolphthalein ____ b. 05 8. 3-10. 0 Colorless_ Red. 

Thymolphthalein c __ b. 04 9. 3-10. 5 ___ do _____ Blue. 

Alizarin-yellow R ___ . 01 10. 1- 12. 1 Light yel- Deep yellow. - low. 

Tropaeolin 0 _______ . 01 11. 1- 12. 7 Yellow ___ Orange. 

a In 60 percent alcohol. b In 50 percent alcohol. c Faded quickly; found of little use. 

SELECTION OF BUFFERS 

Buffer solutions were prepared differing by 0.4 pH 

unit and ranging from pH =4.2 to pH= 12.0, as listed 

below. 26 Their pH was checked at intervals by means 

of the Clark hydrogen-electrode vessel. 

Phthalate-NaOH ____ _____________ _ 

KH2POrNaOH __________________ _ 

Boric acid, KCl-NaOH _____ _______ _ 

Soda-borax _______________________ _ 

Na2HP04-NaOH _________ -~- _____ _ 

Range of pH 

4. 2- 5. 8 

5. 8- 7. 8 

7. 8- 10. 0 

9. 4-11. 0 

11. 0-12. 0 

RESULTS OF COLORUiETRIC DETERMINATIONS 

The results of these colorimetric measurements are 

listed in table 2, approximately in order of increasing 

pH. This grouping of the minerals is in fair agreement with field observations of their stability to water, 

and a relation can be ob$~rved between the alkali content and the pH obtained. This table is discussed 

more fully on page 9. 

Index 

no. 

TABLE 2.-Results of colorimetric tests 

Mineral Source Composition pH 

~ ~:~~it:;~~~~~~~~~~~ -~~~-~~t~~~~:=~~~~~ !~:~~~~~~~:~~6ii;0~~~~~===== ;: ~ 3 Lepidolite __________ -------------------- KLi[Al.2(0H.F)]Al(Si03)3__ 8. 4 

4 Phillipsite __________ -------------------- (K2, Ca)AbSi4012.4~2H 0___ 8.4 

5 Stilbite _____________ Mexico ____________ (Na2, Ca)AbSia016.6H2

0____ 8.6 

6 Muscovite no. L ___ Utah ______ __ ______ H2KAh(Si04)3. _ ------------ a 7. 8 

7 Muscovite no. 3 ____ ----~--------------- _____ do__ ___ __________________ 8. 4 

8 Muscovite no. 6 ____ -------------------- _____ do_______________________ 8. 5 

9 Muscovite no. 2 ____ -------------------- _____ do_________ ______________ 8. 6 

10 Muscovite no. 5 ____ -------------------- _____ do_______________________ 8. 7 

11 Muscovite no. 4 ____ -------------------- _____ do_______________________ 9. o 12 Calamine ___________ -------------------- H 2(Znz0)Si04________ __ _____ 8. 8 

13 Biotite _____________ Philadelphia ______ (H,K)z(Mg, Fe)2Ah(Si04

)3__ 9.0 

a Sericite; fine powdery variety . 

26 flee also Clark, W. M., op. cit., pp. 192-220. 

TABLE 2.-:-Results of colorimetric tests-Continued 

Index Mineral Source Composition pH no. 

14 Pollucite __________ Buckfield, Maine. H2CszAb(SiOa)s------------- 9.0 

15 Anthophyllite ______ Montana __________ (Mg,Fe)SiOa---------------- 9.0 

16 Laumontite . _______ Mordon, Nova H4CaAbSi40J4.2HzO ________ 9. 0 

Scotia . 

17 Orthoclase no. 2 ____ Maine ____________ KA.ISisOs------------------- 8.8 

18 Orthoclase no. L ___ Deadwood, S.Dak. _____ do_______________________ 9. 2 

19 Spodumene ________ Maine ______ _____ _ LiAl(SiOah----------------- 9.2 

20 Clinochlore __ _______ Chester County, HsMgsAbSiaOts- ------------ 9. 2 

Pa. 

21 Albite no. L _______ -------------------- NaAlSiaOs------------------ · 9. 4 

22 Albite no. 2 ___ _____ ---------------- _________ do _____ -------------_____ 9. 8 

23 Wyomingite (leu- Montana __________ KA1Siz06 (?)_______ ______ ___ 9.4 

cite). 

24 Labradorite nQ. 2 ___ -------------------- Na.AlSiaOs.3CaAlzSizOs (?)__ 9. 4 

25 Labradorite no. L __ -------------------- _____ do_______________________ 9. 8 

26 Margarita __________ Chester, Mass _____ H2CaAl4Si20t2--------------- 9. 8 

27 Natrolite ___________ -------------------- Na2AlzSbOJo.2HzO _ --------- 10.0 

28 Epidote ____________ -------------------- Caz(Al.OH) (Al,Fe)2(Si04)3 10.0 

Varies. 

29 Actinolite_--------- Roan Mountain, Ca(Mg, Fe)a(Si0a)4. -------- 10.0 

Tenn. 

30 Phlogopite _________ -------------------- HzKMgaAl(Si04)a (?) _______ 10.1 

31 Diopside ___________ -------------------- CaMg(Si0J)2---------------- 10.1 

32 Hornblende ________ -------------------- CaMga(Si03)4 with NazAlz 10.2 

(Si03)4+Mg2Al4(SiOa)z. 

33 Olivine .. ~------ ---- Willits, N.C ____ __ (Mg, Fe)zSiOt--------------- 10.2 

34 Thulite _____________ -------------------- HCa2AlaSia012 -------------- 10.2 

35 Talc ____ _______ ____ Edwards, N.Y __ __ H2Mga(Si0a)4--------------- 10.2 

36 Pyroxene ____ _______ -------------------- Ca(Mg, Fe) (SiOa)z ___ ______ 10.2 

37 Tremolite no. L ____ -------------------- H2Ca2Mg5(Si0als.----------- 10.2 

38 Tremolite np. 2 _____ -------------------- _____ do. ______________________ 10.2 

39 Pectolite ___ -~------ -------------------- HNaCa2(SiO~)a- _ ----------- 10.4 

40 Prehnite no. 3 ______ --··----------------- H2Ca2AbSbOt2-------------- 10.2 

41 Prehnite no. 2 ______ Michigan __ __ __ __ ______ do _______________________ 10.4 

42 Prehnite no. L _____ ------------- ------- _____ do _______________________ 10.5 

43 Apophyllite no. L_ -------------------- H7KCat(Si0a)s.4HH20 _____ 10.4 

44 Apophyllite no. L. -------------------- _____ do _______________________ 10.4 

45 Wollastonite _______ -------------------- CaSiOa --------------------- 10.8 

46 Glass no. ,106 h ______ -------------------- Si02 75.48; NazO 15t26; CaO 11.2 

9.26. . 

47 Glass no. 135 b ______ -------------------- SiOz 78.40; Na20 16.17; CaO 11.4 

' 0.21. 

b Analyses by F. W. Qlaze, Bureau of Standards. 

ELECTROMETiUC DETERMINATION OF pH '--

KIND OF ELECTRODE USED 

By means of hydrogen-electrode measurements it 

was possible to study wore quantitatively the pH of 

many mineral solutions. Solution of the mineral was 

brought about by grinding in an agate mortar, as in 

the colorimetric test, with two drops of carbon-dioxidefree distilled water. Increasing the quantity of water 

and mineral indiscriminately was considered unwise, 

and with so small a volu:fne of solution the use of s01ne 

form of microelectroqe was necessary. 

The electrode desohbed by Bodine and Fink 27 was 

simplified and found to give good results. Pyrex glass 

tubing of about 3 millimeters inside diameter was 

drawn to a 1-millinieter capillary below ari opening in 

the tube for inserting the electrode. The electrode of 

platinum wire, B. & S. no. 25, sufficiently long to reach 

well into the capillary, was sealed into a Pyrex glass 

tube to be filled with mercury for making electrical 

. 21 Bodine, J. H., and Fink, p. E., A simple microvessel with electrodes for determining the hydrogen-ion concentration of small amounts of fluid,: Jour. Gen. Physi· 

ology, vol. 7, p. 735, 1925. · · 
6 SHORTER CONTRIBUTION TO GENERAL GEOLOGY, 1934-35 

eontact. It was flattened out at the point to increase 

the surface and thinly platinized by electrolysis in 

platinic-chloride solution. The capillary was then 

filled with water to keep the electrode n10ist, and the . 

electrode was inserted and sealed in with Kotinsky 

cement. The seal was then coated with collodion to 

insure its being airtight. This electrode is shown in 

figure 1, A. 

SOURCE AND PURIFICATION OF THE HYDROGEN 

Traces of impurities in the hydrogen, especially 

oxygen, cause gross errors in potentiometric pH determinations of alkaline solutions. Fo~ this reason 

the hydrogen wns generated electrolytically in caustic 

type K potentiometer, in conjunction with a Weston 

standard cell (unsaturated), two 2-volt storage batteries in parallel, and a needle-type galvanometer. In 

the preliminary work a normal calomel electrode was 

used with a saturated potassium-chloride bridge, but 

it was later replaced by the tenth normal electrode 

recommended by Clark. 

At the beginning of a series of measurmnents the 

apparatus was first thoroughly swept out with a 

stream of hydrogen from the generator a (fig. 1, B), 

the capillary point of the electrode being kept beneath 

the surface in a beaker of distilled water. The electrode was then removed from the water and placed in 

the solation und~r measurement, and the flow of 

Mercury 

column 

A ,, -

:, To 

+ / "' .I:~ pote[1tio.meter· CirCUit 

-t I+ I 

~ 

c 

~ ·~ 

a 

B FIGURE 1.--A, Microves:>el and electrode. B, Assembly of apparatus for pH determination: a, Hydrogen generator; b, c, stopcocks; d, electric furnace at 500° C. with -

copper gauze; e, hydrogen electrode, calomel half cell, and saturated KCI bridge. 

soda solution with 110-volt direct current and passed 

over copper gauze held at 500° C. to remove traces of 

oxygen. The flow of hydrogen was controlled with a 

24.5-ohnl rheostat. Between measurements the capillary point was kept below the surface of some distilled water, thus allowing a· steady stream of hydrogen 

to bubble through under a slight pressure. 

ASSEMBLY AND MANIPULATION OF APPARATUS 

In figure 1, B, is shown diagrammatically the assembly of apparatus. The measurements of electromotive force were n1ade with a Leeds & Northrup 

hydrogen was continued for a few· seconds. By 

switching off the hydrogen generator and opening 

stopcock b for a second, the hydrogen was reduced 

to atmospheric pressure, and the solution was drawn 

into the electrode vessel by capillarity. The capillary 

point was then placed just below the surface in the 

saturated potassium-chloride bridge, and readings 

were begun. 

The apparatus was frequently checked with an 

alkaline buffer solution ~ borax, described by Kolt-
ALKALINITY OF SOME SILICATE MINERALS 7 

hoff. 28 These check readings on buffer solutions were 

consistently a trifle low by 0.02 to 0.05 of a pH number, and rarely, as in the case pf a poorly activated 

electrode, they were low by as much as 0.10 of a pH 

number. Checks upon an acid buffer (phosphate, 

pH=6.81) gave results equally low, and the presence 

of oxygen as an impurity in the hydrogen did not 

appear to be the source of this small error. In most 

M 

of the determinations as many checks on - borax 

20 

were made as measurements on the mineral solutions, 

with a view to adding a small correction factor to the 

results. Further study, however, showed that the 

mineral solutions came into equilibrium with the electrode much more quickly than the buffer solution, and 

slightly low buffer readings were attributed to the 

slowness of the electrode in reaching equilibrium. 

In the measuring of all solutions readings were 

taken at intervals of 1 minute until four check readings 

were obtained. Initial readings were consistently low, 

then showed an upward creep in electromotive force 

until a constant value was obtained. Owing to the 

low conductivity of the mineral solutions no attempt 

was made to make readings closer than 0.5 millivolt 

or in some readings 1 millivolt. 

EFFECT OF CARBON DIOXIDE 

That atmospheric carbon dioxide affected to some 

extent, though perhaps slightly, the pH of the mineral 

solutions prepared for the colorimetric determinations 

was . obvious, and elimination of ear bon dioxide was 

attempted by preparing the mineral solutions in carbon-dioxide-free nitrogen. For this purpose a thin 

rubber membrane was fitted tightly over the top of 

the mortar by means of passepartout, with small openings for inserting the pestle and the nitrogen inlet. 

The mineral particles were placed in the mortar, 

after whieh nitrogen was allowed to flow slowly 

through, sweeping out all air. Three drops of carbondioxide-free distilled water was then added, followed 

by insertion of the pestle in the opening. The rubber 

·membrane became inflated by the nitrogen, the overflow leaking out around the sides. The flow of nitrogen 

was continued during the grinding and solution of the 

mineral and until the solution was drawn into the 

electrode vessel. 

To be sure of the absen~e of carbon dioxide the nitrogen was tested by bubbling it through baryta solution. 

After 5 minutes no precipitation was noticed. 

To show the effect of carbon dioxide if no special 

precautions are taken to exclude it, one of the more 

alkaline materials, glass no. 135, was tested in air and 

in nitrogen, with the results given in table 3. 

28 Koltboff,' l. M., pH and electrotitrations, p. 100, New York, John Wiley & Sons, 

1931. For temperature corrections seep. 40. 

TABLE 3.-Effect of carbon dioxide on pH 

[Material, glass no. 135, ground 2 minutes; normal calomel electrode used] 

Ground in air Ground in carbon-diox~e-free nitrogen 

Temperature E.m'.f. pH Temperature E.m.f. pH (oC.) (oC.) 

25 ___________ 0. 9468 11. 21 27 _________ 0. 9660 . 11. 47 

25 ___________ . 9555 11. 36 27 _________ . 9675 11. 49 

26 ___________ . 9491 11. 21 27 _________ . 9650 11. 46 

28 ___________ . 9340 10. 90 28 _________ . 9680 11. 47 

28 ___________ . 9490 11. 15 28 _________ . 9710 11. 51 --· Average ____________ 11. 17 Average ________ ___ 11. 48 

Maximum deviation __ . 27 Maximum deviation . 03 

The results in the presence of atmospheric carbon 

dioxide were evidently low by 0.31 pH unit and showed 

poor agreement. When the effect of carbon dioxide 

was eliminated, however, the results showed remarkable precision. In all subsequent determinations, 

therefore, the action of carbon dioxide was eliminated. 

Its effect is discussed in the comparison of electrometric with colorimetric determinations (p. 9). 

TIME OF GRINDING 

In preparing solutions the time of grinding was, of 

course, limited by evaporation of water in the carbondioxide-free nitrogen, but the period could be extended 

to 3 minutes without the suspension becoming so 

thick as to make measurements impossible. The 

comparison of results obtained by grinding for 1, 2, 

and 3 minutes with glass no. 135 and with spodumene. 

is given in table 4. 

TABLE 4.-Time of grinding needed to reach equilibrium 

[Normal calomel electrode used] 

Glass No. 135 

1 minute 2 minutes 3 minutes 

Temper- Temper- Temper- ature E.m.f. pH I ature E .m.f. pH ature E .m.f. pH 

(oC.) (00.) (00.) 

-------·· --- - ---- 30 _____ 0. 9680 11. 40 27 ____ 0. 9660 11. 47 30 ___ 0. 9745 11. 50 

30 _____ . 9685 11. 40 27 ____ . 9675 11. 49 30 ___ . 9760 11. 53 -- --- -- Average _____ ll. 40 Average ____ II. 48 Average ____ 11. 51 

Spodumene 

30 _____ [o. 828019. o8 30 __ __ 0. 8410 9. 28 31_ __ 0. 8470 9. 36 

30 _____ 1. 83301_!1:_1_5 30 ____ . 8450 9. 35 31_ __ . 8420 9. 27 

30 ____ . 8430 9. 32 31_ __ . 8430 9. 29 -- -- Average _____ 9. 11 Average____ 9. 32 Average ___ 9. 31 

It appears from these result~ th.at equilibrium is 

almost established by grinding for 1 minute, the results 

being only 0.08 to 0.20 of a pH unit lower than after 
8 SHORTER CONTRIBUTION TO GENERAL GEOLOGY, 1934-35 

grinding for 2 or 3 minutes. In view of the results of 

grinding for 2 or 3 minutes, the establishment of 

equilibrin~ in 2 minutes seems a logical conclusion, 

the agree~ent being remarkably good. With glass 

no. 135 a small difference, 0.03 of a .pH number, was 

found between the results of grinding for 2 minutes 

and 3 minutes. It must be remembe-red, however, 

that this d~fference is well within the limits of accuracy 

of the method and that it may have been caused by 

the shift of temperature from 27° to 30° C. The test 

for spodun1ene, a material less easily attacked by 

\Vater, was h1ade at nearly constant ten1perature and 

showed a better agreement between the results for 

2 nlinutes and for 3 minutes. 

EFFECT OF AGATE MORTAR 

The action of the silica of the agate mortar in neutralizing the alkalinity of the mineral solutions remained a n1atter of doubt. There are several reasons 

for thinking that any such action would be slight. 

Presumably the silica would first have to becon1e 

hydrated to silicic acid before its acid properties could 

becon1e manifest, and this reaction occurs but slowly 

at ordinary temperatures. Moreover, the silication of 

an alkali reduces the pH number only slightly. The 

hardness of the agate a.nd the small surface exposed to 

\Vater in comparison to that of the nlineral were other 

reasons for thinking the action negligible. Under the 

conditions of experin1entation the strongneutralizing 

action of the n1inerals was found to develop quickly, 

as noted for orthoclase in the colorimetric work. 

(See p. 4.) 

The silica of the agate mortar, if effective in neutralizing mineral alkalinity, should have an action on 

a buffered solution, neutralizing alkali and reducing 

the pH. The use of an unbuffered alkaline solution 

in this test was thought undesirable because of possible 

changes in concentration due to evaporation. A 

buffer solution was accordingly used, as such solutions 

are but little affected by dilution. 

Measurements of~ borax ground in an agate mortar 

for 2 nlinutes in carbon-dioxide-free nitrogen gave a 

reading within 0.02 of a pH number of the correct 

value, while the untreated buffer gave a reading 

within 0.03. Later, the effect of grinding the buffer 

with powdered quartz was determined. The electrode 

in this test was a little less active. The untreated 

buffer gave results 0.05 of a pH number lower than 

the correct value; the buffer ground for 2 minutes 

with quartz powder gave results 0.07 of a pH number 

lower than the correct value. 

Indicator solutions ground alone in the agate mortar 

showed no change of color. 

These results, although not conclusive enough to 

show that the agate is completely ineffective, indicate 

that the resulting error, if present, is very small. 

EFFECT OF IRON IN MINERAL 

Two factors limited the nun1ber of n1inerals that 

could be used for electrometric pH measurements. 

Minerals of low alkalinity gave solutions so low in 

conductivity that measurements of electromotive force 

could not be made with any certainty; therefore, for 

the most part, only the more alkaline minerals were 

tested. The presence of oxidiz;able materials-chromium, manganese, and most particularly iron-causes 

gross errors. Minerals containing more than a trace 

of iron were found unsuitable for test by this method. 

Substances present in an oxidized condition not only 

give an oxidation-reduction potential but reduce the _ 

pressure of hydrogen at the electrode surface and thus 

decrease the hydrogen-electrode potential. 29 -This 

effect seems particularly pronounced in the use of the 

microelectrode. 

The action of iron as an in1purity upon the measurenlents on preh:riite is shown in table 5. For prehnite 

no. 1 and no. 2, the results are an average of three 

tests; for prehnite no. 3, an average of two tests. 

TABLE 5.-Effect of iron on electrometric measurements of pH 

ColoriI Electrometric determination 

Mineral metric Percent Maximum determina- of Fe Terntion of pH perature pH deviation 

from (00.) average 

Prehnite no. 3 _____ _ 10. 5 0. 37 28 10. 38 0. 04 

Prehnite no. ------~ 10. 4 3. 661 27 10. 06 . 01 

Prehnite no. L _____ 10. 2 I 5. 1 26 8. 5 . 8 

EFFECT OF TEMPERATURE 

No attempt was made to study the effect of temperature, but in each test the temperature of the electrode 

was measured with a thermometer dipping beneath 

the surface of the potassium-chloride bridge. -The 

ionization of water increases rapidly with increase in 

temperature, and that would have art effect on the 

hydrolysis of the mineral and the resulting pH. However, there seems no simple method of controlling the 

temperature in the rather complex operations of grinding and of transferring the solution to the electrode. 

In many of the determinations it was noted that small 

fluctuations in the temperature affected the pH readings but little. This is especially apparent in the tests 

-on glass no. 135. During· the seven determinations 

the temperature ranged from 27° to 30° C., but the 

maximum deviation from the averagfj was only 0.05 

of a pH unit. 

RESULTS OF ELECTROMETRIC DET~RMINATIONS 

The results of electrometric measurements on a 

number of minerals are listed in t~ble 6, with the 

colorimetric results from table 2 for ~~mparison. For 

29 See Clark, W. M., op. cit. p. 371. 
/ AL:KALINITY OF SOME SILICATE MINERALS 9 

minerals giving solutions with pH near 9.0, the two 

methods of determination gave results in close agree~ 

tnent. For the more alkaline minerals, especially those 

containing calcium and magnesium, the effect of atmospheric carbon dioxide in the colorimetric method 

becomes more apparent, and the results of the two 

methods differ by as much as 0.4 of a pH unit. Solutions of calcium carbonate were found to be especially sensitive to atmospheric carbon dioxide, and as 

the color of the indicator slowly faded, the colorimetric 

tests were unsatisfactory in consequence. Atkins 30 

determined colorimetrically the pH of calcite solutions 

made by boiling and cooling to be a little above 9.0. 

This shows close agreements with the electrometric 

result listed in table 6, pH=9.03. 

Especially worthy of note is the high pH value ol;>-

tained for wollastonite (pH=ll.17). This shows a 

good comparison to the electrometric titration curve 

of silicic acid with calcium hydroxide as determined 

by Britton. 31 

TABLE 6.-Results of electrometr·ic tests 

Electrometric determination 

Colorimetric Maxi- Index Mineral deter· Tern- Num- mum no. min a- ber of 

tion of pera- deter- pH de via- ture tion pH (0 C.) mina- from tions average 

------

Calcite __ ___ _____ a 9. 0 25 2 9. 03 0. 02 

14 Pollucite ___ __ ___ _ 9. 0 26 2 8. 96 . 02 

17 Orthoclase no. L _ 9. 2 28 3 9. 18 . 03 

19 Spodumene ______ 9. 2 30- 31 6 9. 31 . 05 

22 Albite no. 2 ______ 9. 8 28 2 9. 84 . 04 

27 N atrolite ____ __ __ 10. 0 28 2 10. 05 . 03 

37 Tremolite no. L __ io. 2 22 3 10. 50 . 02 

38 Tremolite no. 2 ___ 10. 2 24 3 b 10. 17 . 06 

44 Apophyllite no. 2_ 18 ~ I ~~ : 

3 10. 79 . 02 

45 . Wollastonite ____ -~ 3 11. 17 . 01 

47 Glass no. 135 _____ 11. 4 27- 30 7 11. 49 . 05 

a Determined by Atkins. b Probably too low. Contained a trace of iron and 1.15 percent of manganese. 

COMPARISON OF COLORIMETRIC AND ELECTROMETRIC METHODS AND FACTORS INVOLVED IN 

TESTING 

In a general way the two methods of detennining 

the hydrogen-ion concentration of mineral solutions 

described in the: preceding pages yield similar results, 

but each method has certain limitations. 

In the colorirnetric test as actually used there are 

two sources of error- the effect of atmospheric carbon dioxide and the great turbidity of the solutions. 

However, as shown by comparison with electrometric 

results obtained in the absence of carbon dioxide, 

the error due to its presence is small where the pH 

of the solution is under 9.0, although it becomes 

larg~r with increasing alkalinity. This error would 

likewise not be the same for different n1inerals but 

3° Atkins, W. R. G., Some geochemical applic:1tlons of measurements of hydrogenion concentration: Roy. Dublin Soc. Sci. Proc., vol. 19, pp. 455-460, 1930. 

31 Britton, H. T. S., Electrometric study of the precipitation of silicate.s: Chern. 

Soc. London Jour., 1927, p. 425. 

would be more pronounced with n1inerals containing 

calcium and magnesium because of the insolubility 

of the carbonates formed. Elimination of carbon 

dioxide and compensation for turbidity n1ight yield 

results of greater precision, such as those determined 

electrometrically and given in table 6. However, 

the results obtained colorimetrically are roughly 

comparable when sn1all differences in pH are not 

considered, and they are capable of yielding many 

useful generalizations. 

It n1ust. be remembered that n1inerals are seldom 

pure, and in solutions of the less alkaline minerals 

the pH may be the result of alkaline impurities such 

as calcite. This is shown most clearly with calamine, the composition of which is given by the formula H2 (Zn02)Si04 • Britton 32 has determined the 

pH precipitation curves of various metall~c hydroxides. 

He found that the precipitation of zinc from zinc 

sulphate solution began at pH= 5.2. Son1e uncertainty exists as to whether this precipitate was the 

hydroxide or a basic sulphate. Nevertheless, in 

view of these results, it does not seem possible to 

obtain a pH value as high as that observed for calamine, pH~ 8.8, without the presence of an alkaline 

impurity. 

In the electrometric tests the chief source of difficulty has been the presence of iron and other oxidizable material in the minerals. Errors due to the 

presence of iron seem to be particularly pronounced 

with the microelectrode, soine fonn of which was 

required for the tests. It is true that in numerous 

reports in the literature fairly accurate determinations · of pH were obtained in the presence of ferric 

iron by use of the Hillebrand electrode or the Ciark 

hydrogen-electrode vessel, but in both of these designs 

the hydrogen at the platinum electrode surface is 

constantly renewed. With the . microelectrode no 

means is provided for keeping the hydrogen on the 

electrode surface at atmospheric pre3sure, and con:-

sequently good results could be obtained only when 

iron was absent, or nearly so, · 

There is need for an electrode that is not affected 

by oxidizable materials. Such an electrode is the 

glass electrode, described by Macinnes and Dole, 33 

but its use in this study would be difficult, owing to 

the low conductivity of the mineral solutions. Furthermore, tests ~ith the electrode for solutions with 

pH above 9.0 show a large error, which varies in the 

presence of different alkali and alkaline earth hydroxides, and, as previously shown, figures as high 

as pH=11.17 (wollastonite) are obtained for these 

mineral solutions. The limitations of the quinhydrone electrode and of the oxygen electrode ,· in 

alkaline solutions- are even greater than those of the 

hydrogen electrode. 

az Britton, H. T. S., Electrometric studies of the precipitation of hydroxide3, 

pt. 2: Chern. Soc. London Jour., vol. 127, pp. 2120- 2141, 1925. 

33 Macinnes, D. A., and Dole, Malcolm, The behavior of glass electrodes of 

different compositions: Am. Chern. Soc. Jour., vol. 52, p. 29, 1930. 
10 SHORTER CONTRIBUTION TO GENERAL GEOLOGY, 1934-35 

It is noteworthy that the results listed in table 6, 

determined with the hydrogen microelectrode, were 

obtained with an excellent degree of reproducibility, 

the maximum deviation from the average being 0.01 

to 0.05 pH unit. Three determinations on wollastonite 

showed a maximum deviation of only 0.01 pH unit 

from the average. Six tests on spodumene yielded a 

maximum deviation of 0:-05 pH unit from the average, 

and seven determinations on glass no. 135 showed a 

maximum deviation of 0.05 pH unit from the average. 

This reproducibility, together with the study of the 

grinding period needed to reach equilibrium (see table 

4), shows that the true equilibrium is approached 

sufficiently close to give comparable results for the 

extent of decomposition possible with different mineral 

silicates. 

The effectiveness of high temperature and pressure 

in increasing solubility of minerals has been noted by 

Bouyoucos,34 and, presumably the resulting pH 

would thereby also be increased. The ionization of 

water increases with temperature, as does also the pH 

of many buffer solutions. No control of these factors 

bas been attempted, but the conditions of the experiments were such as to give c01nparable results. For 

the electrometric 1neasurements the approximate temperature of the electrode was noted. 

GENERAL APPLICATIONS 

It is well known that the measurement of hydrogenion concentrations has found numerous applications 

in biology and industrial chemistry, but the subject 

does not seem to have received as much consideration 

in geochemical studies as it deserves. The present 

study has shown that many silicates give intensely 

alkaline solutions when the mineral is sufficiently exposed to the action of water. Solutions percolating 

through the earth's crust come into contact with immense quantities of these silicate materials, which may 

neutralize acidity and maintain a definite pH. 

Numerous investigators have shown that the maintenance of a definite and limited range of pH is a 

condition necessary for certain chemical reactions and 

for the deposition of many minerals and ores. Wells 35 

has noted such lirnitations in the precipitation of metallic copper by ferrous chloride. Furthermore, the 

deposition of sulphides and carbonates is affected to a 

great degree by changes in pH. 36 Many chemical reactions are catalytic in their nature; in such reactions 

the pH of the soluti~n may determine the speed of the 

reaction, although it does not, in general, influence 

the final equilibrium. Related to catalysis are adz• Bouyoucos, G. J., Rate and extent of solubility of minerals and rocks under 

different treatments and conditions: Michigan Agr. Exper. Sta. Tech. Bull. 50, 1921. . 

as Wells, R . C., Chemistry of deposition of native copper from ascending solutions: 

U .8. Geol. Survey Bull. 778, p. 65, 1925. 

36 Stieglitz, Julius, The elements of qualitative chemical analysis, pt. 1, pp. 1()()-

202, New York, Century Co., 1911. 

sorption and the precipitation of colloids, phenomena 

conditioned by the pH of the solution. pH frequently 

determines the crystal form and the water of hydration 

of material separating from solution. A great multitude of studies deal with the influence of pH on inorganic equilibria. The 1nany applications of pH are 

discussed by Clark. 37 

However, the neutralizing power of silicates .is subject to certain limitations, which must be remembered 

in applying the results under geologic conditions .. The 

writer has already mentioned the protective action of 

colloidal complexes of silica and alumina formed in 

the decomposition of silicates and the need of exposing 

a large . surface in order to reach an equilibrium between the mineral and the solution. In any case, an 

equilibrium is established between the surface exposed 

and the solution, and if complexes containing a 

smaller amount of base are formed on the mineral 

surface the equilibrium will be between these complexes and the solution. Only by removal of this 

colloid or the exposure of fresh material can the true 

equilibrium of mineral to solution be established. 

The principles of buffer action need also to be 

borne in mind, with the conception of pH as an intensity factor in acidity or alkalinity and not as a 

quantitative measure of the total acid or alkali in the 

solution. A brief discussion of .buffer action may 

serve to emphasize the limited neutralizing action of 

silicate solutions from this viewpoint and to contrast 

the intensity factor pH with the quantitative relations 

involved in neutralization. 

Buffer solutions are solutions of weak bases and 

their salts or week acids and their salts-for example, 

a solution of acetic acid and sodium acetate. The 

addition of strong hydrochloric acid, HOI, to this 

solution would form sodium chloride and acetic acid. 

The acetic acid thus produced is weakly ionized, and 

for that reason the pH of the solution is altered but 

slightly by small additions of the strong acid. Similarly, the addition of small quantities of the strong 

base sodium hydroxide, N aOH, results in the destruction of acetic acid, which is so weakly ionized that by 

its removal from solution the pH is but slightly 

changed. In contrast to this, the titration of a strong 

alkali with a strong acid results in the destruction 

· of the strong alkali, and the pH is shifted greatly 

by small additions of the acid. 

The silicate minerals under consideration are salts 

of strong bases and an exceedingly weak acid, and if 

they were present in solutions in large quantities the 

solutions would have a strong buffer action. But the 

solubility of these minerals is very slight, their high 

alkalinity being the result of nearly complete hydrolysis into a strong base and a very weak acid (silicic 

acid). Because the quantity of mineral dissolved is so 

a7 Clark, W. M., op. cit., p. 549. 
ALKALINITY OF SOME SILICATE MINERALS . 11 

small the neutralizing power of the solutioJ?- is limited, 

though this neutralizing power may be replenished by 

the exposure of fresh surfaces to the solution. 

Neutralization with carbonates results only in the 

evolution of a gas, carbon dioxide, and in the formation of soluble salts, products which under the conditions existing at the surface of the earth are quickly 

removed and in no way retard the neutralizing action. 

In contrast to this, silicate decomposition forms colloidal films of silica and alumina and their complexes 

on the mineral surface, and this retards the neutralizing activity. Therefore it follows logically that under 

surface conditions the principal neutralizing agents 

for acid ground waters are the carbonate minerals. 

But in deep-seated waters, which are under pressure, there is no nieans for the removal of the gaseous 

product, carbon dioxide, and the carbonates become 

less potent as sources of alkali. It would seem that the 

great mass of silicate minerals might then cause solutions previously neutralized by carbonates to become 

extremely alkaline. 

The effect of neutral salts on the alkalinity of silicates 

is another factor that needs brief consideration. These 

salts are always present in underground waters, being 

dissolved in the downward flow or formed by the 

neutralization of ground-water acidity. The effect of 

salts on the pH of solutions and in increasing the 

solubility of other salts has been shown in many experiments; their influence on the pH of mineral solutions needs study. 

Consider the hydrolysis of calcium metasilicate in its 

simplest form, the products being calcium hydroxide 

and silicic acid with the ions and decomposition products formed from them. This is expressed by the 

equation given below, in which the relative size of the 

arrows indicates the relative speed of the reaction in 

each direction and the numbers at the arrows designate 

the separate reactions for later reference. 

(1) 

CaSi03 + 2HOH :;-- Ca (OH)2 

tl (2) 

CaOH++OHT"' (3) 

Ca+++OH-

+ 

H++HSiOat l (4) 

H2SiOa 

tl (5) 

H20+Si02 

According to the law of mass action the extent of 

reaction 1 is expressed by K 1: 

K products 1 reactants 

[Ca(OH)2] [H2SiOa] 

[CaSiOa] [HOHh 

This expression may be simplified, as the reactants are 

present in excess, so that their concentration in solution 

is constant. For the same reason the concentration of 

silicic acid may be neglected, the solution being saturated with it at all times owing to its removal as colloidal silica according to reaction 5. The extent of renction 1 is then given by the constant 81 = [Ca(OH)z]. 

Thus under conditions where an excess of CaSi03 i~ 

present, no matter what the concentration or nature of 

the · dissolved salts, the concentration of Ca(OH)2 

in the solution is constant. 

Reactions 2 and 3 are expressed by the constants K2 

and K3, respectively: 

K _ [CaOH+] [OH-] 2- [Ca(OH)2] 

·[Ca ++] [OH -] 

Ka= - [CaOH+] 

Reaction 4 may be neglected, as it occurs to so small 

an extent. 

The addition of an acid to the solution results in the 

formation of water and of the calcium salt of the acid 

and in the temporary removal of OH ions. However, 

the removal of OH ions causes further reaction, and 

equilibrium is again established- in accordance with 

81, K2, and K3. Thus the OH-ion concentration would 

remain essentially constant after small additions of 

acid, when an excess of CaSi03 is present. The 

addition of a salt with acid properties but containing no 

ion in common with CaSi03 would have a like effect. 

But the presence of calciun1 salt in the solution, either 

formed by neutralization or present by direct addition 

to the solution, would result in a large concentration 

of the ion Ca++, and reactions 2 and 3 might be thus 

repressed in accordance with the constancy of K2 and 

K3 • In this discussion c~lcium metasilicate has been 

used as an example, but the relations apply equally 

well to other silicates. 

However, calculations show that a slight change in 

ionization of a strong base would affect the pH but 

slightly. For example, a solution of sodium hydroxide 

at pH= 10 would be at p0H=4 and [OH-] would be 

equal to 0.0001. If half of the hydroxyl ions were 

removed from solution by a mass-action effect, [OH-] 

1 

would be 0.00005. As pOH=log [OH-j' pOH would 

then be about 4.3 and pH about 9. 7. Thus by removing half of the hydroxyl ions the pH of the solution is 

decreased by only 0.3 of a pH unit. It would seem, 

therefore, that the pH resulting from the hydrolysis of 

silicate minerals would be decreased but little by any 

salts that the solutions may contain. 

As a general principle the presence of additional 

salts would tend to increase the solubility of the relatively insoluble minerals and might cause slight changes 

in the pH figure, the direction and extent of which 

would depend on the nature of the salts. 

The neutralizing activity of silicates and the effect 

of salt solutions on their alkalinity were not studied 

further in the present investigation. Such a study 

would obviously be preceded by experiments on the 
12 SHORTER CONTRIBUTION TO GENERAL . GEOLOGY, 1934- 35 

action of pure water, such as are described in this 

paper. 

The results presented in table 2 show decided differences between minerals, but too much weight should 

not be given to small differences in pH, on account of 

the effect of possible impurities in some of the specimens, incipient alteration, possible varieties in the 

species, and similar variations inherent in natural 

products. Such variations with different samples of 

the same mineral were found for muscovite orthoclase ' ' albite, labradorite, tremolite, prehnite, and apophyllite. 

As the basic constituents are progressively leached 

from these materials in weathering, their reaction to 

water would become less alkaline. In this connection 

several clays and decomposed rocks were tested colorimetrically and were found to give acid or nearly 

neutral reactions to indicators- for example, a decomposed granite collected near the surface in the 

District of Columbia gave approximately pH= 3. It 

cannot be definitely stated whether this acidity is due 

to organic acids or to aluminosilicic acids formed by 

decomposition. Soil chemists are not in close ao-ree- o 

ment as to which of these materials causes the acidity 

of many soils. 38 

In any event, alteration seems a logical explanation 

for the small difference in pH found for different 

samples of the same mineral and suggests the need 

for further study of rocks and minerals regarding this 

point. 

In a broad way, the order of the minerals in table 2 

shows their potential alkalinity. A general increase 

in pH with increase in alkaline content may be observed, and this suggests a correlation between the 

pH values of silicates and their composition. The 

content of silica, alumina, and the different bases 

should have some definite effect on the resulting pH, 

and with improved methods of testing it might be 

useful as a rough check on analyses. However, the 

chemical structure is not known for all silicate minerals, and, as illustrated by organic compounds, properties may be determined more by chemical structure 

than by percentage composition. But the relation 

between composition and alkalinity is definitely indicated in table 1 and is evidence that in chemical 

structure the mineral silicates are relatively simple in 

contrast to the complexity of organic compounds. 

A brief consideration of the measurernents in connection with the law of mass action results in a further 

understanding of weathering. Consider the solution 

and hydrolysis of sodium silicate and of calcium silicate together in the same solution, thus: 

Na2Si03 + 2HOH._=:2Na++ 20H- + H 2Si03 ______ (1) 

CaSi03+ 2HOH~Ca++ + 20H- + H2Si03 ______ (2) 

38 On the pH of soils see Wherry, E. T., Soil acidity, its nature,. measurement, and 

relation to plant distribution: Smithsonian Inst. Ann. Rept. for 1920, p. 247, 1922. 

For simpl~city the sodium and calcium hydroxides 

formed are assumed to be completely ionized. 

According to the law of mass action the extent to 

which these reactions may occur is determined by the 

relative concentrations of the products and reactants 

and is expressed by the equilibriu-m constant K. 

Thus in reaction 1 

K _products 

1 - reactants 

and in reaction 2 

[N a+]2 [OH-]2 [H2SiOaJ 

[N a2Si03] [HOH] 

K =products= [Ca++] [OH-]2 [H2Si03] 2 reactants [CaSi03] [HOH] 

The concentrations of the reactants and of the 

silicic acid may be considered constant, as they are in 

excess, and the extent of the reaction is given more 

simply by the product S, where 

81 = [N a+] 2 [OH-] 2 

82 = [Ca++] [OH-F 

On account of solubility and possibly for other 

reasons reaction 1 occurs to a greater extent than 

reaction 2, so that sl is greater than s2. When 

sodium silicate and calcium silicate are acted upon by 

water together the sodium silicate hydrolyzes to an 

extent expressed by Sr, giving a relatively large concentration of OH ions to the solution. The presence 

of a large concentration of OH ions obtained bv 

reaction 1 retards reaction 2 to a degree commensurat~e 

with the figure for 81• In this way the presence of 

sodium silicate tends to protect the calcium silicate 

from the action of water. Thus when water acts on 

mineral silicates together in a rock the alkaline coU:-

stituents may serve as a protection to those that are 

less alkaline. 

Measurements of the pH of mineral solutions 

indicate the quantity of alkali removed and thus serve 

as an index of the decomposition taking place. It also 

follows from the consequences of mass action shown 

above that the order of the minerals as listed in table 2 

gives an approximate idea of their relative stability. 

Thus we find minerals such as beryl, lepidolite, and 

muscovite giving solutions with relatively low pH; 

it is generally considered that they are less readily 

attacked in weathering processes. Biotite, laumontite, 

orthoclase, and spodumene are slightly more attacked, 

and they are found to give higher pH. The plagioclase feldspars albite and labradorite are next in the 

list, along with such minerals as leucite, margarite, 

natrolite, and others. The alkalinity and case of decomposition of the more basic rock constituents such 

as prehnite, pyroxene, tremolite, apophyllite, and 

especially wollastonite are shown by their high pH. 
ALKALINITY OF SOME SILICATE MINERALS 13 

The ·results are in fair agreement with field observations on the decomposition of rocks and stability of 

n1inerals, and it is hoped that further field observations 

will serve to supplement the conclusions that have been 

tentatively drawn from the laboratory observations. 

In addition to geochemical applications, it may be 

possible that the method of testing above described 

could be used practically in determining the weathering qualities of ceramic materials and glass. Tests 

of two glasses are shown in table 2, and the more 

accurate electrometric tests of glass no. 135 are given 

in table 6. Several glasses containing a larger percentage of alkali gave pJI numbers above 12. The 

colorimetric test on some Pyrex glass was about 

pH=8. This Pyrex glass and the more alkaline glasses 

tested represent extremes of durability and show a 

~ood contrast by the method of test. 

0 

The effectiveness of silicate minerals as agents in 

det'ermining the pH of aqueous solutions in geochemical processes seems to be clearly jndica te~ by the 

experiments described in this paper, although it should 

be again pointed out that the abundance of carbonate 

rocks near the surface leads to their being the source 

of alkalinity in Inany natural waters. In d~ep-seated 

waters the silicates would have a major infhience. 

In brief, the results of these tests are a rough index 

of the weathering qualities of different silicate minerals 

and show how the silicates in nature may affect the 

pH of solutions.with which they are in contact. The 

pH of the solutions thus produced by the hydrolysis 

of these minerals may, in turn, have an influence on the 

chemical reactions taking place and in determining the 

materials deposited from the solutions.