DEPARTMENT OF THE INTERIOR

UNITED STATES GEOLOGICAL SURVEY

GEORGE OTIS SMITH, DIRECTOR

BULLETIN 609

THE FRACTIONAL PRECIPITATION OF

SOME ORE-FORMING COMPOUNDS AT

MODERATE TEMPERATURES

BY

ROGER C. WELLS

WASHINGTON

GOVERNMENT PRINTING OFFICE

1915

CONTENTS.

Page.

Introduction............................................................. 5

Scope of the investigation.............................................. 5

The mass law........................................................... 8

Ions and the solubility product........................................ 8

Fractional precipitation................................................... 10

Previous investigations............................................... 10

Method of experimentation............................................ 11

Sulphides................................................................ 12

Solubility of sulphides...................................'............. 12

Experiments with the nitrates of copper and lead...................... 15

Experiments with the sulphates of copper and zinc...................... 15

' Work of Anthon and Schiirmann....................................... 16

Compound sulphides.................................................. 18

Reducing action of sulphides.......................................... 20

Electric activity in sulphides.......................................... 22

Summary of results of experiments with sulphides...................... 23

Hydroxides.............................................................. 24

Previous work ........................................................ 24

Solubility of hydroxides................................................ 24

Fractional precipitation of hydroxides....:............................ 25

Ferrous and cupric hydroxides ........................................ 26

Summary of results of experiments with hydroxides..................... 28

Carbonates .............................................................. 28

Previous work........................................................ 28

Solubility of carbonates...:........................................... 29

Effect of hydrolysis................................................... 30

Experiments with carbonates at ordinary temperature................... 31

Results of experiments with bicarbonates.............................. 37

Experiments with carbonates at higher temperatures................... 40

Summary of results of experiments with carbonates..................... 42

Silicates................................................................. 43

Previous work........................................................ 43

Experiments with silicates at ordinary temperature..................... 43

Summary of results of experiments with silicates....................... 44

General view of the results of the investigation............................. 45

ILLUSTRATIONS.

Page.

FIGURE 1. Select!ve precipitation of copper before zinc by ammonium sulphide. 15 

2. Apparent selective precipitation with ferrous and cupric sulphates. 26

3

THE FRACTIONAL PRECIPITATION OF SOME OREFORMING COMPOUNDS AT MODERATE 

TEMPERATURES.

By ROGER C. WELLS.

INTRODUCTION.

SCOPE OF THE INVESTIGATION.

The experiments described in this bulletin were made to aid in 

elucidating the chemistry of ore deposition. They were confined to 

aqueous solutions at moderate temperature, and they have shown 

the order of solubility of the compounds of each of the classes investigated sulphides, hydroxides, carbonates, and silicates. The solubility of most of these compounds in water is so slight that its direct 

and accurate determination is difficult. The literature of the subject 

accordingly contains many determinations of their solubility by indirect physicochemical methods, and some of these are given in this 

bulletin in order that they may be compared with the results obtained in the experiments here reported. Three or four complications 

have been met in the otherwise simple relations of solubility, namely, 

that sulphides may act as reducing agents, that carbonates and silicates may precipitate hydroxides, and that solids may require considerable time to assume their most stable forms.

Considered chemically, the precipitation of an ore from a solution 

may require only a single condition the constituents of the ore must 

be present in the solution in certain forms and in concentrations above 

certain minima. A closer examination of this requirement, however, 

shows that it is conjoined with other minor requirements, all more 

or less related in such a way that a certain lack of one may exist if 

another is more than fulfilled. A statement of the relative solubility 

of two metallic compounds, based on the final table given in this 

paper (p. 45) need not, for example, correspond with the order of 

their deposition in nature, although the solubility of an ore has a 

direct relation to its deposition. The solubilities presented were determined by starting with solutions containing the competing metallic

5
6 FRACTIONAL PRECIPITATION" OF ORE-FORMING COMPOUNDS.

constituents in equal concentrations, whereas in nature the constituents occur in various concentrations and the deposition of ores is 

governed very largely by the mass law, a point that will be more 

fully discussed elsewhere in this bulletin.

The present experiments do not show the behavior of solutions that 

may exist in magmas or that may constitute magmatic exudations. 

Such solutions are not only at high temperature and pressure but 

some of them are also supposed to contain liquid carbon dioxide, 

chlorides, fluorides, and sulphides, together with water. As some of 

these compounds are mutually incompatible, the assumption that 

such solutions are a source of ore deposits requires a great deal of 

explanation, one of the first requirements being the ascertainment of 

groups of metallic and nonmetalhc elements that can be appreciably compatible in such solutions. Though such solutions seem to 

have played important roles in the deposition of certain ores, as indicated by geologic evidence, yet, as aqueous solutions at moderate 

temperatures form not only the solvent but the transporting medium 

and distributing agent of enormous quantities of metalliferous material, especially in the upper parts of the earth's crust, any generalizations that may be established with reference to such solutions will 

have wide application.

It is interesting to trace the cyclic courses of the various elements 

in the numerous operations that are going on unceasingly in the 

earth. The continents are elevated and eroded; the sediments 

consolidate and possibly fuse together anew to form new magmas; 

intrusions and .extrusions are constantly occurring: The following 

quotation from T. Sterry Hunt well expresses the continuity of these 

processes: 1

We learn * * * that amid all the changes of the face of the globe the economy 

of nature has remained the same. We are apt, in explaining the appearances of the 

earth's crust, to refer the formation of ore beds and veins to some distant and remote 

period, when conditions very unlike the present prevailed; when great convulsions 

took place and mysterious forces were at work. Yet the same chemical and physical. 

laws are now, as then, in operation; in one part dissolving the iron from the sediments 

and forming ore beds, in another separating the rarer metals from the ocean's waters; 

while in still other regions the consolidated and buried sediments are permeated by 

heated waters, to which they give up their metallic matters, to be subsequently 

deposited in veins. These forces are always in operation, rearranging the chaotic 

admixture of elements which results from the constant change and decay around us.

The courses of the valuable metals in these cycles, as well as those 

of the nonmetals most frequently associated with them, are of particular significance in respect to theories of ore deposition.

1 Hunt, T. S., Chemical and geological essays, chapter 12, pp. 235-236,1878.
SCOPE OF THE INVESTIGATION. 7

In discussions of the origin of ore deposits the nonmetallic constituents are frequently neglected. If an ore contains more than 

one element the term "source of the ore" is equivocal, as it is possible 

that the metal may have come from one source and the nonmetal 

from another. The presence of the nonmetallic constituent needed 

to precipitate the metallic constituent may in fact explain the presence or absence of an ore at a given locality. Moreover, the proportion of such a constituent may determine the character of the ore; 

for example, in a solution containing iron and sulphur an excess 

of sulphur may yield pyrite and an excess of iron may yield pyrrho tite. 

This essential factor must be borne in mind in considering the mode 

of formation of ores. In the treatment adopted in this bulletin 

the constituents discussed are primarily arranged according to the 

nonmetallic constituent that combines with the metal, the sulphides 

in one group, the carbonates in another, and so on; that is, all the

possible compounds of each metal are not considered together. 

Both arrangements have advantages. The one adopted here is at 

least equal in value to that followed in the older method of studying 

all the salts of one metal together.

In making a laboratory investigation it may happen that the 

methods employed do not reproduce natural conditions closely. 

For instance, in studying the behavior of gold, platinum, and silver, 

soluble salts are used in the laboratory, whereas in nature soluble 

salts are rarely found/ and the essential problem is to determine the 

conditions of solution and transportation of those metals rather 

than their precipitation. Frequently, however, the assumption 

is warranted that mixtures of soluble metallic salts in noteworthy 

concentrations may come under the influence of precipitating agents 

such as sulphides, sulphates, carbonates, and hydroxides and be hi 

part or completely precipitated. The precipitating action of solids 

will differ from that of solutions only in respect to the concentration 

of the precipitant available and the rate of the action, the action 

approaching that of an extremely dilute precipitant in solution.

In considering precipitation several elemental concepts have 

wide applicability, such as the law of mass action, ionization, and 

the constant solubility product. For the benefit of those who 

may not have access to detailed treatises on these subjects it seems 

appropriate to explain these concepts briefly, in so far as they are 

applicable to the present investigation, an investigation of solutions 

in which partial precipitation occurs. Thelawof mass action accounts 

for the equilibria occurring in the solution, the ions determine in 

the most general way the properties of inorganic solutions, and 

solubility finds its best interpretation through the solubility product.
8 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

THE MASS LAW.

The law of mass action has very wide application. Suppose that 

in any homogeneous system at a definite temperature there is possible 

a general reaction or equilibrium * of the type.

then, according to the law of mass action,

TT x

ambn......

where a, &, c, d represent the concentrations of the molecular species 

A, B, C, D, and K is a constant, the equilibrium constant. It follows 

from this relation that an increase in the concentration of any substance on one side of the equation must result in an increase in the 

concentration of the substances on the other side and a decrease 

in the concentration of the other substances on the same side.

IONS AND THE SOLUBILITY PRODUCT.

This is not the place to present arguments for and against the 

theory of electrolytic dissociation suggested by Arrhenius in 1887. 

Almost every recent work on chemistry shows abundantly the fruitfulness of the conception of ions and presents the generalizations 

based on it.2 Moreover, the phenomena of electric conduction 

warrant the assertion that the existence of ions is not only theoretically possible but is practically demonstrated.

The statement that a substance " ionizes" in solution does not 

mean that it ionizes completely but only that its characteristic ions 

appear among the "molecular species" contained in the solution and 

are potential factors of chemical reactions. The written symbols for 

ions indicate the nature of their electric charges that is, they show 

whether they are positive or negative; for example, H+ , Cl~; for of

course such ions are entirely different from free atoms. Moreover, 

every solution must contain equivalent numbers of positive and negative ions.

As "molecular species" contained in the solution the ions as well 

as the molecules may be considered with reference to the mass law. 

Such a consideration leads to the conception of the "dissociation 

constant." 

Thus,

gives

, [H+] paco.-]

~ [H2C03 ] 

the dissociation constant for the primary dissociation of carbonic acid.

1 The sign ^^ indicates that the reaction may occur in either direction, according to conditions, and, 

according to the usual interpretation, equilibrium is merely the condition attained when the rate of the 

two opposing reactions is equal.

2 For example, see Stieglitz, Julius, The elements of qualitative analysis, The Century Co., 1911 .
IONS AND THE SOLUBILITY PKODTJCT. 9

With respect to chemical precipitates it was early assumed that the 

concentration of the nonionized dissolved portion is not changed by 

the presence of other salts/ but it has since been discovered that 

large deviations from this rule may be shown by moderately concentrated solutions. 2 The great number of solutions in which it holds 

fairly well; however, warrants its consideration, especially as it loads 

to a principle of even wider applicability, namely, that of the constant 

solubility product.

The solubility product (S. P.) is the product (result obtained by 

multiplication) of the concentrations of the ions yielded by the precipitate in the solution, which must be exceeded before precipitation 

can occur. Dr. John Johnston, of the Geophysical Laboratory of 

the Carnegie Institution of Washington, has suggested to me the 

need of a simpler term than " solubility product" for the mathematical function under discussion. The words "solubility product" 

merely denote a number that is a characteristic property of each 

specific substance. The evaluation of these numbers is one of the 

constant aims of chemists. There will be as many powers in this 

product as there are ions formed from the- salt under examination. 

Thus for a uni-univalent salt 

S.P. = [B][A]

in which [B] and [A] are the concentrations expressed as gram molecules or gram ions per liter. For unibivalent salts the solubility 

product takes the form.

S. P. = [B] 2 [A]or[B][A]2

and so on for more complicated salts.

Therefore if to a precipitate in a saturated solution a salt having a 

common univalent ion is added, the concentration of the other ion or 

ions of the precipitate will be decreased. Recent investigations have 

shown, however, that the effect of the addition of a common bivalent 

ion is insignificant.3 It is not yet clear whether these results indicate 

a failure of the theory or whether, as seems more likely,- correct 

evaluations of the ionic concentrations have not yet been obtained. 

However, though in many solutions the relations are still considered 

only qualitatively, the more insoluble the precipitate considered the' 

more certainly the theory appears to bo substantiated.4 An application of the principle to sulphides, hydroxides, and carbonates of the

1 Noyes, A. A., Ueber die gegenseitige Beein'lussung cler Loslichkeit von dissociierten Korpern: Zeitschr. 

physikal. Chemie, vol. 6, p. 243, 1890.

 2 Arrhenius, S., Ueber die Aenderung der StfLrke schwacher Siiuren durch Salzzusatz: Idem, vol. 31, 

p. 224,1899.

3 Noyes, A. A., and Bray, AV. C., The effect of salts on the solubility of other salts: Am. Chem. Soc. Jour., 

vol. 33, p. 1648,1911.

* Harkins, W. D., The solubility of unibivalent salts in solutions of salts of different types: Am. Chem. 

Soc. Jour., vol. 33, p. 1827,1911.
10 FRACTIONAL PRECIPITATION" OF ORE-FORMING COMPOUNDS.

heavy metals would thus appear not wholly unwarranted, and some 

applications of it have therefore been made in this bulletin.

Of course it remains true that if a substance with no common ion 

is added to a solution the solubility of the precipitate is, in general, 

slightly increased, for a little of it goes to form every possible new 

species of molecule by metathesis; and furthermore the solubility of 

a substance, even in the presence of a common ion, may be greatly 

increased under special conditions on account of the formation of complex ions.

FKACTIONAL PRECIPITATION.

PREVIOUS INVESTIGATIONS.

Most of the separations made in analytical chemistry depend on 

pronounced differences in solubility and the use of an excess of reagent, 

but in geochemistry we have to consider associated substances whose 

solubility may differ but slightly and whose formation was due to 

fractional precipitation. A fractional precipitation is one in which 

only a part of the dissolved substances pass from solution to solid.

The theoretic side of fractional precipitation was treated extensively by Berthollet over a century ago. As regards fractional precipitation Berthollet was correct in contending that the composition 

of the precipitate might be indefinite, and Proust's view as to the 

definite composition of the precipitate was correct in regard to pure 

simple chemical compounds. Berthollet held that the fundamental 

factors in. chemical reactions are cohesion, elasticity, and mass relations, governing, respectively, solubility, volatility, and precipitation.

Fractional precipitation of the rare earths has long been familiar to 

chemists. In this way the constituents of the gadolinite earths were 

separated by Mosander; 1 praseodymium and neodymium. were 

separated from the cerite earths by Welsbach; the constituents of the 

yttria earths were separated by Crookes;2 and the sulphide of polonium by Curie.3 Debus,4 in 1853, made quantitative experiments on 

the fractional precipitations of barium and calcium carbonate, and 

Chizynski 5 later studied the phosphates of these two metals. Chroustchoff and Martinoff 6 studied the precipitation of sulphates and chromates by a barium salt and Kiister and Thiel the fractionation of 

bromides and thiocyanates by silver nitrate. 7 Recently Golblum

1 Mosander, C. G., On yttria, terbium, and erbium: Philos. Mag., 3d scr., vol. 23, p. 252,1843.

2 Crookes, William, On the methods of chemical fractionation: Chem. News, vol. 54, pp. 131,155,1S86. 

s Curie, Sklodowska, Radioactive substances: Chem. News, vol. 88, p. 146,1903.:

4 Debus, Heinrich, Ueber chemische Verwandtschaft: Liebig's Annalen, vol. 85, p. 124,1853.

5 Chizyriski, Anton, Ueber die chemische Massenwirkung: Annalen der Chcmie und Pharmacie, Suppl. 

Bd. 4, p. 226,1865.

6 Chroustchofl, P., and Martinoff, A., Des coefficients d'affinite" chimique: Compt. Rend., vol. 104, 

p. 571,1887.

7 Kiister, F. W.,and Thiel, A., Ueber Gleichgewichtserscheinungen bei Fallungsreaktionen: Zeitschr. 

anorg. Chemie, vol. 33, p. 129,1902.
METHOD OF EXPERIMENTATION. . 11

and Stoffella * have studied the system involving lead carbonate and 

chromate.

The advantage of studying fractional precipitation for a theory of 

ore deposits is that the results show the net effect of all the factors 

that enter into play, including even unsuspected factors. Thus, if 

the salt of one metal is more hydrolyzed in solution than another, or 

if a metal has a tendency to form basic salts, the effect will be manifest in the fractionation.

An objection that may be made to many laboratory experiments 

is that chemical precipitates are formed instead of natural minerals. 

For some purposes this objection is valid; for others not. In general 

such experiments will shed light on the relations to be expected in 

the formation of minerals. In fact in many experiments it has been 

shown that in the course of a few hours the precipitate alters even at 

ordinary temperature into the mineral, as does lead sulphide.2 In 

other experiments the rate of alteration is very slow, as Alien and 

Crenshaw have shown for zinc sulphide,3 and solubility determinations are needed for the minerals as well as for the amorphous precipitates. I have, however, found, in every experiment in which the 

effect of time has been noted, that when equivalent quantities of two 

metallic salts are taken the experiment shows almost at once which 

of the two metals forms the more insoluble compound. There may 

be a decrease in solubility with time or a relative change in the solubilities on account of the very gradual development of the most stable 

compounds, but in no experiment did the metal first precipitated in 

excess later become the more soluble.

METHOD OF EXPERIMENTATION.

The metallic compounds will be discussed in the order of their solubility sulphides, hydroxides and oxides, carbonates, bicarbonates, 

and silicates.

The procedure in the experiments on fractional precipitation was 

simple. A dilute solution containing two metallic salts in equivalent 

quantities was precipitated by reagent enough for one metal only. 

After a time an aliquot part of the mother liquor was analyzed and 

the composition of the precipitate was determined by difference. In 

some experiments the proportions were varied slightly, as noted.

Since the quantities of salts used are frequently expressed in equivalents a table has been prepared to assist in making conversions from 

units of weight to " equivalents." An equivalent is the molecular

1 Golblum, H., and StofEella, G., Contribution a Pe'tude de 1'affinite' chimique: Jour. chim. physique, 

vol. 8, p. 135,1910.

2 Weigel, 0., Die Loslichkeit von Schwermetallsulflden in reinem Wasser: Zeitschr. physikal. Chemie, 

vol. 58, p. 293,1907.

8 Albn, E. T., and Crenshaw, J. L., The sulphides of zinc, cadmium, and mercury: Am. Jour. Sci., 4th 

ser., vol. 34, p. 358, 1912.
12 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

weight of a radicle (or atomic weight of an element that constitutes a 

radicle) reduced to a univalent basis. The equivalents thus calculated correspond to the usual chemical equivalents of salt radicles in 

aqueous solution referred to eight parts of oxygen. Of course, 

equivalent weights of salts may be taken in milligrams, grams; pounds,' 

or tons, but wherever no qualification is expressed gram equivalents 

are always understood.

Table for computing "equivalents."

For weight given of Divide by 

Ag............................. 107.87

Al.............................. 9.033

Ba..........:................... 68.68

Cd............................. 56.20

Ca............................. 20.03

Co............................. 29.48

Cu............................. 31.78

Fe"............................. 27.92

Fe"'.......................... 18.61

H.............................. 1.008

Eg7 .........................'.. 200.6

K.............................. 39.09

For weight given of Divide by 

Mg.............................. 12.16

Mn"..,......................... 27.46

Na.............................' 23.00

NL............................. 29.34

Pb............................. 103.59

Zn.............................. 32.68

01............................... 35.45

C03 "..........-................. 30.00

ECO/........................... 61.01

N03 ............................ 62.01

SO/7 ............................ 48.04

ESO/........................... 97.08

SULPHIDES.

SOLUBILITY OF SULPHIDES.

As these experiments were made to determine the order of solubility 

of the compounds studied, it will be advantageous to present first such 

data as already exist on the subject.

The sulphides of the heavy metals are very insoluble. Although 

they are ordinarily considered quite insoluble their solubilities can be 

determined approximately by physicochemical methods, and, in fact 

it has been shown that mercuric sulphide is many thousandfold more 

insoluble than manganous sulphide. The solubility of the sulphides 

in water at 18° C. was determined by Weigel 1 by the conductivity 

method, but the determination of the solubility by this method is 

complicated by the fact that the exact nature of the dissolved substances must be known, and as the' sulphides produce solutions differing in alkalinity the method could hardly be expected to yield 

more than mere approximations to the true solubility.

i Weijjel, Oskar, Die Loslichkcit von Schwermetallsulfiden in reinem Wasser: Zeitschr. physikal. Chemie, 

vol. 58, p. 293,1907.
SOLUBILITY OF SULPHIDES. 13

Solubility of the sulphides of the heavy metals in distilled water at 1S° C. according to Weigel.

[MolesXKH per liter.]

Crystallized sulphides.

SnS. 0 14

0.55

1 13

1 18

1 21

3 1

6.63

6 55

8,99

L6. 29

W 82

tO. 84

18.89

>3. 6

>4.5

Freshly pr( 

HgS........

Bi2S3 . ......

Ag2S......

As2S3 ......

CuS.. ......

PbS.......

SbjSj.......

CdS.. ......

NiS- ......

CoS.. ......

FeS.......

ZnS. .......

MnS.......

jcipitated (probably amorphous) 

sulphides.

...................... 0.05

...................... 0. 35

..................... 0. 55

..................... 2. 1

...................... 3.51

...................... 3. 6

...................... 5.2

...................... 9.00

...................... 39.87

...................... 41.62

...................... 70. 1

...................... 70.6

....................... 71.6

SnS2 ............................

Galena (from precipitated PbS)... 

Galena (Freiberg)................. 1. 21

Cu2S............................ 3.1

Zinc blende (artificial)............. 6. 63

Zinc blende (Santander).......... 6. 55

Greenockite (artificial)............. 8. 99

Millerite (artificial)............... 16. 29

Wurtzite (artificial)................ 28. 82

Pyrite (artificial)................. 40.84

Pyrite (Freiberg).............-.

Pyrrhotite....................... 53. 6

MnS...........................

The solubility products of sulphides have been determined by other 

investigators from time to time.1

For a univalent, bivalent, and trivalent metal, respectively, the 

solubility products (S. P.) would be:

S. P. = [M+]2 [S--]

S. P. = [

The values of the solubility products already published were collected by Bruner and Zawadzki, who also calculated them from the 

heat of formation of the sulphide and the electrolytic potential of the 

metals and of sulphur. In the table below I have placed their calculated values of the solubility products, those found by the observer 

noted, and finally the "solubility of the sulphide in water" calculated 

from the solubility product. It will be noted that the solubilities 

in this table are far smaller than those obtained by Weigel.

i Bernfeld, I., Studien uber Schwefelmetallelektroden: Zeitschr. physikal. Chemie, vol. 25, p. 46, 18C8. 

Knox, Joseph, A study of the sulphur anion and of complex sulphur anions: Faraday Soc. Trans., vol. 4, 

p. 48, 1908. Imraerwahr, C., Beitrage zur Kenntnis der Loslichkeit von Schwermetallniederschlilgen auf 

elektrochemischen Wege: Zeitschr. Elektrochemie, vol. 7, p. 478, 1901. Glixelli, S., Zur Theorie der 

H2S Fallung der Metalle: Die Eimvirkungen von SchwefelwasserstofE auf Zinksalze: Zeitschr. anorg. 

Chemie, vol. 55, p. 306, 1907. Lucas, R., Gleichgewichte zwischen Silbersalzen: Zeitschr. anorg. Chemie, 

vol. 41, p. 211, 1904. Bruner, L., and Zawadzki, J., Ueber die Gleichgewichte bei der Schwefelwassersto3follong der Metalle: Zeitschr. anorg. Chemie, vol. 65, p. 136, 1910.
14 FBACTIOFAL PBECIPITATION OF OBE-FOBMING COMPOUNDS.

Solubility products of several sulphides and their solubility in water.

[Based on determined solubility product, if known; otherwise on solubility product calculated by Bruner

and Zawadski.]

Sulphide.

MnS.. .............

ThS. ..............

FeS................

«ZnS. .............

NiS....... .........

CoS................

PbS...............

CdS. ..............

Ag2G.. ..............

Ag2S.. .............

Ag2S.. .............

CuS...............

CuS................

HgS...............

Solubility product 

calculated.

7.0X10-16

(4.5X10-23)

7.0X10-22

6.0X10-24

7.0X10-25

2.1X10-23 i.sxio-29

I 1.4X10-48

\ 7.0X10 45

a 2 OXIO"52

a i OXIO"49

Solubility product 

found.

(4.5X10-23)

3.7X10-19

1.1X10-24

5.0X10-26

3.4X10-28

5.0X10-29

4.8X10-53

1.5X10- 

3.9X10-50

5.9X10-42

1.2X10-42

2.8X10-54

6.7X10-48

Observer.

wadski. 

..... do. ...........

.....do............

Glixelli...........

.....do............

wadski. 

.....do............

.....do............

.....do............

.....do............

Solubility in water.

Moles per liter (.calculated).

9 RV1ft-8

2.2X10-8

1.9X10-11

1.0X10-12

2.2X10-13

8.4X10-' 2

1.8X10-14

4.2X10-15

2.3X10-18

1.9X10-"

2.2X10-17

2.4X10-21

1.1X10-21

1.7X10-2?

2.4X10-24

a Based on different heats of formation.

The order of solubility obtained from experiments on fractional 

precipitation should agree with the order of solubility given in the 

last column of the table. As a matter of fact, the agreement is excellent except for two metals, silver and cobalt. The discrepancy for 

silver requires explanation; that for cobalt is not surprising in view 

of the small difference between its solubility and those of the adjacent 

sulphides. The discrepancy for silver may be due to a failure of the 

principle of the solubility product, as silver sulphide is a unibivalent 

compound, or possibly to the precipitation of silver by sulphides in 

part as free metal. A study of the fractional precipitation of thallous 

sulphide would assist in deciding between these two possibilities. 

Weigel's series seems to place silver in its proper position, but it 

shows several discrepancies for other metals.

While the solubility of sulphides is under consideration attention 

may be paid briefly to the enormous effect of changes in acidity on 

the precipitation of sulphides. This effect is due to the nature of the 

ionization of hydrogen sulphide, the sulphide ion concentration varying inversely as the square of the hydrogen ion concentration for a 

given concentration of total sulphide. As is well known, the precipitability of the sulphides varies so greatly that a complete separation of 

certain groups is possible. Thus in an acid solutioii.hydrogen sulphide 

added to a dilute mixture of the sulphates of iron and copper precipitates almost wholly copper sulphide. This is not strictly a 

fractional precipitation, since in an acid solution of hydrogen sulphide the concentration of sulphide is probably not sufficient to 

exceed the solubility product of ferrous sulphide. A fractional precipitation would be obtained with a small amount of precipitant
EXPERIMENTS WITH SULPHATES OF COPPER AND ZING. 15

under conditions which would permit all the metals to be precipitated with an excess. With iron and copper sulphates, for example, a fractionation might be made with sodium sulphide, since 

an excess of sodium sulphide would precipitate both sulphides completely.

. EXPERIMENTS WITH THE NITRATES OF COPPER AND LEAD.

The first experiments on this point showed that the composition 

of the precipitate depends a great deal on the manner of precipitation 

and time of standing. A mixture of the nitrates of lead and copper, 

partly precipitated with a solution of hydrogen sulphide in the cold 

and filtered within a short time, yielded a precipitate containing the 

sulphides of both metals, although copper was slightly in excess. 

On warming, however, copper sulphide was the chief product, being 

precipitated even at the expense of the lead sulphide first formed.

EXPERIMENTS WITH THE SULPHATES OF COPPER AND ZINC.

A good illustration of selective precipitation was obtained with 

mixtures of cupric sulphate and zinc, sulphate, precipitated by 

ammonium sulphide. The data are given in the table below and 

the curves in figure 1. The evidence is very clear that even in

-.-£ 2-

Cu 0

Zn 4 Taken'

FIGURE 1. Selective precipitation of copper before zinc by ammonium sulphide. The abscissae represent 

the content of the solution in milligram ecmivalente of cupric sulphate and zinc sulphate before the addition of 2 milligram equivalents of ammonium sulphide; the ordinates represent the composition of the 

resulting precipitate.

alkaline solution more cupric sulphide is precipitated than zinc sulphide. The quantities of substances used are expressed in milligram 

equivalents. (See p. 12.) The solutions were dilute at the time of 

precipitation, the final volume being 40 cubic "centimeters for the 

quantities stated in the table.
16 FRACTIONAL PRECIPITATION OF OEE-FOEMING COMPOUNDS.

Fractional precipitation of cupric and zinc sulphates by ammonium sulphide. 

[2.00 milligram equivalents in each precipitation.]

Cupric 

sulphate 

taken.

Milligram

equivalents.

0.00

.50

1.00

1.50

2.00

3.00

3.20

4.00

Zinc 

sulphate 

taken.

Milligram

equivalents.

4.00

3.50

3.00

2.50

2.00

1.00

.80

.00

Copper in 

precipitate.

Milligram

equivalents.

0.00

.49

.87

1.29

1.70

1.97

2.05

2.00

Zinc in 

precipitate.

Milligram

equivalents.

2.00

1.51

1.13

.60

.41

.02

.00

.00

These results are stated fully as illustrations of relations which 

hold for most pairs of sulphides. Both precipitate together at first, 

then one largely redissolves, leaving the sulphide radicle in combination with a single metal.

These facts may find application in theories of the formation of 

ores. If large bodies of precipitable solutions come into sudden 

contact, an indefinite mixture of sulphides will be precipitated. On 

the other hand, if the solutions are somewhat warm or react slowly 

the sulphides tend to be precipitated in a definite succession.

WORK OF ANTHON AND SCHtJBMANN.

The precipitation of soluble salts of metals by the insoluble sulphides of other metals was studied long ago by Anthon.1 He obtained the series silver, copper, lead, cadmium, iron, .nickel, cobalt, 

and manganese. A sulphide of any one of these metals, he found, 

would be precipitated at the expense of the sulphide of one lower in 

the series. The subject was studied further in 1888 by E. Schurmann,2 who enlarged the series as follows: Palladium, mercury, silver, 

copper, bismuth, cadmium, antimony, tin, lead, zinc, nickel, cobalt, 

ferrous iron, arsenic, thallium, and manganese. The experiments of 

Anthon and Schurmann lead to exactly the same conclusions as 

those on fractional precipitation. The further investigation of sulphides was therefore restricted to some of the side issues involved. 

(See pp. 18-23.)

Schurmann found that some of these transpositions, of which he 

studied a very large number, occurred more easily than others, particularly those where the metals are widely separated in the series. 

It was especially difficult to obtain complete transformations in the 

case of nickel sulphide treated with zinc sulphate and between 

thallium salts and the neighboring members of the series. The

1 Anthon, E. F., Jour, prakt. Chemie, vol. 10, p. 353,1837.

2 Schurmann, Ernst, Ueber die Verwandtschaft der Schwermetalle zum Schwefel: Liebig's Annalen, 

vol. 249, pp. 326-350,1888.
WOEK OF ANTHON AND SCHUBMANN. 17

behavior of iron was peculiar in that the sulphides of zinc, nickel, 

and cobalt formed at the expense of an excess of ferrous sulphide, 

but they were in turn dissolved when treated with ferrous sulphate. 

The positions of arsenic, tin, and antimony in the series are subject 

to qualifications depending .on the valence of the metal.

This series for the order of solubility of sulphides has thrown great 

light on the development of sulphide ores.1 It indicates the order 

of deposition of sulphides from a mixture of metallic salts when they 

compete under equal conditions for a sulphide. The first metals of 

the series will be precipitated first, leaving the solution richest in the 

last members; or, if solution is going on, the last members will tend, 

other things being equal, to go into solution first.

Schurmann was chiefly interested in determining the place of each 

metal in the series. He considered the series to represent the order 

of the " affinities" .of the metals for sulphur. A more empirical statement is that it represents the order of the solubilities of the sulphides 

in water.2

If the mechanism of these reactions . is considered according to 

present theories, we should have the following scheme of ionization 

in solution, for example, for the couple silver-copper:

S- +Cu++

11 U

Ag2S CuS04

Since the solubility product of silver sulphide is smaller than that 

of copper sulphide, silver sulphide forms at the expense of copper 

sulphide, and the equilibrium lies well toward the right in the following equation:

Ag2S04 + CuS ^ Ag3S + CuS04

The actual relations may be more complex than the last equation 

indicates, because there may also be a reaction of oxidation and reduction, but the scheme shows the general relation. Every sulphide 

has a "solubility product"; that is, the product of the concentrations 

of its two ions, which must be exceeded before precipitation can occur. 

Conversely, if the concentrations of the ions of any sulphide are less 

than its solubility product the sulphide will dissolve. That Schtirmann's series is very nearly the order of solubility of the sulphides 

in pure water is evident from a comparison of the last column in the 

table on page 14 with Schurmann's series on page 16.

1 Emmons, W. H., The enrichment of sulphide ores: U. S. Geol. Survey Bull. 529, p. 56, 1913.

2 Wells, R. C., Econ. Geology, vol. 5, p. 7, 1910.

95921° Bull. 609 15 2
18 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

The idea of the solubility product allows us at once- to explain 

certain facts noted by Schiirmann and to predict others with reasonable confidence. For instance, it is not a matter of primary 

importance what acidic radicles are present in the solution as long 

as they do not form insoluble salts with the metals present. Again, 

the greater speed of adjustment of equilibrium for metals widely 

separated in the series accords with our present idea that the speed 

of reaction is a function of the remoteness of removal from 

equilibrium.

If we let K! and K2 represent the solubility products of two sulphides involved in a fractional precipitation and write

K^OVIJCS) 

K2 =(M2 )(S)

where (MJ, (M2 ) and (S) represent concentrations of metal 1, metal 

2, and sulphide ions, then when equilibrium is established, since the 

sulphide ion is common, we should have

K! (MJ (Mt) ^

ET = (Ml or = 3 ' a new constantThat is, the ratio of the concentrations of the two metallic ions in 

a solution saturated with both sulphides should be a constant, independent of the quantity of sulphides but probably somewhat dependent on the temperature. Schurmann as a rule gave only qualitative 

results, and where he says one metal was "completely" separated 

from the other more exact research might show that a distribution 

occurs. As a matter of fact, however, the relation expressed by 

equation (1) above has not yet been experimentally illustrated in 

a quantitative way for sulphides.

A consideration of the solubility of sulphides helps to explain certain phenomena observed by Stokes.1 He found that pyrite and 

marcasite are slightly attacked by a solution of alkali carbonate 

but much more completely when a carbonate of copper, lead, silver, 

or zinc is present. This is no doubt due in part to the greater insolubility of the sulphides of copper, lead, silver, or zinc, which form at 

the expense of the .sulphide of iron in the pyrite and marcasite 

while the alkali dissolves the excess sulphur, the two concurrent 

reactions being much more effective than either singly.

COMPOUND SULPHIDES.

One of the objects in studying fractional precipitation was to gain 

light on the conditions of formation of compound sulphides, such as 

chalcopyrite. It is easy to perceive that the solubility of one sul1 Stokes, H. N., Experiments on the action of various solutions on pyrite and marcasite: Econ. Geology 

vol. 2, p. 19, 1907.
COMPOUND SULPHIDES. * 19

phide might be changed by the presence of a second in solution and 

that a compound sulphide might have an intermediate solubility, 

following the usual behavior as described by Van't Hoff for double 

salts. The sulphides are so insoluble, however, that no adequate 

study of their behavior over a wide range of conditions has as yet 

been made.

Grout 1 has suggested that compound sulphides may have been 

produced in some instances by the admixture of alkaline and acid 

solutions. This seems very probable. Malfatti 2 has described certain complex sulphides of iron of the formula RFeS2 where R is a 

univalent radicle. But in chalcopyrite copper has usually been 

regarded as bivalent, if the formula is written CuS.FeS; to write it 

Cu2S.Fe3Ss suggests trivalent iron. However that may be, if chalcopyrite could be formed by a simple precipitative reaction its formation would remove the constituents from a solution in definite proportions. Such removal would modify curves like those shown in 

figure 1 (p. 15) into lines with horizontal sections over certain regions. 

As no such horizontal sections were discovered in the fractional 

precipitation of copper and iron sulphide the conditions for the 

formation of chalcopyrite were not revealed by that method.3 It 

was therefore concluded that the formation of chalcopyrite by a 

precipitative reaction at moderate temperature and pressure demands at least more precipitating sulphide than corresponds to the 

copper present; otherwise no iron will be precipitated. Further investigation of this subject is needed.

By microscopic examination of the lean Butte ores J. F. Simpson 

has shown that the pyrite contains the copper as the copper minerals 

chalcocite, bornite, chalcopyrite, and enargite, and that of such minerals chalcopyrite is the oldest.4 Graton'and Murdock have reached 

similar conclusions after an examination of many specimens from 

other localities.5 The typical minerals of enrichment processes appear 

to have been formed by the progressive alteration of one mineral into 

another, not strictly a fractional precipitation, although the fractionating tendency is revealed by the removal of the iron and the increasing deposition of copper. For the formation of the compound sulphides by enrichment the iron sulphide is already present. Of course, 

chalcopyrite may also result from magmatic differentiation, but for 

aqueous solutions we can not avoid the conclusion that its formation

1 Grout, F. F., On the behavior of cold acid sulphate solutions of copper, silver, and gold with alkaline 

extracts of metallic sulphides: Econ. Geology, vol. 8, 1913.

2 Llalfatti, H., Beitriige zur Kenntnis des Eisensulflds: Zeitschr. anal. Chemie, vol. 48, p. 352,1909.

8 For details see Econ. Geology, vol. 5, p. 1,1910.

« Simpson, J. F., The relation of copper to pyrite in the lean copper ores of Butte, Mont.: Econ. Geology, 

vol. 3, p. 635,1908.

6 Graton, L. C., and Murdock, Joseph, The sulphide ores of copper: Am. Inst. Min. Eng. Bull., May, 

1913, pp. 741-812.
20 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

indicates the presence of an excess of available sulphide beyond the 

requirements of the copper. Briefly, a small amount of a soluble sulphide, say sodium sulphide, acting on a considerable quantity of 

ferrous and cupric sulphates does not produce chalcopyrite but only 

copper sulphide. »

REDUCING ACTION OF SULPHIDES.

Besides taking part in purely metathetic reactions both soluble and 

insoluble sulphides may act as reducing agents. This action may be 

referred to the slight electroaffinity of sulphur, which produces the 

partial reaction:

This reaction has a strong tendency to occur whenever it can be 

linked with a reaction of positive ions yielding positive electricity. 

For example, we are familiar with the action of hydrogen sulphide 

in acid solution on ferric salts :

2FeCl3 + H2S = 2FeCl2 + S + 2HC1

If the reaction is considered to be essentially one between ions we 

have :

-*- + S = 2Fe++ + S (2)

This reaction shows the neutralization of the positive and negative 

charges and the precipitation of free sulphur.

The application of the reducing tendency of sulphide ions here 

suggested and illustrated should lead to interesting conclusions with 

other than iron salts. With cupric salts, for example, we should 

expect that the sulphide ions would tend to reduce cupric ions, along 

the line of the following equation:

2Cu++ + S- = 2Cu+ + S (3)

Data on this possibility are scanty. Several investigators l have 

called attention to the fact that hydrogen sulphide produces in a 

solution of cupric sulphate a precipitate in which there is some free 

sulphur and a deficit in the amount of combined sulphur required 

by the formula CuS. According to Schweizer, there is usually about 

7.2 per cent cuprous sulphide (Cu2S) that is, if the complex is interpreted as a mixture of cupric and cuprous sulphide but, according to the more recent work of Clark, if such a precipitate is filtered 

off and digested for a long time with a solution of hydrogen sulphide

 5 

1 Thomsen, Julius, Ueber die Zusammensetzung des auf nassem Wege gebildeten Schwefelkupfers:

Deutsche Chem. Gesell. Ber., vol. 11, p. 2043, 1878. Ditto, A., Action du sulfure de cuivre sur le sulfure de

' potassium: Compt. Rend., vol. 98, p. 1429, 1884. Brauner, Bohuslav, Action of hydrogen sulphide on

cupric salt solutions: Chem. News, vol. 74, p. 99, 1896. Schweizer, E., Diss., p. 34, Erlangen, 1908. Clark,

J. D., A chemical study ol the enrichment of copper sulphide ores: Hew Mexico Univ., Diss., p. 106, 1914.
EEDUCING ACTION OF SULPHIDES. 21

in water the residue after removal of free sulphur contains about 71 

per cent cuprous sulphide.

The study of the ionic reaction suggested by equation (3) as a possible equilibrium governed by the mass law is rendered difficult by 

the fact that sulphur and the two sulphides of copper are very insoluble. Of the different concentrations in equation (3) that of Cu++ 

will be determined by the solubility product of cupric sulphide and 

that of Cu+ by the solubility product of cuprous sulphide. The concentration of free sulphur may presumably be varied from zero to 

that of a saturated solution of sulphur. The concentration of sulphide 

ions, S , may be varied most widely, both by changes in total concentration and by the use of alkaline solutions in which the ionization 

of hydrogen sulphide is enormously increased. The use of an alkaline solution also serves to keep the concentration of free sulphur, S, 

below that of a solution saturated with sulphur on account of the

formation of alkali polysulphide.

If we increase the concentration of the sulphide ions, S , there 

should be an increase in the cuprous ion concentration, Cu+ , if equation (3) is significant, and presumably more cuprous sulphide might 

then be precipitated.

A number of experiments were made, directed wholly toward 

securing a precipitate with a high percentage of cuprous sulphide, 

especially by the use of alkaline solutions. The results are summarized in the following table. The precipitations were made hot by 

adding the cupric sulphate to the alkaline sulphide. The precipitate 

was washed by decantation, finally with alcohol and ether, well 

pressed between filter papers, and dried over sulphuric acid.

Action of alkaline soluble sulphide solutions on cupric sulphate.

Experiment.

126. ..........................................................

127...........................................................

128...........................................................

130. ..........................................................

131. ..........................................................

129. ...........................................................

132. ..........................................................

Cram equivalents taken per 

liter of the mixture.

CuS04 .

0.091 

.059. 

.035 

.001 

.016 

.014

CuCl 2

.05

Na2S.

0.091 

.112 

.069 

.002 

.155 

.028

.10

NaOH.

0.017 

.022 

.130 

.200 

.146 

.158

.10

Percentage 

of Cu2S in 

precipitate.

18.3 

31.1 

47.6 

55.8 

61.8 

69.3

49.8

These results prove that the immediate precipitate produced by 

alkaline sulphides contains a much higher percentage of cuprous; 

sulphide than that formed by hydrogen sulphide in acid solutions. 

The results strengthen the supposition that the equilibrium shown 

in equation (3) determines the composition of the precipitate to some
22 FRACTIONAL PRECIPITATION OP ORE-FORMING COMPOUNDS.

extent, since the sulphide ion concentration varies with the alkalinity.1 It appeared impossible, however, by this method to obtain 

a precipitate of pure cuprous sulphide. The experiments of Ditte 

show that a concentrated solution of potassium sulphide will convert 

copper entirely into the cuprous state, but he says that a dilute 

solution is without effect.2 My experiments seem to modify Ditte's 

conclusions. The further solution of the problem would seem to require a more careful study of the reversibility of the action with 

respect to variations in the concentrations of sulphide ions'and free 

sulphur.

The reducing action of sulphide ions on other metallic salt solutions 

has not been studied.

ELECTRIC ACTIVITY IN SULPHIDES.

Since insoluble metallic sulphides are conductors of electricity the 

equalization of chemical differences, especially differences in degree 

of oxidation, can take place through electric action over appreciable 

distances. Thus the deposition of such metals as gold, silver, and 

copper may occur at one end of a conducting filament while the oxidation of a sulphide or any other oxidizable substance is occurring at 

the other end of the system constituting the metallic portion of the 

circuit. The associations of products produced in this way will 

naturally be different from those produced under conditions of 

thorough admixture. Further, not only will dissimilar sulphides 

and metals manifest reciprocal influences on each other through 

electric activity but effects from active solutions will be transmitted 

as far as the complete circuits extend. I have discussed these matters in detail elsewhere 3 and need refer to only one phase of the 

subject further in this connection.

In studying chalcocite enrichment' A. C. Spencer has reached the 

conclusion that the change of pyrite or chalcopyrite to chalcocite 

may be considered as an alteration involving a series of steps, or 

perhaps even a continuous progression through indefinite compounds 

or mixtures of iron-copper sulphides. He succeeded in bringing 

about the change of chalcopyrite at least into chalcocite through 

bornite and covellite in several artificial ways. One of these ways 

consisted in touching the specimens immersed in cupric sulphate by 

an iron nail. Since there is every reason to believe that the iron 

affected the sulphides through electric action it seems worth while 

to quote his account of this experiment as an illustration of possible 

electric action.4

1 Knox, Joseph, A study of the sulphur anion and of complex sulphur anions: Faraday Soc. Trans., 

vol. 4, p. 47,1908.

2 Ditte, A., Compt. Rend., vol. 98, p. 1429,1884.

s Wells, R. C., Electric activity in ore deposits: U. S. Geol. Survey Bull. 548,1914. 

j 4 Spencer, A. C., Chalcocite enrichment: Econ. Geology, vol. 8, p. 629, 1913.
SUMMARY OF RESULTS OF EXPERIMENTS WITH SULPHIDES. 23

The colors * * * referred to above may be readily obtained by another simple 

procedure, and the observer can hardly fail to conclude that the colors obtained 

indicate the formation of the minerals bornite, covellite, and chalcocite. If any 

member of the series chalcopyrite, bornite, covellite is touched by a piece of iron 

while it is immersed in a cupric sulphate solution the mineral changes color almost 

instantly and in a short time becomes coated with the mineral next above it in the 

series. The brilliant indigo of covellite changes to the dull gray so characteristic of 

chalcocite, bornite assumes a blue color unmistakably like that of covellite, and 

chalcopyrite takes on a bronzy hue resembling that of bornite. Furthermore, within 

a very short time the bronze plating on chalcopyrite gives place to or is hidden by a 

film of covellite; then within an hour or so the surface turns to a chalcocite gray, and 

finally metallic copper is deposited. When pyrite is used in place of a cupriferous 

sulphide the effect of the iron ia sufficient to throw down metallic copper rather 

quickly, but by scraping away the metal and again placing the mineral in contact with iron in the solution it is possible to obtain deposits of copper sulphides. 

In this way spots having somewhat the same color as chalcopyrite and others 

bronzy like bornite may be developed on pyrite along with unmistakable films 

of covellite and of chalcocite. 'If copper is used instead of iron the results are 

essentially the same with chalcopyrite, bornite, and covellite. For instance, 

covellite may be coated with chalcocite by contact with copper in a solution of 

cupric sulphate. It is obvious that the speed of reaction may be varied by employing different metals as inductors, or by employing minerals to cause electrolytic 

action. Very pleasing results have been obtained by placing in a solution of cupric 

sulphate a polished specimen of intergrown chalcopyrite and pyrrhotite. Here a 

pinkish bronze color resembling that of bornite appeared within a few days, but 

gradually changed to purple, deep purple, and finally to indigo blue. On the most 

reactive grains the covellite color was fully developed in about eight weeks, but the 

surface in general became blue only after 12 weeks, and even then certain areas were 

still bronzy. At the end of four months no gray color had developed to indicate the 

formation of chalcocite, but the color was a paler blue than that of natural covellite.

It may be suggested that the results described, which were obtained under ordinary 

temperatures, may actually epitomize the course of reaction between the primary 

sulphides and copper salts held in oxidized solutions penetrating from the surface. 

Even if the means employed to produce the results in a short time are not comparable 

with those involved in natural processes, perhaps the conditions under which the 

experiments were made may be considered to be less unnatural than those prevailing 

in mineral syntheses effected under high temperatures.

I 

SUMMARY OF RESULTS OF EXPERIMENTS WITH SULPHIDES.

1. The concentration of the sulphide ion is so greatly affected by 

change of acidity that this change is the principal factor determining 

the precipitation of sulphides.

2. The solubility of the several sulphides, however, differs so 

greatly that complete separations by fractional precipitation are 

possible; for example, copper from zinc.

.3. A mixture of two metallic salts yields, by fractional precipitation, an initial precipitate containing the sulphides of both metals, 

but as a rule, if the mixture is heated or is permitted to stand, one 

sulphide largely or wholly redissolves.

4. The order of precipitation, beginning with the metal that 

separates first, is palladium, mercury, silver, copper, bismuth,
24 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

cadmium, antimony, tin, lead, zinc, nickel, cobalt, ferrous iron, 

arsenic, thallium, manganese.

5. Attempts to form chalcopyrite by fractional precipitation of 

ferrous and cupric sulphate were unsuccessful.

6. Soluble sulphides may act as reducing agents. This action 

affects the composition of precipitates of copper sulphide and may 

explain the anomalous position .of silver in the precipitation series.

7. At one stage of the action produced by an admixture of sulphides 

and metallic salts electric activity may determine the products 

formed.

HYDROXIDES.

f- PREVIOUS WORK.

Only a few experiments are found in the literature on the fractional 

precipitation of the common metallic hydroxides. Mills and Smith,1 

in -working on nickel and cobalt, found a very narrow margin of 

difference in precipitability, but this is probably one of the most 

difficult pairs to separate, although my own experiments indicate 

that manganese and magnesium are similarly close together. Fink 2 

obtained the precipitation series copper, zinc, nickel, cobalt, manganese, and magnesium for the hydroxides of these metals. Pickering 3 

studied the partial' precipitation of copper sulphate solutions only 

by calcium hydroxide, but this was not strictly a fractionation.

SOLUBILITY OF HYDROXIDES.

The preliminary data available in regard to the solubility of the 

hydroxides at room temperatures of 20°-25° C., are as follows:

Solubility of hydroxides.

Gram equivalents per liter. 

Zn(OH)2 .................................................. 6.0X10-*

Cd(OH)2 ........................:......................... 8.7X10-6

Mn(OH)2 .................................................. 3.0X10-5

Ag30...................................................... 4.6X10-5

Pb(OH)a ................................................... 9.3X10-5

Fe(OH)2 (Lamb 4 ).......................................... 2.7X10~5

Mg(OH)a .................................................. 7.5X 10~4

HgO....................................................... 2.5X10-4

Ca(C)H)2 .................................................... 4.5X10-2

Sr(OH)2 ..........................:........................ 1.3 X10-1 - Ba(OH)2 .................................................. 4.6X10-1

iMills, E. J., and Smith, J. J., Researches in chemical equivalence: Roy. Soc. Proc., 5thser., vol.29, 

p. 181,1879.

2 Fink, R., Ueber die Affinitat der Vitriolmetalle zur Schwefelsaure: Deutsche chem. Gesell. Ber., vol. 20, 

p. 2106,1887.

a Pickering, S. U., The chemistry of Bordeaux mixture: Chem. Soc. Trans., vol. 91, p. 1989,1907; vol. 97, 

p. 1852, 1910.

* Lamb, A. B., The potential of iron calculated from equilibria measurements: Am. Chem. Soc. Jour., 

vol. 32, p. 1218,1910.
FBACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS. 25

FRACTIONAL PRECIPITATION OF HYDROXIDES.

By the procedure used in the study of the sulphides the hydroxides 

of the important metals were fractionally precipitated as follows: 

To a dilute solution containing two metallic salts in equivalent 

quantity, heated nearly to boiling, was added a dilute solution of 

sodium hydroxide sufficient to precipitate one metal only. The 

sodium hydroxide contained some barium hydroxide, used to free it 

from carbonate. Oxidizable salts were precipitated in an atmosphere of hydrogen. The precipitate was kept at a temperature 

close to 100° for several minutes and finally filtered off. An aliquot 

portion of the filtrate was analyzed and the composition of the precipitate was thus determined. In general one metal was precipitated to a greater extent than the other, and after working out a 

number of pairs it - was found possible to arrange the metals in a 

series, which may be termed the hydroxide precipitation series. 

Although in the experiments with the sulphides the proportions in 

the precipitate were determined, in nearly every experiment with the 

hydroxides the results were limited at first to the determination of 

the order in the series. This order was found to be as follows:

Order of solubility of hydroxides.

1. Ferric hydroxide.

2. Aluminum, hydroxide.

3. Cupric hydroxide.

4. Zinc hydroxide.

5. Lead hydroxide.

6. Nickelous hydroxide.

7. Silver hydroxide.

8. Ferrous hydroxide.

9. Manganous hydroxide.

10. Magnesium hydroxide.

11. Calcium hydroxide.

The pairs determined are shown in the next table.

Pairs of metals compared in fractional precipitation of hydroxides.

Experiment.

113

111......................................

Q7

98. ......................................

109......................................

%

100 10.......................................

102......................................

99. ......................................

inr>

101......................................

103......................................

105......................................

109

110......................................

Major constituent of precipitate.

.....do................................

.....do................................

.....do................................

.....do................................

.....do................................

.....do................................

.....do................................

.....do................................

Silver. . ...............................

.....do................................

Minor constituent of 

precipitate.

Do.

Zinc.

Nickel.

Nickel.

Silver.

Do.

Do.

Do.

It will be seen that the precipitation series does not follow the 

'solubility series very closely. Possibly some of the solubility determinations are incorrect, or the matter of hydrolysis may not have 

been sufficiently considered. However that may be, it is significant
26 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

that the precipitation series may be correlated with a heat effect, 

namely with the heat of the reaction

Sodium hydroxide+metallic sulphate=metallic 

hydroxide+sodium sulphate.

These heats of reaction are shown in the table below.

Order obtained for the heat of reaction of sodium hydroxide with metallic sulphates.

M

Hg+ + . .......,......:......"..

Hg+.. ........................

Fe+++... .....................

Ag. ..........................

Cr ............................

Cu. ..........................

Al... .........................

1.

10. 05

7.47

6.47

5.19

2

11.78

10.57

8.,24

6.25

M

Pb. ....... ...................

Zn. ..........................

Cd. ..........................

Co............................

Fe++.. .......................

Ni... ...... ...................

Mn. ..........................

1

3.99

3.59

3.35

3.23

2. fi4

2.45

2

4 7Q

3.77

3.56

^ 13

3.01

2.42

2.23

1. NaOnaq+M(SO4 )i/2aq=Na(SO4 )1/ 2a'ci+MOH.

2. NaOHaq+MN03aq=NaN03aq+MOH.

Although the order here shown does not yet agree exactly with the 

order obtained in the fractional precipitations the agreement is closer 

than that between the precipitation series and the order of the 

solubilities determined directly (p. 24).

FERROUS AND CUPRIC HYDROXIDES.

For the purpose of studying one fractionation in greater detail the 

precipitation of iron and cupric hydroxides was studied over a wide

4 Fe Taken

Cu4

0

FIGURE 2. Apparent selective precipitation with ferrous and cupric sulphates. The abscissa; represent 

the content of the solution in milligram equivalents of ferrous and cupric sulphate before the addition 

o f 2 milligram equivalents of ammonium hydroxide, the ordinates the composition of the resulting precipitate.

range of concentrations. The fractionations in the table below were 

carried out with ammonium hydroxide as the precipitant. The 

results are also plotted in figure 2. Dilute solutions were used for
FERROUS AND CUPRIC HYDROXIDES. 27

these precipitations, the final volume in every experiment being 40 

cubic centimeters for the quantities stated in the table.

Precipitation of ferrous and cupric sulphate by ammonium hydroxide (milligram equivalents).

Cupric 

sulphate 

taken.

0.00

.80

1.10

1.50

1.70

2,00

2.50

2.85

2.90

3.00

3.20

4.00

Ferrous 

sulphate 

taken.

4.00

3.20

2.90

2.50

2.30

2.00

1.50

1.15

1.10

1.00.

.80

.00

Copper in 

precipitate.

0.00

.80

1.08

1.03

1.08

1.09

1.12

1.20

1.26

1.33

1.66

2.75

Iron in 

precipitate.

.2.43

1.30

1.01

1.01

1.05

1.07

.97

.95

.91

.87

.68

.00

Color of precipitate.

Pale green.

Black.

Do.

Brownish.

Orange.

Do.

Do.

Green.

These results show a very peculiar selective precipitation. At the 

extreme concentrations basic salts are precipitated. Although it 

seemed at first as if the results indicated the formation of a precipitate 

containing one atom of iron to one of copper, it was afterward concluded that the oxidation of the ferrous salt had not been wholly prevented, and in fact that it could not be wholly prevented on account 

of the presence of cupric sulphate. As a result of this oxidation ferric 

hydroxide was precipitated instead of ferrous hydroxide, and therefore the assumption of the formation of a complex compound is unwarranted.

Similarly solutions of cupric sulphate and ferrous sulphate in varying proportions were precipitated hot out of access of air by sodium 

hydroxide.

Precipitation of ferrous and cupric sulphate by sodium hydroxide (milligram equivalents).

Ferrous 

sulphate 

taken in 

250 cubic 

centimeters.

1 fiO

A on

C.20

Cupric sulphate 

taken in 

250 cubic 

centimeters.

6.50

4.00

1.80

Iron precipitated 

by 4 milligram equivalents of 

sodium hydroxide.

1.15

1.78

2.03

Copper precipitated.

1.79

These results show that the copper has a greater tendency to precipitate where equal amounts of both salts are present, and this tendency causes the practical removal of the copper where it is the minor 

constituent. What looks like evidences of the formation of a molecule FeCu(QH)4 may again be due simply to unpreventable oxidation 

of a portion of the ferrous salt.
28 . FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

Recently Curtman and St. John l have studied the sensitiveness of 

the hydroxide reactions for the common metals. They find an order 

of sensitiveness which differs considerably from the precipitation 

series here deduced, but it appears that their results are based on optical effects, such as the visibility, color, and density of the precipitates, 

so that an exact agreement is not to be expected.

SUMMARY OP RESULTS OF EXPERIMENTS WITH HYDROXIDES.

1. The precipitation series "obtained for the hydroxides, beginning 

with the most insoluble compound, is as follows: Ferric hydroxide, 

aluminum, cupric hydroxide, zinc, lead, nickel, silver, ferrous hydroxide, manganous hydroxide, magnesium, calcium.

2. The series is very nearly the same a§ that for the heat required 

for the formation of the hydroxides, namely: Mercury, ferric hydroxide, silver, copper, aluminum, lead, zinc, cadmium, cobalt, ferrous hydroxide, nickel, manganese.

3. When cupric and ferrous salts are precipitated together by a 

soluble hydroxide the iron tends to appear as ferric hydroxide, and 

this fact explains the apparent exceptional precipitation of the two 

metals together as hydroxides in spite of the fact that cupric 

hydroxide is more insoluble than ferrous hydroxide.

CARBONATES.

PREVIOUS WORK.

In 1853 Debus 2 published some quantitative experiments on the 

fractional precipitation of barium and calcium carbonate. Although 

some of his conclusions are faulty, an examination of his results 

shows that calcium competing with barium is precipitated to a 

greater extent than barium by a soluble carbonate. No other experiments on the fractional precipitation of carbonates have come-to my 

attention except those reported in 1882 by Mills and Bicket 3 on the 

fractional precipitation of nickel and manganous sulphates and nickel 

and cadmium sulphates by a slight deficiency of sodium carbonate. 

Their experiments clearly show that cadmium is precipitated to a 

greater extent than nickel and that nickel1 is precipitated to a greater 

extent than manganese under equal terms of competition. Their 

experiments were made in the cold and no time was allowed for any 

possible readjustment. Other experiments by them on the separate 

precipitability of these salts show that rather complex precipitates 

are obtained.

1 Curtman, L. J., and St. John, A. D., Am. Chcm. Soc. Jour., vol. 34, p. 1679, 1912.

2 Debus, H., Ueber chemische Verwandtschaft: Liebig's Annalen, vol. 85, p. 124, 1853. 

s Mills, E. J., and Bicket, J. H., Researches on chemical equivalence: Philos. Mag., 5th ser., vol. 13, 

pp. 169, 177, 1882.
SOLUBILITY OF CARBONATES. 29

T. Sterry Hunt l states that a gypsum solution loses no calcium to 

crystalline dolomite or crystalline carbonate of magnesium; and that 

even in the presence of carbon dioxide very little magnesium is taken 

into solution; but with hydrated magnesium carbonate or hydrate of 

magnesium there is considerably more magnesium dissolved.

To determine the behavior of certain heavy metals with an excess 

of carbon dioxide at ordinary temperatures Raikow studied the action 

of that gas on their hydroxides. 2 He found that basic carbonates 

were formed with glucinum and yttrium, that acid carbonates were 

formed with magnesium, thallium, nickel, and cobalt and slightly with 

manganese and ferrous iron, and that all the other carbonates were 

normal. His results do not include any statement of the solubilityor insolubility of the carbonates so formed, although they indicate 

that with an excess of carbon dioxide only silver, cadmium, mercurous, lead, and copper salts among the common elements form insoluble precipitates to any great degree of completeness.

SOLUBILITY OF CARBONATES.

The preliminary data available in regard to the solubility of the 

carbonates are as follows:

Solubility of carbonates in pure water.

PbC03 . ....................................

FeCO3 . .....................................

ZnCO3 . .....................................

BaCOs-..-- ................................

SrCOa-. ....................................

MnCO3 .. ...................................

NiCOs.... .............:....................

MgC03 . ....................................

ThCOs.. ........................... ........

Li2C03 .....................................

Na2C03 . ....................................

Temperature.

°C.

18

15

15

16

16

15

25

100

15 25'

100

15

15

15

15

15

15

15

15

Gram equivalents 

per liter.

1.6 X10-«.........

6.2 XIO-5 ... ......

1.7 X10-4.........

1.9 XIO-1 ... ......

2.6 XHH.........

3.4 X10-4 -........

2.86X10-4 -. .......

3.5GX10-4 -........

3.0 X10-4......... 

3. 06X10-4 . ........

3.80X10-».........

3.2 X10-4... ......

4.6 XIO-4 -..-.....

4 7 V 1 0-1

2.4 XIO-3 .........

1 7 Y10~!

3.7 XIO-1 .........

2.6................

Authority.

Do.

T. Schloesing.c

Do.

J. Kendall.d

Do.

Do.

Do.

Do.

Do.

« Uebcr die Loslichkeit einiger Bleiverbindungen in Wasser: Arbeiten aus dem Kaiserl. Gesundheitsamt, Bd. 26, Heft 3, Berlin (Julius Springer), 1907.

b Recherches exp<5rimentales sur la solubilit6 des carbonates et des bicarbonates, Thesis, Geneva, 1897. 

c Compt. Rend., vol. 74, p. 1555. 1872. 

d Philos. Mag., vol. 23, p. 958, 1912. 

e Abegg, R., Handbuch dor anorganischen Chemie, Bd. 2, Abt. 2, p. 157,1905.

The uncertainty of some of these data is illustrated by the figures 

for silver carbonate, which does not appear in the table. Cox 3

1 Chemical and geological essays, p. 107,1878.

2 Raikow, P. W., Weitere Untersuchungen fiber die Einwirkung der Kohlensaure auf die Hydrate 

der Metalle: Chem. Zeitung, vol. 31, p. 55,1907.

8 Cox, A. J., Ueber die Loslichkeitsverhaltnisse einiger schwerloslicher Silbersalze: Zeitschr. physikal. 

Chemie, vol. 46, p. 11,1903.
30. FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

gives 17 parts per million, whereas Moissan 1 gives 30 at 15° C. 

Moissan states that this is increased to 1,040 parts per million when 

the water is saturated with carbon dioxide. Discrepancies in many 

of the determinations are probably due to the fact that the acidity of 

the solutions was not exactly controlled. Even the slight amount 

of carbon dioxide in the air increases the solubility of carbonates 

enormously.

EFFECT OF HYDROLYSIS.

Owing to hydrolysis hydroxides, not carbonates, are precipitated 

with aluminum, ferric, and chromic salts. " Basic carbonates" are 

precipitated under ordinary circumstances with zinc, nickel, cobalt, 

copper, and mercuric salts, and boiling removes the carbon dioxide 

from these salts to a greater or less extent. When hydrolysis is not 

so pronounced or the hydroxide is more soluble, normal carbonates 

arc formed as with calcium, manganese, ferrous iron, cadmium, 

silver, lead, and mercurous salts; these are unchanged on boiling, 

but the carbonates of lead, mercury, and silver are changed to basic 

carbonates in time. Lead carbonate suspensions become alkaline at 

70° according to Abegg.

For carbonates the chief hydrolytic reaction would be

CO3 - - + HOH = HCO3 - + OHand to a less extent

HCO3 - + HOH = H2C03 + OHThe degree of hydrolysis of sodium carbonate at 24.1° is shown below 

for the concentrations stated, the degree expressing the percentage 

of total sodium present in the form of the hydroxide.

Hydrolysis of sodium carbonate at 24.1° C.

Moles Na2CO3 per liter. ........ 0.1900 0. 0940 0. 0477 0.0238

Per cent hydrolysis............. 2.12 3.17 4.87 2 7.10

The hydrolysis of sodium carbonate may be increased by boiling 

the solution and the removal of carbon dioxide as follows, starting 

with a normal solution: 3

Hydrolysis of sodium carbonate produced by boiling.

Time of boiling (hours) ...........

Remaining Na2C03 (per cent)

. . . 1. 25

... 96.7

. . . 3. 3

3.5

94.6

5.4

6.0

92.7

7.3

8.5

91.4

8.6

23.3

87.0

13.0

38.0

83.8

17.2

1 Moissan, H., Traits" de chimie min&ale, vol. 5, p. 5,80,1904-1906.

2 Shields, John, Ueber Hydrolyse in wasserigen Salzlosungen: Zeitschr. physikal. Chemie, vol. 12, p. 167, 

1893.

s Kiister, F. W., and Griiters, Max, Ueber den Zerfall von geloster Soda in Kohlendioxyd und Natriumhydroxyd: Deutsche chem. Gesell. Ber., vol. 36, p. 748,1903.
FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS. 31 

EXPERIMENTS WITH CARBONATES AT ORDINARY TEMPERATURE.

In making fractional precipitations of carbonates two metallic salts 

were precipitated by less than their chemical equivalent of soluble 

carbonate and the ratio of the metals in the precipitate determined.

The first precipitations were made in flasks which stood at the 

temperature of the laboratory and were shaken from time to time. 

The analyses were carried out by difference that is, by analyzing an 

aliquot part of the solution after filtering off the precipitates and subtracting the quantities found from the quantities taken. The general 

scheme of this ionization and reaction in solution may be illustrated 

by a mixture of zinc sulphate, calcium sulphate, and sodium carbonate, as follows :

(Zn,Ca)S04 = (Zn,

Na2C03 = C03- - + 2Na+ (4)

Since, in the fairJy dilute solutions used, the soluble salt, in this 

experiment sodium sulphate, plays an insignificant part, the scheme 

may be simplified even further, as follows:

(Zn,Ca) ++ +00," = (Zn,Ca)C03 (5)

Information can be gained in the first place by considering what 

will happen when equivalent quantities of the two metals are present, 

with just enough carbonate for one, and afterward by noting whether 

the result so found is changed for other proportions of the metals.

From equation (5) it is evident that the equilibrium to be determined may be obtained by precipitating both the metallic salts with 

a soluble carbonate or by treating a carbonate of one metal with a 

soluble salt of the other metal. The first method was chiefly 'used, 

and the results are presented below.

In these results the radicles and metals are expressed in milligram 

equivalents, an equivalent being equal in milligrams to the molecular 

weight of the metal or radicle reduced to a univalent basis. The 

solution was made up to 1 liter in every experiment, or was calculated 

to that basis.
32 FRACTIONAL PRECIPITATION OF OKE-FOEMING COMPOUNDS.

Preliminary results of experiments with carbonates at 25° C., with equivalent quantities of two 'metals present and carbonate enough for one.

Quantity taken (milligram equivalents). Quantity precipitated (milligram equivalents).

Experiment 4.

........................... 8.0

........................... 8.0

Na2C03 ........................... 8.0

Time, 17 days.

Experiment 5.

Zinc............................... 6.52

Calcium......................... None.

............................. 20.0

ZnSOi-......................'....... 20.0

Na2C03 ........................... 20.0

Time, 19 days.

Experiment 6.

Cadmium......................... 15. 5

Zinc.............................. 4.5

FeS04 ............................ 10.0

CaS04 ............................. 10.0

Na2C03 ........................... 10.0

Time, 3 days.

Iron.............................. 6.25

Calcium......................... None.

Experiment 7.

OuS04 ............................. 10.0'

FeSO4 ............................ 10.0

Copper........................... 5. 30

Iron...........:.................. 5. 35

Na2C03 ........................... 10.0

Time, 2 days.

Experiment 8.

CuS04 ............................. 8.0

ZnS04 ............................. 8.0

Na2C03 ........................... 8.0

Time, 26 days.

Copper........................... 6. 84

Zinc.............................. .76

Frescott and Johnson 1 state that calcium carbonate has no action 

on zinc salts at ordinary temperature. Since the above experiment 

shows a considerable precipitation of the zinc to be exact, 6.52 milligram equivalents in 8 or 81.5 per cent in 17 days it appears that the 

duration of the experiment is a factor to be considered. This was 

confirmed by further experiments, in which finely powdered calcite 

was introduced into zinc sulphate instead of precipitating the two 

salts with sodium carbonate.

In one experiment (No. 12) 8 milligram equivalents of calcite were 

introduced into 8 milligram equivalents of zinc sulphate in a liter. 

The flask was shaken occasionally at room temperature. After 2 

days 16.5 per cent of the zinc was precipitated, and after 24 days 

81.2 per cent was precipitated.

In another experiment (No. 9) precipitated calcium carbonate, 

which had been previously dried and heated, was digested at 20°-25° 

for 3 hours with dilute zinc sulphate. No zinc was precipitated.

1 Prescott, A. B., and Johnson, 0. C., Qualitative chemical analysis, 5th ed., p. 179,1903.
EXPERIMENTS WITH CARBONATES AT1 ORDINARY TEMPERATURE. 33

These results show that hasty chemical experiments may lead to 

greatly mistaken conclusions in geologic investigations. At ordinary 

temperature several days are required to show the effect; and, in fact, 

it is not certain that equilibrium was attained in these, experiments,, 

but it is certain that calcium carbonate will precipitate considerable 

zinc in time, even at ordinary temperatures. On the other hand, too 

long continuation of experiments in glass.vessels introduces an uncertainty on account of the slow production of basic salts due to the 

action of the glass.

The fact that in the first experiment with zinc sodium carbonate 

was added, whereas in those just mentioned calcite was used, necessitates a consideration of the varieties of calcium carbonate. It is 

generally true that a freshly precipitated substance is more soluble 

than the same substance after the lapse of time. With some substances this is due to the size of the particles, with others to the 

crystalline form, and these facts must be considered in precipitations 

with solids. As regards calcium carbonate, W. Meigen 1 has shown 

that when calcium chloride solutions are precipitated by sodium 

carbonate aragonite separates, in globules at low temperatures, in 

needles at high temperatures, and that on standing under some conditions this aragonite passes into calcite. Since aragonite is more 

soluble than calcite, it would be expected to have a slightly greater 

precipitating power on solutions of metallic salts than calcite. Meigen 

has shown that this is the case. This study has been continued by 

his students. 2

Similar conclusions can be drawn from my own experiments with 

freshly precipitated calcium carbonate. As a variation of the fractionationin one experiment (No. 13), the calcium sulphate was first 

precipitated with its equivalent of sodium carbonate, then, without 

filtration, zinc sulphate was added, and in 2 hours 17.6 per cent of the 

zinc was precipitated. With the calcium carbonate which had been 

previously dried and heated no zinc was apparently precipitated, 

even in 3 hours, but after 24 days the zinc was practically all 

precipitated.

The conclusion from all the experiments with zinc and calcium at 

ordinary Jbemperature is that calcium carbonate in any form will ultimately precipitate a considerable amount of zinc from zinc sulphate 

solutions but that the metathesis takes tune to occur. Moreover, it 

occurs slowly with well-dried calcium carbonate, more quickly with 

aragonite than with calcite, and most quickly with freshly precipitated calcium carbonate.

> Naturf. Gesell. Freiburg im Breisgau Ber., Bd. 30, p. 64,1903.

2 Gassner, L., Weitere Beitriige zur Kenntniss des kohlensauren Kalks: Inaug. Diss., Freiburg 1m 

Breisgau, 1900.

95921° Bull. 609 15 3
34 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

These results immediately suggest new experiments concerning the 

action of calcium carbonate on salts of iron, aluminum, chromium, 

cobalt, nickel, and manganese. It is held in qualitative analysis that 

a suspension of freshly precipitated calcium carbonate will precipitate 

ferric, aluminum, and chromic salts in the cold, but have no action on 

ferrous, cobalt, nickel, manganese, and zinc salts. Chromic salts are 

said to precipitate more slowly than aluminum or ferric salts.1

According to Meigen's researches,2 a hundredfold excess of calcite 

precipitated a 0.045 normal solution of manganese to the extent of 

2.9 per cent in 72 hours at 17 P, and aragonite precipitated 70.2 per 

cent under the same conditions.

In experiment 21 calcite was tried with nickel sulphate; 5.10 milligram equivalents of nickel sulphate and 4 milligram equivalents of 

sodium sulphate in 500 cubic centimeters of water were shaken occasionally with 4 milligram equivalents of powdered calcite for 17 days. 

The precipitate then showed 0.36 milligram equivalents of nickel and 

3.89 milligram equivalents of calcium. Hence the precipitation of 

the nickel by calcite after 17 days was slight, only about 7 per cent 

of the whole quantity, but experiment 87 shows that time is an important factor in this action and that eventually nickel is very largely 

if not completely precipitated.

Magnesium carbonate, like calcium carbonate, is known to exist in 

several forms which differ in solubility and which must therefore also 

differ in their capacity for precipitating other carbonates. The most 

soluble form is the trihydrate, MgC03 .3H2O. A liter of pure water 

saturated with this salt contains 0.23 gram of magnesium and 0.42 

gram carbon dioxide at 20°, according to Auerbach,3 in the absence 

of any excess of carbon dioxide. I have found its solubility under 

atmospheric conditions, that is, its solubility in water in equilibrium

with air containing the normal amount of carbon dioxide, to be somewhat greater, namely, 0.36 gram magnesium per liter, total carbon 

dioxide 1.01 gram per liter. Amorphous magnesite was found to be 

much less soluble than this. No experiments were made with crystallized magnesite, which is rather difficult to obtain.

Other fractional experiments were made, beginning with experiment 22, the results of which, together with the preceding experiments, 

are shown in the following table, all results corresponding to a volume 

of 1 liter. In most experiments 8 milligram equivalents of sodium 

carbonate were used, but a few variations were made, as follows: 

Experiment 32, 10.28; experiment 34, 8.80; experiment 39, 8.30; 

experiments 71 and 74, 16 milligram equivalents of sodium carbonate.

1 Prescott, A. B., and Johnson, O. C., Qualitative chemical analysis, 5th cd., p. 154,1903.

2 Op. cit., p. 83.

a Auerbach, F., Zeitschr. Elektrochemie, vol. 10, p. 161,1904.
EXPERIMENTS WITH CARBONATES AT ORDINARY TEMPERATURE. 35 

Summary of fractional precipitations of carbonates at ordinary temperature.

Experiment.

4............

5............

6............

7............

8............

22...........

23...........

24...........

25...........

26...........

27...........

28...........

29...........

30...........

31...........

32...........

33...........

34...........

39...........

41...........

42...........

55...........

60. ..........

64...........

71...........

74...........

78...........

87...........

90...........

91...........

92...........

93...........

Quantity of mixture taken (milligram 

equivalents per liter).

20CdS04 ..........

10FeSO4..........

8Pb(N03) 2 .......

8Pb(N03 )2........

8Zn(NO3)2........

8AgN03 ..........

8AgN03 .. ........

8ZnS04 ...........

8AgN03 .... ......

8FeSO4 ..... ......

8FeS04 ...........

8Cd(N03 ) 2........

8MnS04..........

8NiSO4 ...........

8NiS04 .. .........

SAgNOj..........

8AgN03 .. ........

8HgN03 ..........

8CuS04 ...........

8NiS04 ... ........

16ZnSO4 .. ........

16MnS04 .........

8ZnS04 ...........

8NiSO4 ...........

20 Fe(NO3 ) 2. ......

20Fe(N03 )2 .......

20FeS04 ..........

20ZnS04 ..........

8Ca80t ... ........

20ZnSO4 ..........

10FeSO4 ..........

8ZnSO4 ...........

SAgNOs..........

8Cd(NO3 )2........

8Pb(NO3 ) 2........

8HgNO3 ..........

8Cd(N03 ) 2........

8NiS04 ...........

8Zn(NO3 )»........

8NiS04 ...........

8Zn(NOa)s........

8HgN08 ..........

8CaS04 ...........

14.54 MgSO4 .......

8MnS04 . .........

8MnS04 . .........

8CaSO4 ...........

8MgSO4 ..........

8Pb(N03 )2 .......

8CdS04 ...........

8AgN03 ..........

16MnSO4 .........

16CaS04 ..........

8CdS04 ...........

8CaSO4 ...........

20Pb(N03 )2.......

20Cu(NO3 ) 2.......

20CdSO4 ..........

20 FeS04 . .........

Duration 

of experiment 

(clays).

17 

19 

3 

2 

26 

4 

6 

i 6 

24 

4 

10 

8 

1 

1 

11 

2 

5 

9

7

7 

8 

6 

8 

6 fa)

b ) 

15 

800 

1 

1 

1 

1

Quantity of the metal found in precipitate (milligram equivalents).

15.5 cadmium.....

8.00 lead.. .........

7.96 lead.. .........

2.28 silver.........

S.OOzinc...........

4.20 iron...........

3.54 cadmium ..... 

6.56 manganese. . . .

4.04 nickel.........

3.98 nickel.........

6.80 silver... '......

11.96 zinc.. ........

14.80 manganese. . . 

1.26 zinc...........

2.40iron...........

10.65 iron..........

Tr. calcium. 

4.5 zinc. 

Tr. calcium. 

5.35 iron. 

0.76 zinc. 

0.08 silver. 

0.42 cadmium. 

Tr. zinc. 

7.98 mercury. 

7. 18 cadmium. 

1.50 nickel. 

3.96 zinc. 

3.06 nickel. 

2.64 zinc. 

7.36 mercury. 

1.64 calcium. 

Tr. magnesium. 

3.80 manganese. 

5.76 manganese.

1.56 calcium. 

1.12 magnesium. 

0.75 lead. 

0.20 cadmium. 

2.96 silver. 

8.96 manganese. 

5.70 calcium. 

2.34 cadmium. 

1.80 calcium. 

19.0 lead. 

10.20 copper. 

11.6 cadmium. 

4.2 iron.

o 18 hours. & 1 hour.

The results show that very rarely was a complete separation 

effected in the fractionation in the time allowed; in most of the experiments both metals were found in the precipitate and filtrate. 

The separations are far less sharp than those obtained with sulphides. 

The precipitation series deduced from these experiments is shown in 

the table below. The numbers in the third column indicate the 

ratio, in equivalents, of the predominating metal to the one in smaller 

amount in the precipitate. Not all the experiments have been made 

which could be made; conversely some of the orienting experiments 

were later seen to show little except that the two metals were widely 

separated in the series.
36 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

Precipitation series deduced from the experiments on fractional precipitation of carbonates 

and relative precipitation of the two metals.

Experiment.

55....................................................

31...................................................

25...................................................

59...................................................

90............................:.....................-

23...................................................

24...................................................

22...................................................

7, 91 ................................................

8....................................................

60...................................................

78...................................................

92...................................................

67...................................................

26...................................................

28. ................................................. -

27...................................................

4. ...... .............................................

71...................................................

93............................................:......

7....................................................

30...................................................

29...................................................

6......................................:.............

34...................................................

33...................................................

87...................................................

64...................................................

32, 7-1. ..............................................

41...................................................

42...................................................

Major constituent 

of precipitate.

...:.do.. ...........

.....do.............

.....do.............

.....do.............

.....do.............

.....do.............

.....do.............

.....do.............

.....do.............

Zinc...............

.....do.............

.....do.............

.....do.............

.....do.............

Iron...............

.....do.............

.....do.............

.....do.............

Nickel.............

.....do.............

.....do.............

.....do.............

Silver.............

Minor constituent 

of precipitate.

Nickel.............

.....do............

Nickel.............

.....do.............

Ratio of 

major to 

minor constituent.

9 1

^ ^

6.5

19.0

100.0

1.0

9 A

36.0

1.9

4.0

2.8

 ji n

2.5

3.3

6.5

i ^

3.7

1.0

1.0

1.4

6.3

1.1

1.1

4.0

4.4

1.4

Incidental to the determination of relative solubility the experiments showed that metallic silver was precipitated with ferrous and 

mangarious salts and there was also a very slight darkening of the 

precipitate with nickelous salt. The exact position of silver is therefore in some doubt, but it is certainly below zinc.

The positions of manganese and nickel are inconsistent with the

recorded solubilities. Either the solubility of manganese carbonate 

should be less than that of calcium carbonate or there was oxidation 

of the manganese salt in experiment 70. Nickel shows a similar 

discrepancy.

The experiments further throw doubt on the solubility of ferrous 

carbonate given on page 29, although it must be said that the fractionations with ferrous sulphate were not very satisfactory. The 

difficulty was partly due to the ready oxidizibility of ferrous sulphate 

and partly to the apparent tendency of iron to enter into double 

compounds containing an atom of iron to one of the other metal. 

There is also a tendency toward the formation of bicarbonates, which 

is considered on page 37.

The heats of precipitation for the metallic carbonates are so small 

that no apparent regularity exists in the relation between the heats 

and the positions in the precipitation series.
FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS. 37 

RESULTS OF EXPERIMENTS WITH BICARBONATES.

The effect of basic carbonates on the fractionations and the elimination of this effect will now be considered. If in any precipitation of 

carbonate a hydroxide of one of the metals possesses such a small 

solubility that the hydrolysis of the carbonate would result in exceeding the solubility product of the hydroxide, more or less "basic salt" 

will be formed. These precipitates are ordinarily considered to be 

indefinite mixtures of hydroxide and carbonate, but it seems highly 

probable that either a pure carbonate, a hydroxide, or a definitebasic salt would result if time were allowed for the attainment of 

equilibrium, for it has been shown that many days are often required 

for the attainment of a stable state at ordinary temperatures. This 

is another instance where the duration of the experiment determines 

the final product.

It is possible that a basic carbonate of one metal might form at the 

same time as a bicarbonate of another metal. There is some evidence 

of this in experiments 87, 92, and 93. In experiment 92 the mother 

liquor showed 1.1 milligram equivalents of HC03~, and in experiment 

93 it showed 5.3 milligram equivalents of HC03 ~ per liter. 

In order to ensure that carbonates alone were being precipitated 

the experiments might have been carried out in the presence of a 

certain excess of carbon dioxide. As a much simpler means to the 

same end, however, a number of experiments were made with sodium 

bicarbonate as precipitating agent. The only difficulty encountered 

in this procedure was the rather large solubility of the bicarbonates 

of certain metals. The experiments were as follows:

Fractional precipitations ivith sodium bicarbonate.

Quantity taken (milligram equivalents). Quantity precipitated (milligram equivalents).

Experiment 35.

CllS04 . ......................... 8.0

ZnS04 .......................... 8.0

....................... 8.0

Time, 9 days.

Experiment 36.

Copper........................... 6. 08

Zinc............................ 5. 84

CdS04 .......................... 8.0

ZnS04 . ......................... 8.0

NaHCOj........................ 8.0

Time, 3 days.

Experiment 37.

Cadmium....................... 5,46

Zinc............................ 1.60

........................ 8.0

.............:.......... s.o

NaHC03 ........................ 8.0

Time, 20 hours.

Mercury........................ 7. 54

Zinc.......................... None.
38 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

Fractional precipitations with sodium bicarbonate Continued.

Quantity taken (milligram equivalents) Quantity precipitated (milligram equivalents). 

Experiment 38.

Pb(N03)2 ......................... 8.0 Lead............................. 7.82

Cu(N03)2 ........................ 8.0 Copper........................... 0.16

NaHC03 ........................ 8.0

Time, 3 days.

Experiment 44.

ZnS04 .......................... 8.1 Zinc............................ 0.70

'FeS04 ........................... 8.0 Iron............................ 4.36

NaHC03 ........................ 8.0

Time, 24 hours.

Experiment 45.

ZnS04 .......................... 8.1 Zinc............................ 2.14

NiS04 .......................... 8.0 Nickel.......................... 0.82

NaHC03 ........................ 9.0

Time, 8 days.

Experiment 46.

ZnS04 .......................... 8.0 Zinc. ........................... 1.06

AgN08 ........................... 8.0 Silver............................ 2.90

NaHC08 ......................... 8.0

Time, 8 days.

Experiment 47.

AgN08 .......................... 8.0 Silver............................ 3.70

MnS04 .......................... 8.0 Manganese. ..................... 4.32

NaHC03 ........................ 8.0

Time, 11 days.

Experiment 48.

ZnS04 .......................... 40.0 Zinc............................ 12.1

FeS04 ........................... 40.0 Iron. ........................... 21.2

NaHC03 ........................ 40.0

Time, 20 hours.

Experiment 51.

FeS04 ........................... 32.0. Iron............................ 9.16

MnSO4 .......................... 32.0 Manganese...................... 21. 92

NaHC03 ........................ 22.0

Time, 1 day.

Experiment 52.

CuS04 .......................... 8. 0 Copper........................... 5. 64

CdS04 ........................... 8.0 Cadmium. ...................... 0.40

NaHCOg........................ 8.0

Time, 5 days.

Experiment 53.

HgN03 .......................... 8. 0 Mercury................... Nearly all.

Pb(N03)2 .......................... 8.0 Lead....................... Trace.

NaHC03 ........................ 8.0

Time, 5 days.
RESULTS OF EXPERIMENTS WITH BICABBONATES. 39

Fractional precipitations toith sodium bicarbonate Continued.

Quantity taken (milligram equivalents). Quantity precipitated (milligram equivalents).

Experiment 54.

HgN03 8.0

Cd(N03 ) 2 ........................ 8.0

NaHCOg. ................. ---.- - 8.0

Time, 4 days.

Mercury..................... Nearly all.

Cadmium . Trace.

Experiment 57.

MnS04 . ........................

NaHC08 . ......................

Time, 5 days.

8.0 

8. 0 

8.0

Cadmium....................... 7. 60

Manganese...................... 5. 96

Experiment 72.

16.0

16.0

16.0

Time, 18 hours.

Zinc............................ 6.84

Manganese...................... 10. 24

Experiment 88.

NiS04 .......................... 16.0

CaS04 .......................... 16.0

NaHCOg........................ 16.0

Time, 5 days.

Nickel.......................... 1.88

Calcium......................;.. 0. 20

Summary of results of fractional precipitation with sodium bicarbonate at 15°-20° C., 

arranged in the order of solubility deduced from the experiments.

Experiment.

53..................................................

54..................................:...............

37..................................................

38..................................................

52..................................................

50..................................................

35..................................................

57..................................................

49...................................................

36..................................................

89..................................................

51..................................................

72..................................................

46. .................................................

44. 48. ..............................................

45..................................................

88...................................................

Major constituent 

of precipitate.

Mercury.. ........

.....do.............

.....do.............

.....do.............

Cadmium.........

.....do.............

.....do......:......

.....do.............

Manganese. ........

.....do..............

Silver.............

Zinc...............

Nickel.............

Minor constituent 

of precipitate.

Lead..............

Zinc...............

Ferrous iron ......

Zinc.. ... ......

Manganese..... ...

Zinc...............

Silver.............

Ferrous iron .......

Zinc...............

.....do.............

.....do.............

Nickel............

Ratio of 

major to 

minor constituent.

(«)

M

49 0

14.0

1.6

1.0

1.3

6.4

3.4

92.0

2.5

1.5

2.7

1.7

2.6

9.4

Chiefly mercury.

From these results it appears that when bicarbonates are used instead of carbonates the principal changes are that ferrous iron, zinc, 

and nickel fall below silver in the series.
40 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS. 

EXPERIMENTS WITH CARBONATES AT HIGHER TEMPERATURES'

At 100° C. bicarbonates are no longer stable at low pressures, although the last traces of carbon dioxide are held very tenaciously by 

solutions containing calcium or magnesium salts. Under high pressures of carbon dioxide and at elevated temperatures the bicarbonates 

would no doubt persist. The compounds stable under these conditions, however, and their solubilities remain to be determined. A 

few solubilities have been measured in solutions saturated with carbon 

dioxide at a pressure of a few atmospheres, but mostly at ordinary 

temperatures. The crystallization of dolomite at 150° C. has been 

noted. Of course many of the metallic carbonates become basic as 

the temperature rises, as has already been stated, and some lose carbon dioxide on boiling. Hence in most of the experiments the 

precipitation was made cold and the flask and contents then heated 

to 60°-80° C.

In experiment 16 a solution containing 5.43 milligram equivalents 

of cadmium sulphate and 5.43 milligram equivalents of zinc sulphate 

was precipitated with 5.43 milligram equivalents of sodium carbonate 

 in a volume of 250 cubic centimeters. After boiling three hours the 

precipitate contained 4.76 milligram equivalents of zinc and 0.60 

milligram equivalents of cadmium.

In experiment 18 a solution containing 8 milligram equivalents of 

nickel carbonate was boiled one hour with 8 milligram equivalents of 

calcium sulphate. No nickel dissolved.

In experiment 20 a solution containing 8.81 milligram equivalents of 

cadmium sulphate and 8.81 of nickel carbonate was boiled five hours 

in the presence of 8.81 milligram equivalents of sodium sulphate. A 

small amount of cadmium was precipitated.

In experiment 66 the attempted fractionation of manganese sulphate and silver nitrate resulted in a reduction of the silver salt and 

oxidation of the manganese.

In experiment 70 the precipitate formed from the sulphates of manganese and calcium turned brown in about a quarter of an hour, showing a slight oxidation of the manganese.

Further results follow.
EXPERIMENTS WITH CARBONATES AT HIGHER TEMPERATURES. 41

Fractional precipitations with sodium carbonate at 60°-80° C.

Quantity taken (milligram equivalents). Quantity precipitated (milligram equivalents).

Experiment 78.

Hg(N03) 2 ........................ 16.0 Mercury.......................... 8.88

Pb(NO3) 2 ......................... 16.0 Lead............................ 5.2

;Na2C03 ...... ....... ............. 16.0

1 hour at 60° C.

Experiment 79.

Pb(N03) 2 ......................... 16.0 Lead............................ 10.24

Cu(N03) 2 .... .................... 16.0 Copper........................... 6.68

Na2C03 ....... ................... 16.0

1 hour at 80° C.

Experiment 80.

Ag(N03 ) 2 ........................ 12.0 Silver............................ 11.00

MgS04 .......................... 12.0 Magnesium...................... .40

Na2C03 .......................... 12.0

1 hour at 70° C.

Experiment 82.

ZnS04 ............................ 16.0 Zinc............................ 13.80

FeS04 .................... '...-... 16.0 Iron.............................. 4.72

Na2C03 . ......................... 16.0

1 hour at 70° C.

. Experiment 83.

CdS04 - ........................... 16. 0 Cadmium. . ....................... 14. 32

FeS04 ........................... 16.0 Iron.............................. 2.04

Na2C03 .......................... 16.0

1 hour at 70° C.

Experiment 84.

FeS04 ........................... 16.0 Iron.............................. 0.68

NiS04 ............................ 16.0 Nickel.......................... 10.0

Na2C03 . ......................... 18.0

1 hour at 70° C.

Experiment 85.

CaS04 ............. ............... 12.0 Calcium................ Undetermined.

FeS04 ........................... 12.0 Iron.............................. 8.44

Na2C03 . ......................... 12.0

1 hour at 70° C.

Experiment 86.

CuS04 ......... ................... 16.0 Copper........................... 9.08

FeS04 ........ ...... ............. 16.0 Iron.............................. 6.72

............. 16.0

1 hour at 70° C.
42 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

These results are collected in the following table:

Summary of results of fractional precipitation with sodium carbonate at 60°-SO° C., arranged in the order of solubility deduced from the experiments.

Experiment.

78..........................................

76..........................................

59, 62, 79.. ..................................

15..........................................

86..........................................

17..........................................

77..........................................

82..........................................

67..........................................

83..........................................

84..........................................

73..........................................

65..........................................

94..........................................

85..........................................

95..........................................

80...............................:..........

Major constituent 

of precipitate.

.....do.............

Zinc...............

.....do.............

.....do.............

.....do.............

.....do.............

.....do.............

Minor constituent 

of precipitate.

.....do............

.....do.............

Silver.............

Ratio of major to 

minor constituent.

1.7

4.5

1.5

3.6

1.4

2.1

1.1

2.9

2.7

7.0

15.0

10.0

17.0

1.0

1.5

27.0

The results at higher temperatures show only minor changes from 

the series at ordinary temperatures. Zinc appears above cadmium, 

probably owing to the formation of oxide. Nickel appears above 

iron, but as there is some doubt about their relative positions in the 

series at ordinary temperatures, this change is not to be emphasized. 

This shows that the experiments at ordinary temperatures yield conclusions of fundamental significance.

In experiment 86 the ferrous iron was partly oxidized, presumably 

at the expense of cupric sulphate. Of the iron in the filtrate 24 per 

cent was in the ferric condition. The experiments with ferrous salts 

were all made in a special precipitating flask in an atmosphere of 

hydrogen, so that oxidation by air was sufficiently avoided.

SUMMARY OF RESULTS OF EXPERIMENTS WITH CARBONATES.

1. Calcite introduced into a dilute solution of zinc sulphate precipitates the zinc slowly, requiring two or three weeks to precipitate 

most of the zinc.

2. Freshly precipitated calcium carbonate acts more rapidly in 

precipitating metallic salts than any other form, and aragonite and 

calcite follow in the order named.

3. The precipitating action of calcite on manganese,, nickel, and 

silver salts is less marked than on zinc, cadmium, copper, lead, and 

mercury salts.

4. When a mixture containing two equivalents of two metallic 

salts, one equivalent of each, is precipitated by one equivalent of a 

soluble carbonate, a fractionation occurs, and the extent of this 

fractionation has been determined approximately for a large number 

of substances.
SILICATES. 43

5. The series obtained, beginning with the metals precipitated to 

the greatest extent, with sodium carbonate at ordinary temperature 

is mercury, lead, copper, cadmium, zinc, ferrous iron, nickel, manganese, silver, calcium, magnesium.

6. The series,obtained with sodium bicarbonate at ordinary temperature is mercury, lead, copper, cadmium, manganese, silver, ferrous iron, zinc, calcium, magnesium.

7. The series obtained in warm solutions with sodium carbonate is 

mercury, lead, copper, zinc, cadmium, nickel, iron, calcium, silver, 

magnesium.

8. Silver salts are partly reduced by ferrous and manganous salts; 

possibly also by nickelous salts, this effect becoming more marked as 

the alkalinity increases.

9. In the presence of bicarbonates zinc shows relatively greater 

solubility than cadmium, and nickel is likewise relatively more 

soluble than when precipitated with normal carbonate.

10. With rising temperature the precipitates from cadmium, iron, 

and silver salts appear to increase in solubility relatively faster than 

the others.

11. The presence of cupric sulphate greatly intensifies the tendency

of ferrous salts to oxidize. &

SILICATES.

PREVIOUS WORK.

Since many of the natural silicates have been formed under the 

condition of equilibrium with a magma it is obvious that they will 

not be in equilibrium with aqueous solutions. In order to complete 

the study of fractional precipitation, however, it appeared desirable 

to see if there would be any selective action when solutions containing 

more than one metallic salt were precipitated by a soluble silicate. 

Jordis has already studied the precipitation of single metallic silicates 

from soluble salts l with reference to their composition. His experiments with ferric salts,2 however, are probably complicated by the 

phenomena of hydrolysis, but he suggests that the behavior of 

colloids must be considered. :

EXPERIMENTS WITH SILICATES AT ORDINARY TEMPERATURE.

My own experiments have consisted of attempts at fractional 

precipitations of metallic salt solutions by amounts of soluble silicates 

insufficient for both the metallic constituents. Although it was 

expected that the sodium silicate, being hydrolyzed, would behave like

1 Jordis, Eduard, Zur Chemie der Silikatc: Zeitschr. angew. Chemie, vol. 19, p. 1697, 1906; Jour, prakt. 

Chemie, vol. 77, pp: 226-261,1906; idem, vol. 81, p. 303, 1910.

* Jordis, Eduard, and Lincke, P., Beitriige zur Kenntnis der Mctall-Silicate: Jour, prakt. Chemie, vol. 

SI, p. 289, 1910.
44 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

caustic soda, it seemed desirable to determine whether the silicic acid 

would participate in the reaction. After the results with the hydroxides, it was not surprising to find that in general one metal was precipitated to a greater extent than the other. Moreover, most of the 

silica was carried down with the precipitate. The results are shown 

in the following table, where the concentrations are expressed in 

millimoles per liter, one millimole being a thousandth of the molecular weight taken in grams. A diluted water-glass solution was used 

as the soluble silicate. The ratio of Si02 to Na20 in it was 3.09 to 

1.00. It was free from other metals except for a trace of iron and 

alumina. After precipitation the solutions stood at about 30° C. 

for the periods noted in the table. The last column in the table 

shows the silica not precipitated by the metallic salts present. Sulphates were used except with lead and calcium salts, where nitrates 

were employed.

Results of the fractional precipitation of silicates, arranged in the order of the solubility 

deduced from the experiments.

Experiment.

114

115

116

117

118

119

120

121

122

123

o!24

o!25

Time

(days).

97

91

91

107

97

31

102

31

95

97

95

95

Quantities precipitated (millimoles).

4.8 copper, 0.2 lead, 11.2 SiOs ..............................................

6.7 copper, 0.7 zinc, 12.8 SiOz. ..............................................

5.1 zinc, 0.1 silver 11.9 SiO2. ................................................

5.1 zinc, 2.6 nickel, 10.5 Si02 . ..............................................

Very little cadmium and zinc, 12.5 SiO2 . ...................................

5.13 nickel, no calcium, 12.9 Si02.. . ........................................

10.7 magnesium, 2.9 calcium, 27.6 Si02 . ....................................

Si02 not 

precipitated 

(millimoles).

4.0

2.4

3.0

3.2

3.3

4.7

2.3

2.7

2.3

2.2

2.3

2.8

a Double quantities taken.

Quantities taken in each experiment:

9.10 millimoles MiS04 per liter of the mixture. 

9.10 millimoles M2S04 per liter of the mixture. 

4.95 millimoles Na2O,15.2 millimoles Si02 per liter of the mixture.

I SUMMARY OF RESULTS OF EXPERIMENTS WITH SILICATES.

From the above results it appears that:

1. The order of precipitability of the metals by soluble silicate 

solutions is roughly as follows, beginning with the most precipitable: 

Copper, zinc, manganese, cadmium, lead, nickel, silver, magnesium, 

calcium.

2. The metals are precipitated almost as they would be by solutions of caustic soda, so that no marked specific effect appears to bo 

due to the presence of the silicate.

3. Most of the silica is precipitated.

4. Some silica remains in solution, even with an excess of the 

metallic salt, and the quantity that remains appears to be independent of the nature of the metallic salt.
FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS. 45

GENERAL VIEW OF THE RESULTS OF THE INVESTIGATION.

For comparison the several series obtained in the study of fractional precipitation in aqueous solution have been brought together 

in the following table:

Order of solubility for each class of compounds deduced from these experiments on fractional precipitation.

Sulphides.

Mercury .........

Copper..........

Antimony.......

Tin .............

Zinc ............

Nickel..........

Cobalt ..........

Ferrous . ........

Manganese ......

Carbonates.

MinVol

Silver

Calcium ......

Bicarbonates.

Mercury ......

Lead .........

Copper .......

Cadmium. ....

Silver .......

Ferrous. ......

Zinc ..........

Hydroxides.

Aluminum .....

Zinc ...........

Nickel.........

Silver..........

Ferrous ........

Manganese .

Silicates.

Copper.

Zinc.

Cadmium.

Lead.

Nickel.

Silver.

Magnesium.

<£»

It would be possible to work out still other series by fractional 

precipitation. For example, the series of chlorides would be relatively short, whereas that of the phosphates would be longer, for 

most phosphates are insoluble. The compounds studied, however, 

cover the ordinary geologic transformations and the studies show 

the general relations. Furthermore, in the tabulation of the results 

no attempt has been made to do more than arrange the several 

series in parallel columns. By a scheme that would exhibit cross 

relations all possible compounds might be arranged in a single 

series that would show the order of precipitation under like conditions. Such a series, however, would, have little practical application, 

for several reasons. A comparison of two compounds with no 

common radicle would really imply the possible existence of at least 

four compounds, a degree of complexity greater than that considered 

in the present study of fractional precipitation. Furthermore, the 

effects of mass action are always conjoined with the solubility 

requirements, a fact illustrated by the relations of certain silver 

haloids at Tonopah, Nev.* In this locality silver chloride, cerargyrite, was deposited before the chlorobromide, embolite, and the

1 Burgess, J. A., The halogen salts of silver and associated minerals at Tonopah, Nev.: Econ. Geology, 

vol. 6, p. 13,1911.
46 FRACTIONAL PRECIPITATION OF ORE-FORMING COMPOUNDS.

iodide, iodyrite, which is more insoluble than the cerargyrite, probably because of the great excess of chlorides in the original solution. 

This series well illustrates the application of the mass law. Silver 

iodide is more insoluble in water than silver chloride, but if a solution 

of a silver salt meets a solution containing a great deal of chloride 

and very little iodide, after mixing (it being assumed for argument 

that precipitation does not occur instantaneously) the silver ion 

concentration will have a fixed value, the chloride ion concentration 

will be very large, and the iodide ion concentration very small. 

Hence the solubility product of silver chloride will be exceeded 

and that of silver iodide not be exceeded, although the solubility 

product of silver chloride is larger than that of silver iodide. Silver 

chloride may therefore be deposited before silver iodide, in apparent 

contradiction to the stated solubilities of the two compounds in 

water. On the other hand, the solubility of a compound in water 

is one of its most important properties, and it was to determine 

the order of these solubilities that these experiments on fractional 

precipitation were made.

o