DEPARTMENT OF THE INTERIOR

Hubert Work, Secretary

U. S. GEOLOGICAL SURVEY 

George Otis Smith, Director

Bulletin 778

CHEMISTKY OF DEPOSITION OF NATIVE COPPEE 

J FEOM ASCENDING SOLUTIONS

BY

ROGER C. WELLS

WASHINGTON

GOVERNMENT PRINTING OFFICE) 

1925
CONTENTS

Page

Introduction 1

Native copper deposits _ _ _ 1

Source of the copper__ _ _ 2

Acknowledgments_ __ _ 3 ,'

Chemical possibilities relating to the existence of copper at depth 3 ^

Copper in the absence of water_______________________ 3

Copper in the presence of water __ 5 |

Oxidation-reduction reactions ____ _ _____ _ 12 . Solubility of copper and some copper compounds _ ______ 12 j

Transportation of copper 19 (

Precipitation of copper ____ _ 21 j v

Formation of native copper through oxidation of sulphur by ferric com- V

pounds.___ - . 26

Hypothetical reactions___ _ __ 26

Steinmann's experiment__ _ 28

Action of sulphide waters on ferric oxide ____ 28

Ferric oxide, copper sulphide, and acid 29 

Equilibrium between ferric oxide and acidified solutions of ferric

sulphate _ _ _ 29 V

Bleaching accompanying copper _ __ _ _______ 30 <""

Results with an excess of acid ____ _ _______ 32

Cupric sulphate and covellite_________________________ 33 ^

Theory of the reaction___ _ ____________ ________ 34 ,

Application to chalcocite________________________ 35

Cupric sulphate and chalcocite_____________________ 35

Results of experiments _ _ ___ __________ 35

Effect of sulphuric acid________________________ 38 , Cuprous chloride and chalcocite___________.______'.__________ 39

Theories proposed to account for the deposition of native copper from

ascending solutions __________________'. _______________-___ 39

Deposition by cooling____________________________ 39

Deposition by dilution_______________________________ 39

Deposition of fractional oxidation of copper sulphide_________ 40 ,

Deposition by neutralization and reduction_______________ 41 \

Comparison of the theories__________________________ 42

Experiments ____________________________________ 42

Index _____ ______________________ 71 >: 

ILLUSTEATIONS

FIGURE 1. Apparatus used in experiment 52a________________ 50 

2. Apparatus used in experiments 149-153______________ 63 

ii
CHEMISTKY OF DEPOSITION OF NATIVE COPPER FEOM 

ASCENDING SOLUTIONS

INTRODUCTION

NATIVE COPPER DEPOSITS

Many theories have been proposed at one time or another t© 

account for the deposition of native copper in ore deposits, but nearly 

every one of them is unsatisfactory for the reason that some factor 

has been neglected which would, if considered, greatly modify the 

theory if not entirely invalidate it.

As is well known, the Lake Superior district affords a conspicuous example of the occurrence of native copper. In studying this 

field a number of geologists of the Calumet & Hecla Consolidated 

Mining Co., including Mr. B. S. Butler, Dr. T. M. Broderick, and 

Prof. L. C. Graton, arrived at certain views as to the origin of copper.1 These views were based more or less on geologic observations, 

but as they involved chemical changes about which little or no information was available, the problem was brought to the attention of 

the United States Geological Survey, with the hope that chemical 

experiments could be made which would decide some of the questions 

involved. The Survey undertook the investigation, which is really 

a continuation of the fundamental work of Stokes, and this paper 

presents some of the results so far obtained.

The writer visited the mines for a few days in June, 1922, and 

observed typical ores of the conglomerate, the amygdaloid rock, and 

the fissures, but he has necessarily relied chiefly on the work of 

others for the geologic features that bear on the origin of deposits 

of native copper. His endeavor has been principally to broaden thechemical knowledge relating to the problem. The subject is so large,, 

however, that obviously, in spite of the large number of experiments 

made, much still remains to be learned concerning the chemistry of 

the ores and associated minerals.

It would be inappropriate to give an extensive description of thes 

geology of Keweenaw Point, Mich., in this paper, as the literature

1 Graton, L. C., explanatory statement read at the Ann Arbor meeting of the Society o£ 

Economic Geologists, Dec. 29, 1922.

1
2 DEPOSITION OF NATIVE COPPER

on the subject is already very large.2 Suffice it to say that the principal copper deposits are found in extensive lodes that parallel certain of the Keweenaw lava flows, now much tilted. The copper also 

occurs to some extent in faults and fissures.

SOURCE OF THE COPPER

From an extensive study of the minerals associated with the Lake 

Superior copper Pumpelly 3 long ago concluded that the copper was 

reduced from oxidized compounds by ferrous compounds and that 

the copper compounds were already probably in part concentrated 

as sulphide, silicate, carbonate, and sulphate before their burial by 

later lava flows.

On the other hand Smyth 4 advanced the view that the copper 

was brought to its present position from below by solutions that 

entered the formations after they were tilted.

These two ideas have both had their defenders up to the present 

time.

The work described in this paper was done primarily to throw 

light on chemical possibilities in ascending solutions, but much of 

the information obtained will doubtless be suggestive for other conditions as well. Little attention has been given to the well-known 

reactions that pertain to enrichment. The need suggested by 

the geologic evidence was rather for more knowledge of the 

chemical possibilities in deep or ascending solutions, especially 

those at high temperature and pressure or such as might have constituted magmatic emanations. Aqueous solutions have been studied 

up to 350° C. and a few melts up to about 450°. The results show 

that copper might be deposited by ascending solutions under certain conditions, and the work has gone far enough to refer the 

copper-bearing solutions to a source that was of a distinctly reducing character. Certain more or less oxidized products, however, 

such as sulphur, sulphuric acid, or cupric sulphate, have also been 

assumed in the discussion or used in many of the experiments.

The first part of this paper presents a discussion of the experimental evidence, old and new, that bears on the problem, and the 

second part gives a record of the experiments performed by the 

writer.

2 Some of the more instructive treatises are: Irving, R. D., The copper-bearing rocks of 

Lake Superior: U. S. Geol. Survey Mon. 5, 1883; Van Hise, C. R., and Leith, C. K., The 

geology of the Lake Superior region: U. S. Geol. Survey Mon. 52, 1911; Lane, A. C., 

The Keweenaw series of Michigan: Michigan Geol. and Blol. Survey Pub. 6, 1911.

8 Pumpelly, Raphael, The paragenesls and derivation of copper and Its associates on 

Lake Superior: Am. Jour. Scl., 3d ser., vol. 2, p. 352, 1871.

4 Smyth, H. L., On the origin of the copper deposits of Keweenaw Point: Science, new 

ser., vol. 3, p. 251, 1896.
EXISTENCE OF COPPER AT DEPTH

The writer desires to express his appreciation to the geologists of 

the Calumet & Hecla Consolidated Mining Co., who very generously 

presented the available field and other information bearing on the 

subject and did everything possible to assist and expedite the investigation. He also wishes to thank the members of the United States 

Geological Survey who criticized his manuscript, Mr. A. C. Spencer 

especially having made many helpful suggestions.

CHEMICAL POSSIBILITIES RELATING TO THE 

EXISTENCE OF COPPER AT DEPTH

One of the first questions that arose in the investigation concerned 

the most probable form in which copper may exist at depth.

It has long been held that copper may exist in magmas in the form 

of a sulphide, as chalcopyrite and bornite are so generally found as 

primary minerals, but few investigators have considered the changes 

that might intervene between the magmas and the final ore deposits.

Perhaps the most significant point to decide first is whether water 

was present, or, if it is impossible to be sure on this point, it may be 

well to set forth the results to be expected in the absence and in the 

presence of water to see if there are significant differences in the two 

classes of reactions.

COPPER IN THE ABSENCE OF WATER

Some of the reactions that take place in the smelting of copper ores 

show what may be expected under dry conditions. For example,

Cu2Si03+FeS=:Cu2S+FeSi03 (1)

Cu,S+2Cu2 0=6Cu+SO2 (2)

Cu2S+02=2Cu+S02 (3)

The first reaction may explain the fact that copper silicate has not 

been found as a product of igneous origin. The next two reactions 

might explain the development of metallic copper in a dry magma 

if by chance sulphur was deficient, hydrogen was absent, and there 

was not enough oxygen present to oxidize both the copper and the 

sulphur.

The fact that cuprous sulphide and oxide melt at rather high temperatures, together with the oft-repeated statement that the mineralizers in magmas probably include chlorides and fluorides, turned the 

writer's attention to the relatively low melting point of cuprous 

chloride, 434° C., especially in view of a statement by Schnabel 5 that

6 Schnabel, Carl, Handbook of metallurgy, translated by H. Louis, vol. 1, p. 11, Macmillan, 1898.
cupric chloride attacks copper sulphide, producing cuprous chloride 

with the separation of sulphur.

The reaction last mentioned has been verified (experiment 73r 

p. 53). It was also found, as expected, that the melting point of 

cuprous chloride is much lowered by the presence of other chlorides.

A mixture of anhydrous cuprous chloride with 23 per cent of calcium chloride appeared to melt at 384°. Menge,6 who studied this 

system much more carefully, found the eutectic to contain 12 per cent 

CaCl2 and its melting point to be 393°.

Mixtures of cuprous chloride with 10, 20, 30, 40, and 60 per cent 

of sodium chloride all began to melt around 315°, indicating a 

eutectic at this temperature, but the exact composition of the eutectic 

mixture could not be determined by this method. It seems probable 

that other chlorides might lower this point still further.

Truthe 7 has shown that Cu2 S and CuCl form a eutectic with 

16.3 per cent Cu2S, which melts at 392°. The writer hopes to study 

melts of this character further, as soon as opportunity can be found 

to do so, to see if such a study will lead a step nearer to the experimental realization of " magmatic solutions."

Small proportions of native copper in nickel-iron segregations 

in basalts have been described.8 It seems almost necessary to conclude that a very strong reducing agent, such as carbon or hydrogen, 

has acted in these places, providing the segregations are not ascribed 

to such a remote source as the interior metallic core of the earth or 

to meteorites.

The statement of certain investigators 9 that dry hydrogen does not 

reduce cuprous sulphide to copper rapidly, if at all, below 600° 

seems to be supported by work of the writer. A preparation made 

by heating CuS in hydrogen at 600° gave 79.88 per cent copper 

and 19.86 per cent sulphur; the theoretical percentages for Cu2S 

are 79.86 and 20.14, respectively. Microscopic specks of copper 

were visible in some preparations, however, the cause of which is 

not easily explained. The hydrogen sulphide produced when CuS 

is taken might prevent the formation of copper for a time, but the 

quantity of free copper did not seem to increase on continuing 

the heating for six hours.

8 Menge, 0., Die binSren Systeme von MgCl2 und CaCl2 mit den Chloriden der Metalle 

K, Na, Ag, Pb, Cu', Zn, Sn" und Cd: Zeitschr. anorg. Chemie, Band 72, p. 162, 1911.

7 Truthe, W., Ueber das Verhalten der Sulflde von Pb, Cu', Ag und des Cu20 in der 

Schmelzen der zuhorigen Chloride: Zeitschr. anorg. Chemie, Band 76, p. 167, 1912.

8 Beyschlag, F., Krusch, P., and Vogt, J. H. L., The deposits of useful minerals and 

rocks, translated by S. J. Truscott, p. 341, Macmillan, 1916.

9 Wegscheider, R., Bemerkungen zur quantitativen Bestimmung des Kupfers als Sulfur: 

Monatsh. Chemie, Band 14, p. 315, 1893.
Posnjak, Alien, and Merwin find that chalcocite is stable in a 

vacuum even at its melting point 10 that is, 1,130° C.

According to Uhl,11 dry sulphur dioxide reacts with copper at a 

red heat to form cupric sulphate and cuprous sulphide, thus:

3Cu+2S02=CuS04+Cu2S (4)

This action does not seem to go on very fast at 500°, however, 

with dry sulphur dioxide (experiment 90), but the moist gas yields 

copper sulphide readily.

Prescott and Johnson 12 state that cupric sulphate is decomposed 

at a red heat, thus:

2CuS04 =2CuO+2S02+02 (5)

The point of decomposition is probably nearer a white heat, although there is no experimental evidence on this point.

The scanty evidence available therefore points to a few Avays in 

which copper might be formed in a dry environment, mostly at high 

temperatures. They simplify to the oxidation of the sulphide or 

the reduction of oxidized compounds by hydrogen. But reduction 

by hydrogen yields water and illustrates how difficult it is to proceed without supposing water to be a possible component. In fact, 

the improbability that water is ever Avholly absent justifies us in 

passing to other possibilities as more likely and common.

COPPER IN THE PRESENCE OF WATER

Morey 13 has advanced the view that crystallization of a magma 

may cause increasing pressure to develop in the remaining fluids, 

provided these contain dissolved gases, especially such as water. 

Wherever vents exist the fluids will tend to escape. The fluids 

might be " melts," but as water and other mineralizers increase, the 

"melts" would have more and more the character of solutions. In 

what follows, water will be assumed to be one of the constituents.

In recent analyses of the Deccan and other traps by Washington 14 

the water content is uniformly high, and he considers this to represent water originally present in the basaltic lavas.

10 Posnjak, E., Alien, E. T., and Merwin, H. S., The sulphides of copper: Econ. Geology, 

vol. 10, p. 494, 1915.

11 Uhl, J., Ueber Elnwlrkung von Schwefeldioxyd auf Metalle: Deutsche chem. Gesell. 

Ber., Band 23, p. 2151, 1890.

u Prescott, A. B., and Johnson 0. C., Qualitative chemical analysis, 5th ed., p. 325, 

1901.

18 Morey, G. W., The development of pressure in magmas as a result of crystallization: 

Washington Acad. Sci. Jour., vol. 12, p. 219, 1922.

14 Washington, H. S. ( Deccan traps and other plateau basalts: Geol. Soc. America Bull., 

vol. 33, p. 765, 1922.
0 DEPOSITION OF NATIVE COPPER

A few years ago the French chemist Gautier 1B showed that steam 

acting on chalcocite at a bright-red heat produces metallic copper, 

thus

Cu2S+2H2 0 *± 2Cu+2H2+SO2 (6)

This reaction suggests that in hot magmas in which water is a 

prominent constituent the sulphide of copper might be changed to 

the free metal. But it must be noted that the reaction was obtained 

only when the resulting gases, sulphur dioxide and hydrogen, were 

swept away by fresh steam. With little opportunity for the escape 

of gas there would probably be an equilibrium instead of a one-way 

reaction. In fact, we must imagine that at depth there are many 

reversible reactions involving water, sulphur, hydrogen, hydrogen 

sulphide, the two oxides of sulphur, the two oxides of iron, various 

compounds of copper as well as the free metal and probably chlorides also, not to mention silica, carbon dioxide, and the different 

bases. The writer knows of few experiments at high temperatures 

with such systems as these, but certain reversible reactions have been 

studied which throw light on the relations that are to be expected.

Randall and Bichowsky 10 have investigated the following reaction 

from 887° to 1,372° C.:

S02 + 3H2 ^2H20 + H2S (7)

They find an equilibrium, in which water is the predominant constituent; all the other gases are. present, however. The concentration 

of any one constituent depends on that' of all the others. As the 

temperature falls the concentration of H2S increases, whereas that 

of the other gases decreases. Increasing pressure favors the reaction 

toward the right.

In view of this equilibrium it ought to be possible to reverse 

Gautier's reaction (6), starting with copper, hydrogen, and sulphur 

dioxide. The writer easily brought about this reversal, as shown in 

experiment 74, where copper sulphide and water are formed from 

copper, sulphur dioxide, and hydrogen. The results at 900°, then, 

indicate that the formation of either copper or chalcocite is possible 

according to the relative proportions of the gases present. The 

exact proportion of steam in the mixture that would be necessary 

to convert chalcocite into copper at any temperature remains to be 

determined, but it would probably be high. The equilibrium between ferrous oxide and water discussed later, equation (14&), 

would tend to increase the proportion of water to hydrogen as the

15 Gautier, A., Action de la vapeur d'eau sur les sulfures an rouge: Compt. Rend., vol. 

142, p. 1465, 1906.

16 Randall, Merle, and Bichowsky, P. R., Equilibrium in the reaction between water and 

sulphur at high temperatures: Am. Chem. Soc. Jour., vol. 40, p. 368, 1918.
EXISTENCE OP COPPER AT DEPTH 7

temperature rises. On the other hand, the relations demanded by 

the equilibrium corresponding to equation (7) are such that the 

concentrations of S02 and H2 increase along with that of H2 0, 

which would be adverse to reaction (6). On this account as well as 

because of the probable significant proportions of sulphur and its 

compounds in magmatic gases it seems logical to conclude that 

copper would tend to crystallize from magmas as cuprous sulphide. 

At lower temperatures, however, by virtue of other equilibria that 

will be considered presently, equations (9), (10), and (11) and 

experimental work by the writer, the possibility is presented that 

copper might exist in solution as cuprous sulphate.

Hydrogen sulphide exists partly in a dissociated condition at elevated temperatures,17 but this is without effect on the stability of 

cuprous sulphide.

Hydrogen is well known to be an effective reducing agent at elevated temperatures. It will even reduce aqueous solutions of several 

oxidized salts of copper to metal at moderate temperatures under 

pressure, as shown by Ipatief and Verkhovsky.18 The acetate 

yielded the free metal after heating at 92° for four days with the 

hydrogen at a pressure of about 350 atmospheres. Copper was precipitated from the sulphate at a temperature above 90° and the 

hydrogen at a pressure above 120 atmospheres, but none was produced at 20° with the hydrogen under a pressure of 500 atmospheres.

Cupric chloride, 2 normal, was reduced to cuprous chloride at 155° 

by hydrogen at a pressure of 100 atmospheres; whereas at 180° a 

0.2 normal solution was partly reduced to metallic copper, apparently yielding the following reversible equilibrium:

2CuCl + H2 ^2Cu + 2HCl (8)

Thermodynamic calculations show that hydrogen, even at a pressure of 1 atmosphere, should precipitate copper from neutral cupric 

sulphate solutions of ordinary concentration, and higher pressures 

would be still more effective, but according to Ipatief the reaction is 

too sluggish to be within the range of experimental observation 

much below 90° C.

A current of pure hydrogen may be made to reduce sulphuric acid 

completely under certain conditions at 700°-900°, yielding water 

and hydrogen sulphide,18a but the reversibility of reaction (7) would

17 Preuner, G., and Shupp, W., Dissociation des Schwefelwasserstoffes: Zeitschr. physlkal. Chemie, vol. 68, p. 157, 1909.

18 Ipatief, V. N., and Verkhovsky, V., Ueber das Verdrangen der Metalle aus wassrlgen 

LSsungen ihrer Salze durch Wasserstoff bei hohen Temperaturen und Drucken: Deutsche 

chem. Gesell. Ber., Band 42, p. 2078, 1909; Band 44, p. 1755, 1911. Also published in 

Russian in Russ. phys.-chem. Soc. Jour., vol. 41, p. 779, 1909; vol. 43, p. 1746, 1911.

isa Vila. A., Sur la reduction de 1'aclde sulfurique en hydrogene sulfure": Compt. Rend., 

vol. 179, pp. 1163-1165. 1924.

42980 25 2
8 DEPOSITION OE NATIVE COPPEE

doubtless limit this action in nature, at least if the gases were 

confined.

The following facts are noted with special reference to the presence of sulphur dioxide. It follows from reaction (7) that hydrogen 

converts sulphur dioxide partly into hydrogen sulphide. In the 

absence of hydrogen there is another equilibrium which was studied 

at 300° by Bichowsky.19

3S02 + 2H2 0<^S + 2H2S04 (aq) (9)

Bichowsky has found that an excess of water will displace this 

reaction almost wholly to the right, whereas an excess of sulphuric 

acid reverses it. Another equilibrium, which is closely associated 

with this one, is stated thus :

^3H2S + H2SO< (10) 

Lewis and Eandall 20 also give the following equilibrium at 445° :

3S+2H2 0^2H2S+S02 (11)

A consideration of these equilibria shows that the presence of sulphur or sulphur dioxide with water must necessarily yield some 

hydrogen sulphide over a considerable range of temperature. In 

fact, the writer has fo'und that sulphur dioxide in the presence of 

water between 150° and 200° converts copper into cuprous sulphide 

without the necessity of having hydrogen present (experiment 48).

More recently it has been shown that the simultaneous presence of 

acid sulphides and sulphites that is, the hydrosulphide radical,

HS', and the bisulphite radical, HSO3 ' in solution at ordinary 

temperature leads to the formation of thiosulphates, especially when 

the bisulphite is in excess and the acidity is no greater than 

that of the bisulphite itself.21 This observation suggests that thios'ulphates should commonly occur in ore solutions. Of course, such 

solutions are not generally available for laboratory study, but thiosulphates have been reported in one or two so-called " mine waters." 

Tetrathionates might also be expected to occur in such waters.

A consideration of the equilibria referred to above shows that it 

makes little difference whether the substances used are hydrogen 

sulphide and hydrogen, s'ulphur dioxide and hydrogen at a high

19 Bichowsky, F. R.,, Equilibrium in the reaction between sulphur dioxide and water : 

Am. Chem. Soc. Jour., vol. 44, p. 116, 1922.

20 Lewis, G. N., and Randall, Merle, Equilibrium in the reaction between water and 

sulphur at the boiling point of sulphur : Am. Chem. Soc. Jour., vol. 40, p. 362, 1918.

21 Foerster, F., and Mommsen, E. T., Beltrag zur Kenntniss der Thiosulfate : Deutsche 

chem. Gesell. Ber., Band 57, p. 258, 1924.
EXISTENCE OF COPPER AT DEPTH 9

temperature, sulphur dioxide and hydrogen sulphide, sulphur dioxide and. water, or .sulphur and water at 150° to 200°; the tendency 

over a wide range of temperature is to convert copper into chalcocite.

The results when iron compounds are also constituents of these 

systems would be of considerable interest from the viewpoint of a 

geologist. Lack of time has prevented a full investigation of these 

possibilities, but a number of the experiments described in this paper throw light on the problem under certain conditions.

That sulphur and sulphur dioxide should be considered possible 

components of magmatic gases perhaps needs explanation. They are 

well known as volcanic products, but their origin has generally been 

ascribed to reactions at or near the surface of the earth. Their existence in magmatic gases must obviously be governed to a considerable extent by the equilibria represented by such equations as (7) 

and (11), and accordingly it must largely depend on the proportion 

of hydrogen sulphide in such gases. If the partial pressure of hydrogen sulphide was reduced by any reaction to a small value that 

reduction would cause a shift in the equilibrium represented by 

equation (7), for example, and would practically remove S02 from 

consideration.

The experimental evidence on the possible concentration of hydrogen sulphide is very scanty, and much more work will be required to clarify our knowledge of the subject. We have, however, 

some data on the dissociation pressures of sulphides 22 and may also 

feel certain that many sulphides are hydrolyzed by water at high 

temperatures. It is consequently logical to assume that appreciable 

concentrations of hydrogen sulphide may exist in magmatic gases, 

and if so the other components required by the equilibria that 

have been discussed would inevitably be present.

It has not seemed wise to limit the discussion in this paper to compounds which must necessarily have come directly from magmatic 

sources, for products may conceivably have been formed by partial 

oxidation at one period, buried beneath later formations, and at a 

subsequent period subjected to the influence of new accessions of 

steam. Likewise brines that may have been buried at one period 

could act as potent chemical agents under the influence of heat at a 

subsequent period. There is such a long range from oxidized compounds at the surface to deep magmas that the study of the possible 

intermediate conditions seems not only permissible but necessary.

28 Alien, E. T., and Lombard, R. H., A method for the determination of dissociation 

pressures of sulphides and its application to covellite and pyrite: Am. Jour. Sci., 4th ser., 

vol. 43, p. 175, 1917.
10 DEPOSITION OF NATIVE COPPER

The relations involved in mixtures containing sulphur dioxide 

and oxygen will next be considered. Bodenstein and Pohl 23 have 

shown that the equilibrium

2S08^2SO2+02 (12)

is displaced very largely toward the right at 900° but almost wholly 

to the left at 400° and lower. The dissociation of SO2 need not be 

considered here,24 as it barely begins at 1,000° C. We may therefore 

consider the action of sulphur trioxide and sulphuric acid on the 

substances under scrutiny.

It seems possible that concentrated sulphuric acid might be formed 

in nature under certain conditions for example, by the escape of 

the more volatile gases or by the more rapid diffusion of lighter 

gases from mixtures like those considered above, just as hydrocarbons are believed to separate into various fractions in the earth 

through diffusion. The main question is whether the free acid, if 

formed, could long survive uncombined.

It is well known that concentrated sulphuric acid can be heated in 

.glass tubes without attacking them very rapidly. Foerster 25 has 

shown that the attack by concentrated acids is less rapid than that 

by water, and that the attack increases with the dilution of the acid. 

The results found in experiment 115&, in which an acid solution was 

heated 4 days at 300° in a Jena glass tube, indicated that very little 

acid was neutralized by the glass. In experiment 1200, which lasted 

20 days with half-normal acid at 200°, only 0.0293 gram of acid was 

neutralized, and in experiment 120&, which lasted 11 days at 300°, 

only 0.0100 gram was neutralized. 26

On the other hand, the attack of finely powdered minerals by 

acids has even been used in mineral analysis. Mitscherlich 27 found 

that the following minerals, when finely powdered, are decomposed 

at 200° by a mixture of 3 parts sulphuric acid to 1 part water: Some 

of the spinels, tantalite, ilmenite, pyroxenes, talc, asbestos, chlorite, 

beryl, garnet, feldspars, petalite, and cordierite. Even glass, in the 

form of a powder, seems to be decomposed and dissolved by water.28 

"No investigators seem to have reached conditions of equilibrium in

23 Bodenstein, M., and Pohl, W., Gleichgewichtmessungen an. der Kontaktschwefelsaure: 

Zeitschr. Elektrochemie, Band 31, p. 373, 1905.

24 Ferguson, J. B., The thermal dissociation of sulphur dioxide: Am. Chem. Soc. Jour., 

vol. 41, p. 69, 1919.

25 Foerster, F., Ueber die Einwirkung von Siiuren auf Glas: Zeitschr. anal. Chemle, 

Band 33, p. 305, 1894.

26 Compare Alien, E. T., Crenshaw, J. L., and Merwin, H. H.f Effect of temperature and 

acidity in the formation of marcasite and wurzite: Am. Jour. Sci., 4th ser., vol. 38, p. 

397, 1914.

27 Mitscherlich, Alexander, Jour, prakt. Chemie, Band 81, p. 116, 1860.

28 Mylius, F., and Foerster, F., Ueber die Beurtheilung von GlasgefBssen zu chemischen 

 Gebrauche: Zeitschr. anal. Chemie, vol. 31. p. 264, 1892.
EXISTENCE OF COPPER AT DEPTH 11

the attack of glass by any solution. The suggestion that possibly/ 

alunite was formed in experiment 120a seems to be a step in that 

direction.

In experiments 129 and 130 the quantity of sulphuric acid neutralized by finely powdered leucite rock and Jena glass tubing during 2 days heating at 250° was quantitatively determined. The 

leucite rock was found to be attacked over 10 times as fast as the 

glass tubing under similar conditions. The detailed figures are 

given with the description of the experiments elsewhere in this 

paper (p. 59).

That sulphuric acid, if formed in nature, is probably as a rule soon 

converted into sulphates is undeniable; on the other hand, acid 

waters from volcanic sources have been reported from certain localities, so that it is no great stretch of imagination to assume that they 

may have assisted in chemical reactions, even granting that the 

acidity in some of them is due to oxidation of sulphides.29

In this connection it should be noted that cuprous sulphide and 

silver sulphide may be dissolved by concentrated sulphuric acid 

(experiments 35, 58, and 70). The writer believes that such solutions would under certain conditions deposit the free metals, as well 

as the sulphides, on dilution and cooling. As few solvents for copper sulphide are known this one has chemical if not geologic interest.

The early experiments of Pickering 30 on the action of concentrated sulphuric acid on copper sulphide are of historic interest. Iin 

his experiments any gas that formed was allowed to escape at atmospheric pressure. He gives the following equations for the reactions:

Cu2S + 2H2S04 = CuS + CuS04 + S02 + 2H20 (13) 

CuS + 2H2S04 = S + CuS04 + S02 + 2H20 (14)

The time allowed for these experiments was only 30 minutes or 

less, and the temperature ranged from 19° to 220°.

One of the most noteworthy equilibria that may relate to the> 

alteration of rocks and the genesis of ores is that between water and 

the two oxides of iron :

3FeO + H20<^H2 + Fe304 (14o>

This equilibrium has recently been investigated very carefully 

from 675° to 1,000° C. by Eastman and Evans,31 who find that thehydrogen in the gaseous mixture ranges from 46 to 9.9 per cent by

29 .For acid waters see Clarke, F. W., The data of geochemistry, 5th ed.: TJ. S. Geol_ 

Survey Bull. 770, p. 199, 1925.

80 Pickering, S., The action of sulphuric acid on copper: Chem. Soc. Jour., vol. 33, p_ 

118, 1878.

31 Eastman, E. D., and Evans, R. M. ( Equilibria involving the oxides of iron : Am.. 

Chem. Soc. Jour., vol. 46, p. 888, 1924. References to the earlier literature are given.
12 DEPOSITION OF NATIVE COPPER

volume between these two temperatures respectively. The presence 

of such percentages of hydrogen would in consequence of the equilibria that have already been discussed tend strongly to keep copper 

reduced to sulphide, possibly even to metal at the highest temperatures.

In the paper last mentioned Eastman and Evans also refer to the 

equilibrium

\ (15)

At temperatures between 500° and 900°. both carbon dioxide and 

carbon monoxide are present in considerable proportions at equilibrium. It is believed that carbon monoxide, like hydrogen, would 

reduce most of the oxidized compounds of copper to free metal, but 

according to Schnabel it is without action on cuprous sulphide below 

a red heat.

OXIDATION-KEBUCTION REACTIONS

In what has been said and in what follows the frequent occurrence 

of the terms oxidation and reduction will be noted. These terms 

mean more than simply the addition or removal of oxygen, for, in 

the most general sense, any reaction that increases the positive 

charges on an atom, or its positive valence, may be considered an 

oxidation. For example, the changes of Cu+ to Cu++ or of S" to S 

are oxidations, and the reverse changes are reductions. To refer to a 

reaction as an oxidation is generally ambiguous, as it is clear that 

the oxidation of one substance is invariably linked with the reduction of another substance ; this includes the reaction in which oxygen 

passes from free oxygen to oxide. Copper may exist in at least three 

valences zero valence, or the free metal, and cuprous and cupric. 

Sulphur is even richer in forms sulphide, S~~ ; the free element, S ; 

the dioxide, S02 , corresponding to S++++ ; and the trioxide, S03 , corresponding to S++++++. In many oxidation-reduction reactions sulphur passes only from S~ to S ; in others, as will be shown, it passes 

from S~~ to S++++++, thus exhibiting its maximum shift of eight positive charges.

Let us next consider briefly the solubility of copper and certain of 

its compounds in different reagents, returning later to reactions involving the precipitation of copper.

SOLUBILITY OF COPPER AND SOME COPPER

COMPOUNDS

Dilute sulphuric acid, in the absence of air or other oxidizing 

agents, has very little action on copper at ordinary temperature.
SOLUBILITY OF COPPER 13

Even at 200° C. experiment 49 showed solution to the extent of only 

8 milligrams per liter. Above this temperature the attack goes 

somewhat further and yields a number of products, including 

cuprous sulphide. At least three reactions take place simultaneously, in order of prominence, as follows:

8Cu + 4H2S04 = Cu2S 4- 3Cu2S04 + 4H20 (16)

5Cu+4H2S04 = Cu2S + 3CuS04 + 4H20 (17)

Cu+2H2S04 = S02 +CuS04 + 2H20 (18)

In experiment 137 a tube containing copper and sulphuric acid 

was heated and observed while hot. Some Cu2 S was seen after 

heating two days at 200°. The solution was not noticeably blue, 

which may have been because the copper was largely in the cuprous 

state. At 250° the solution first turned a very pale yellow and later 

a pale blue.

That both cuprous and cupric sulphates are formed when copper 

dissolves in sulphuric acid is probably connected with the fact that 

an equilibrium tends to be reached between these two salts in the 

presence of metallic copper, thus 9

Cu2S04<=±Cu + CuS04 (19)

It seems probable that a rapid shift of this equilibrium with 

changes of temperature explains a feature noted in many experiments the separation of copper on cooling, as observed in experiments 104, 128, 132a, and 1326. In experiment 133 all the copper 

was reerystallized. Although a shift of equations (17) and (18) 

with changes of temperature should eventually be considered, the 

shift of reaction (19) appears to be more rapid. The copper appears in the form of beautiful crystals, generally separate or easily 

separable from the copper originally taken. In several experiments 

it was determined separately.

If the reerystallized copper was formed by the shift of an equilibrium according to equation (19) on cooling, its amount should 

measure the copper that existed as cuprous sulphate when hot, except for the small amount which does not precipitate on cooling. 

As the. cuprous salt and total dissolved copper left "after cooling 

are known, the copper in each state when hot was calculated by the 

following equations:

Cu" (hot)=total dissolved Cu(cold)-

(Cu(cold)+Cu reerystallized) (20) 

Cu' (hot) = (reerystallized CuX2)+Cu' (cold) (21)
14 DEPOSITION OF NATIVE COPPER

The results obtained in this way are shown in Table 1, and, derived from them, the mass-law " constant " for reaction (19),

[CuSOJ 

*"-[CuaSOJ

as well as the constant used by Abel,32

K -VtCuSOJ 

A [Cu2S04]

(22)

(23)

the theoretical basis for which may be found in his paper. The 

bracketed quantities in these expressions represent concentrations of 

the molecular species written in the bracket, and the "constants" 

are supposed to be constants at any given temperature.

TABLE 1. Mass-law constant for reaction (19)

[Concentrations in gram-molecules per 1,000 grams of water]

Experiment

o 

Abel..... ................ ...

Abel... 

Abel.-.-. .. -

AhAl

1320 ... ..... .

138..... 

139..... - . 

Temperature 

(°C.)

0

25

40

100

250

300

350

[CuSO«]

0.0154

.0151

.0149

.0145

.0041

.0264

.0729

[Cu2S04 ]

0.00045

.00077

.00095

. 00143

.025

.258

.324

Kit

34.2

19.5

15.7

10.1

.16

.10

.22

KA

275

159

128

84

2.5

.63

.84

It appears from Table 1 that both " constants" are in general 

smaller at high temperature than at ordinary temperature, although 

the figures do not fall on perfectly smooth curves. It is evident that

equilibrium (19), at least over a wide range, is displaced strongly 

to the left with rising temperature and in the opposite direction as 

the temperature falls.

It is known from Abel's work that K19 is not strictly constant at 

constant temperature. Doubtless the reason for this failure in constancy is that the measured concentrations do not represent the concentrations of the molecular species written in reaction (19). Part 

of each salt is ionized but in different degrees. The reaction may be 

expressed in ionic form, thus

(24)

(25)

giving the mass-law constant

^Abel, E., Ueber das Gleichgewlcht zwlschen den verschiedenen Oxydationstufen desselben Metalles: Zeitschr. anorg. Chemle, Band 26, p. 396, 1901.
SOLUBILITY OF COPPER 15

but here again the difficulty is immediately encountered of evaluating 

the ionic concentrations. Abel has pointed out the necessity of considering that the cuprous copper enters very largely into a complex 

ion, and more recently it has been found that cupric sulphate is also 

unusually abnormal in its ionization.33 The correct evaluation of 

K24 and its variation with temperature is an interesting problem for 

the physical chemist, but to make any considerable progress on it 

would require more space than should be given to it in this paper.

From Table 1. it is possible to arrive at the total dissolved copper 

calculated to exist at each temperature for the strength of acid then 

present, according to the preceding assumptions. This amount may 

be considered to represent th?, solubility of copper in sulphuric acid 

and is shown in Table 2 in grams per 1,000 grams of water.

TABLE 2. Solubility of copper in sulphuric acid at high temperatures

[All concentrations are in grams per 1,000 grams of water]

Experiment

49

132 

138 

139

Temperature 

(°C.)

200

250 

300 

350

H2S04

25

22 

38 

24

Cuprous 

copper

3.4 

35.4 

44.4

Cupric 

copper

.26 

1.7 

4.6

Total 

copper

0.008

3.66 

37.1 

49.0

The total copper dissolved must also be a function of the concentration of acid at each temperature, and in order to arrive at some 

idea of this relationship the mass-law constants for equations (16) 

and (17) were computed separately. The constants, omitting the 

solids and water, as usual, and neglecting ionization, are as follows:

[Cu2SOJ3

 *MG fTT Qn I* v^ D/

[CuSOJ3

(27)

The results are given in Table 3.

TABLE 3. Mass-law constants for reactions (16) and (17) 

[Concentrations are in gram -molecules per 1,000 grams of water]

Experiment

49

132 

138 

139

Temperature (°C.)

200

250 

300 

350

[H8SO<]

0.25

.22 

.39 

.24

[CwSOd

5X10-»

0.025 

.258 

.324

[CuSOi]

1.2X10-*

0.0041 

.0264 

.0729

KH

6.7X10-' 

0.75 

3.0

K.7

3.0X10-* 

7.8X10-" 

3. 5X10->

33 Lewis, G. N., and Randall, Merle, Thermodynamics, p. 362, McGraw-Hill Book Co., 

1923.
16 DEPOSITION OF NATIVE COPPER

Table 3 clearly shows that both reactions (16) and (17) go 

further as the temperature rises. It was found, however, that the 

sulphuric acid was reduced only to about the same concentration at 

each temperature studied. Both the cuprous and cupric copper increase with rising temperature, but the cuprous is decidedly in 

excess at the high temperatures. In fact, at 350° the concentration 

of the cuprous sulphate exceeds that of the sulphuric acid present.

So far as the writer has been able to determine, reaction (16) can 

not be reversed, although the sulphuric acid is never entirely consumed. Eeaction (17), however, appears to be partly reversible, 

and much study has been devoted to it, the results of which will be 

presented later. This reaction was given by Pickering 34 for certain experiments made by him to show the action of concentrated 

sulphuric acid on copper, but generally his results were represented 

by equation (18). In his experiments sulphur dioxide, if formed, 

was allowed to escape at atmospheric- pressure. In the writer's 

experiments, on ' the other hand, weaker acid was Used, and the 

products were confined in closed tubes. Under these circumstances 

the quantity of sulphur dioxide formed is very small, being generally scarcely enough to determine (experiment 132). In fact, 

it has been shown that sulphur dioxide and cupric sulphate react, 

as described in experiments 61 and 124, or at lower temperatures 

reaction (18) is simply reversed, according to qualitative tests (experiments 46 and 47). Copper will not consume all the available 

sulphuric acid in dilute solutions, which, from the preceding facts, 

appears to be due to the accumulation of cupric salt, sulphur dioxide, 

and perhaps even water itself all substances that tend to reverse 

the action.

Doubtless the reactions might also be explained as caused by certain small concentrations of hydrogen, but the final i-esults would 

not be changed. ' 

Concentrated hydrochloric acid dissolves copper markedly, forming hydrogen and cuprous chloride, which is fairly soluble in the 

acid or other chloride solutions. The more dilute acid solutions dissolve a little copper. Thus, half-normal acid, containing 18.7 grams 

per liter, when heated with copper to 165° for 17 hours contained 

copper equivalent to 0.7 gram per liter (experiment 52). The results of Ipatief have already been mentioned.

Solutions of calcium chloride have been studied because such 

solutions are now found in the deep levels of the copper mines in 

Michigan.

84 Op. cit., p. 11. See also Cundall, J. T., and Fairgrieve, M. McC., The action of 

sulphuric acid on copper: Chem. Soc. Jour., vol. 105, p. 60, 1914.
SOLUBILITY OF COPPER 17

A concentrated solution of calcium chloride, specific gravity 1.5. 

containing approximately equal parts of water and calcium chloride, 

after heating with copper to 175° for 20 hours gave the equivalent 

of 0.07 gram of copper per liter (experiment 55). On the theory 

that this result was caused by the partial hydrolysis of calcium 

chloride into hydrochloric acid and lime, a similar experiment was 

performed with small quantities of silica and ferric oxide present 

also. After 54 hours heating at 190° the copper found in the solution was equivalent to 0.19 gram per liter.

Dilute sulphuric acid which is sufficiently concentrated to dissolve

ferric oxide forms in the presence of ferric oxide acid ferric sulphate,

which dissolves copper rapidly, the iron being reduced to the ferrous

state and the solution becoming blue from the cupric sulphate.

; Hydrochloric acid under the same circumstances also dissolves

i copper. In one experiment at 90° (experiment 57) the equivalent of

1.3 grams of copper per liter was found.

/ Copper is soluble in solutions of ammonium nitrite at ordinary

temperature.35 It is also appreciably soluble in cyanide solutions

^ in the presence of air, but these two solvents need not be considered

1 in geochemical studies. It is more or less attacked and dissolved

by fused caustic soda at 350° C. and above.

/ Cuprous sulphide is somewhat soluble in solutions of alkali cya- 

/ nides. It is also slightly soluble in solutions of ammonium sulphide 

\ or ammonium carbonate. It can also be brought into a colloidal 

condition by water containing hydrogen sulphide. Colloidal solutions have been suggested as possible aids in forming some ore deposits, the escape of gas being suggested as the cause of precipitation.36 The writer has prepared many colloidal suspensions of 

cuprous sulphide but has generally found them too unstable to pass 

through filter papers or layers of powdered minerals (experiment 

86). They are also coagulated by ferric oxide (experiment 89).

It has already been indicated that if sulphur dioxide and oxygen

are components of magmatic emanations, then at certain temperaI tures sulphuric acid and soluble copper compounds may be present.

> Hydrochloric acid is well known as a volcanic emanation. It is

formed when hydrogen acts on the heated chlorides of the copper and

iron groups, or when sulphuric acid acts on chlorides, and also from

the decomposition of calcium chloride and magnesium chloride at

high temperatures. To assume that the last source might account

for the origin of hydrochloric acid in Michigan i? no stretch of the

88 Dhar, N. R., Einwirkung von Metallen, wie Kupfer und Zinc, auf eine wassrige 

Losung von Ammoniumnitrit: Zeitschr. anorg. Chemie, Band 119, p. 174, 1921.

^Tolman, C. P., jr., and Clark, J. D., The oxidation, solution, and precipitation of 

copper in electrolytic solutions and the dispersion and precipitation of copper sulp'hldes 

from colloidal suspensions, with a geological discussion: Econ. Geology, vol. 9, p. 559, 

1914. For the r51e of absorption in the genesis of ores see Kohler, E., Zeitschr. prakt. 

Geologic, vol. 11, p. 49, 1903.
18 DEPOSITION OF NATIVE COPPER

imagination, for concentrated solutions of calcium chloride occur in 

the present deep mine waters, and the writer has actually heated some 

rock from one of the lower levels of the Calumet & Hecla lode and 

obtained water followed by hydrochloric acid.

The temperature at which calcium chloride begins to decompose 

in the presence of moisture is not known exactly. It is probably 

over 300°. Magnesium chloride, on the other hand, shows a tendency 

to decompose around 180°, which may explain the absence of magnesium chloride in the deep mine waters.

Although " alkaline sulphide solutions " have long been considered 

the normal type of deep mineralizing solutions, the waters that have 

been collected and tested probably do not exactly represent magmatic 

emanations but rather mixtures of these emanations with large proportions of circulating waters, generally very dilute and frequently 

not even hot, which in passing through long rocky channels become 

at least neutral if not alkaline.

In a single experiment (79) liquid sulphur dioxide could not be 

seen to dissolve any cuprous sulphide at ordinary temperature or at 

64°, although Prescott and Johnson state that liquid sulphur dioxide 

dissolves sulphur.

Copper sulphate is soluble in cold water to the extent of 210 grams 

to the liter at 25° C., the solid phase being CuSO4.5H2O. This solubility increases regularly up to about 105°, where the solid phase 

becomes CuS04 .3H2 0, and on further increase of temperature the 

solubility changes very slightly. The addition of sulphuric acid 

decreases the solubility of cupric sulphate markedly at all temperatures. On account of its significance for the problem in hand 

the writer made some determinations of the solubility in rather 

concentrated solutions of sulphuric acid at 200° C., with the results 

shown in the following table:

TABLE 4. Solubility of cupric sulphate in sulphuric acid at 200° C.

Volume of

solutio n 

(cubic 

centimeters)

1,000

1,000

1,000

Weight of 

solution 

(grams)

1,727

1,690

1,649

H2SO« 

(grams)

1,615

1,344

1,205

HsO

(grams)

104

294

365

CuSOi 

(grams)

8

52

81

Metallic copper reacts with all cupric solutions to form more or 

less cuprous salt. The proportion of cuprous to cupric salt varies 

with the salt, with the concentration, and with the temperature; the 

relations for the sulphate at different temperatures are indicated 

in Table 1.
TRANSPORTATION OF COPPER 19

On account of the possible connection between solutions of copper 

and calcium chloride in the Lake Superior region the writer determined the solubility of cuprous chloride in such solutions at 25° C. 

(experiments 149-155). This solubility is very considerable when 

the concentration of calcium chloride is high, resembling the behavior with ferrous chloride and certain other salts. The data 

on solubility are given in Table 13 on page 69. The determinations 

might well be extended to higher temperatures. That such solutions 

deposit copper on dilution is shown in experiment 52a.

TRANSPORTATION OF COPPER

It is obvious that a number of the solutions referred to in the preceding pages could serve as transporting agents. This alone, however, can not explain deposits of copper; it should also be shown 

that precipitation is a logical final step.

To illustrate this point it may be said that cold dilute solutions 

can carry cupric sulphite or, what is the same thing, cupric sulphate 

and sulphur dioxide. Such solutions deposit copper on warming. 

But as it is necessary to assume that ascending solutions would be 

hot at the start, this particular reaction seems inapplicable to account for the deposition of copper. The same may be said of the 

simple hypothesis of precipitation caused by ferrous compounds 

acting on ascending solutions,37 for ferrous compounds are just as 

likely to occur at depth, and how could the copper have got away 

from them? Another factor is needed to complete such a hypothesis.

The transfer of copper from hot to cold regions is the particular 

problem for consideration here. Such a transfer is possible where 

sulphates are present and rests on the formation and subsequent 

decomposition of cuprous sulphate. The transfer would of course 

be much increased by convective currents. It seems entirely logical 

to assume that some metallic copper is deposited in this way in 

nature, especially where the temperature gradient is high.

The quantity of copper deposited by a single cooling of a hot 

solution of the kind just referred to is rather small. Larger quantities have been obtained in laboratory experiments by maintaining a 

continuous circulation between hot and cold points by means of 

closed circuits or the use of large tubes favorable to convective circulation. In nature there usually can be considered only a single 

movement of a solution or diffusion, in the strict sense, through 

relatively narrow channels.

87 Smyth, H. L., On the origin of the copper deposits of Keweenaw Point: Science, vol. 

3, p. 251, 1896.
20 DEPOSITION OF NATIVE COPPEE

Diffusion is generally held to be too slow a process to account for 

the formation of most ores, but there appear to be circumstances 

under which it might have some effect. It appears difficult to suppose that the Lake Superior copper has collected into lodes by diffusion from small particles extensively disseminated in the traps. 

There is little if any reason to assume any concentration or temperature gradient toward the lodes as they are situated at present, and 

consequently none to assume concentration by diffusion from the 

source mentioned.

If diffusion from depth is postulated it seems necessary to suppose 

that such diffusion was restricted to fairly continuous water-filled 

channels; otherwise the copper would have been more widely distributed horizontally. The amount of diffusion possible in such 

channels can easily be approximated by making certain simple 

assumptions.

For example, let it be assumed arbitrarily that on account of differences in temperature the concentration of copper falls from 40 

grams per liter to a trace within a vertical distance of 2 miles. Let 

it also be assumed that copper in solutions such as we are now considering will diffuse at the same rate as. sodium chloride under" 

similar conditions, that the average temperature was 200° C., and 

that all the copper which diffused was deposited. Then, according 

to the usual laws of diffusion,38 the results are as set forth in the following table:

TABLE 5.- -Diffusion of copper in aqueous solutions

Period

1 day............ ....................... ... ...................

1 year ___________________ . ....................

2,900 years... __________________________

Concentration

gradient

(gramequivalent

per

centimeter)

1

1

0. 000008

. 000008

. 000008

Cross

section

(square

centimeter)

1

1

1

1

1

Temperature

(°C.)

18

200

200

200

200

Quantity

diffused

(gramequivalents)

1.06

5.6

.000045

.016

46

The quantity 46 gram-equivalents, or 3.2 pounds, is sufficient 

copper to yield a column of "ore" half a mile high and 1 square 

centimeter in cross section, carrying 0.6 per cent copper. Under 

the conditions assumed in the calculation it would require 2,900 

years for 3.2 pounds of copper to cross a square centimeter section. 

The richness of the ore deposit would therefore depend simply on 

the size of the available channels as related to the extent in height 

and cross section of the zone in which deposition occurred.

88 Lewis, W. C'., A system of physical chemistry, vol. 1, p. 428, Longmans & Co., 191G.
PRECIPITATION OF COPPER 21

The calculations were made on the assumption of a uniform concentration gradient. Actually the gradient would probably be much 

steeper near the region of high temperature. For this reason the 

period calculated is most likely a minimum, and the maximum may 

be much greater. The process of diffusion would also doubtless be 

somewhat complicated by absorption of the dissolved constituents 

by the channel walls. However, the figures show the probable importance of the phenomenon of diffusion and also the desirability 

of learning as much as possible concerning the concentration of 

copper in various solutions at high temperature.

An actual movement of solutions from the hot region to the cool 

region would serve the purpose just as well as diffusion, provided 

that sufficient copper in solution can be postulated and that appropriate chemical reactions can be given to explain the deposition of copper, and the rate of transfer might be much more rapid 

than it could be by diffusion.

PRECIPITATION OF COPPER

It should be noticed first of all that hot acidified cuprous sulphate 

solutions, which can be considered cupric sulphate solutions saturated with copper, deposit copper on cooling. This fact is shown 

by many experiments in the literature, especially the so-called hot 

and cold tube experiments of Foerster and Seidel,39 Richards, Collins 

and Heimrod,40 and Stokes.41 It is immaterial how the cuprous 

sulphate is produced. For example, it may result from the action 

of cupric sulphate on chalcocite (experiments referred to on p. 36). 

Attempts to obtain the deposition of copper on cooling solutions of 

cuprous sulphate in the presence of cuprous chloride were not very 

successful, however, as basic salts and cuprite formed instead, but 

it is believed that this was because hydrochloric acid was driven 

into the vapor phase during heating of the tubes (experiments. 135& 

and 1355). The experimental conditions could not be framed to 

agree exactly with the probable natural conditions. The effect of 

sodium chloride has not been sufficiently studied to say that it 

wholly prevents the deposition of copper from these solutions on 

cooling, but it at least lessens such deposition greatly, as the 

cuprous sulphate merely forms cuprous chloride and sodium sulphate.

38 Foerster, F., and Seidel, 0., Zur Kenutniss der Elektrolyse von Kupfersulfatlosungen: 

Zeitsclir. anorg. Chemie, Band 14, p. 120, 1897.

40 Richards, T. W., Collins, E., and Heimrod, G. W., The electrochemical equivalents of 

copper and silver: Am, Acad. Proc., vol. 35, p. 124, 1899.

, to Stokes, H. N., Experiments on the solution, transportation, and deposition of copper, 

silver, and gold : Econ. Geology, vol. 1, p. 645, 1906.
22 DEPOSITION OF NATIVE COPPER

Another general reaction less often referred to is the precipitation 

of copper from cuprous salt solutions by dilution.42 Particularly 

solutions containing cuprous chloride dissolved in hydrochloric acid 

or other chlorides, such as calcium and sodium chloride, and in 

equilibrium with copper, precipitate copper when diluted with water 

(experiment 52&). One is strongly tempted to conclude that some 

of the " mass " copper in fissures in the Lake Superior mines may 

have originated in. this way at places where concentrated cuprous 

solutions mingled with more dilute waters.

Small flakes of copper may also be observed as the result of diffusion in capillary spaces or through semipermeable membranes under 

certain conditions.43 If a concentrated solution of alkali sulphide 

is poured into a solution of cupric sulphate with as little stirring 

as possible the floes of copper sulphide become more or less coated 

with metallic copper in the copper solution. This result is no doubt 

due to electrolytic action in which a circuit is formed, so that an 

electric current passes through the precipitate and returns through 

the solution, reducing copper ions to copper at one point and oxidizing 

sulphide ions to polysulphide at another point. Electric activity is 

probably responsible for some of the treelike growths and filaments 

of copper and silver iii nature, but it would be foolish to attempt to 

explain thus all these occurrences. The electrolytic action depends 

on the passage of a current through solutions or moist rocks, and 

hence the oxidations and reductions involved may occur at points 

more or less separated. Many interesting results may be observed 

by anyone who will take the trouble to set up the combinations, 

when pieces of sulphide ore or strips of various metals are suspended 

at the boundary line of two dissimilar solutions, one being oxidizing 

and the other reducing. The difficulty of applying the observations, 

as well as lack of space, however, compels us to pass to the more 

usual phenomenon of chemical reaction.

In addition to the reactions already cited are those involving 

ferrous compounds, which have been repeatedly given in the literature to explain the origin of metallic copper.44

Nearly all chemists who have described the reduction of copper 

salts by ferrous salts leave the inference in the reader's mind that 

only a little copper was obtained in their experiments, but that

«Abegg's Handbuch, Band 2, abt. 1, p. 581, 1908.

a For references, see Wells, R. C., U. S. Geol. Survey Bull. 548, p. 8, 1914.

"Pumpelly, Raphael, The paragenesls and derivation of copper and its associates on 

Lake Superior: Am. Jour. Sci., 3d ser., vol. 2, pp. 188-198, 243-258, 347-355, 1871 

(especially p. 353). Stokes, H. N., Experiments on the solution, transportation, and 

deposition of copper, silver, and gold: Econ. Geology, vol. 1, p. 648, 1906. Fernekes, 

Gustave, Precipitation of copper from chloride solutions by means of ferrous chloride: 

Econ. Geology, vol. 2, p. 580, 1907. Lane, A. C., Keweenaw series of Michigan: Michigan 

Geol. and Biol. Survey Pub. 6, p. 847, 1911. Lindgren, Waldemar, Mineral deposits, p. 

399, 1913. '
23

enough was obtained to show that it might be formed that way in 

nature.

Thus Knop obtained only a mirror on the glass vessel in the course 

of time.45 Wibel 46 could not repeat the experiment, but obtained 

traces of copper at 210° C. Braun obtained a " partial deposition." 

Hunt 47 states, however, that an excess of ferrous chloride converts 

one-third of the copper of cuprous oxide to metal. He also states 

that an excess of ferrous oxide or ferrous carbonate reduces all the 

copper of cupric chloride to metallic copper.

2CuCl2+6FeO=2Cu+2FeCl2+2Fe2O3

The experiment described by Stokes, starting with cupric and 

ferrous sulphate, was repeated by the writer (experiment 14) and 

yielded some copper on cooling, after heating at 200° for about 3 

hours.

The equilibrium involved here is apparently the well-known one 

between ferric and cuprous salts on the one hand and ferrous and 

cupric salts on the other. For example,

FeCl3 + CuCl^FeCL, + CuCl2 (29) 

or as ions

(30)

Eeasoning based on E. M. F. measurements leads to the expression 

for the equilibrium constant

++ **

[Cu+] [F

In other words, the reaction with soluble salts lies very largely 

toward the right, but in neutral or alkaline solutions, because ferric 

hydroxide is much less soluble than ferrous hydroxide, the equilibrium may be displaced toward the left, when a second reaction 

occurs, forming metallic copper, thus :

Fe++ + Cu+ <=±Fe+++ + Cu (32)

The formation of the copper, however, may conceivably occur on 

account of the accumulation of cuprous ions, thus :

2Cu+<^Cu + Cu++ (24)

48 Knop, A., Ueber die Kupfererz-Lagerstatten von Klein-Namaqualaud und Damaraland ; ein Beitrag zur Entwickelungs-Gescbichte der Kupfer-Erze : Neues Jahrb., p. 514, 

1861. (Also especially p. 543.)

46 Biddle, H. C., The reduction of copper by solutions of ferrous salts : Am. Chem. Jour., 

vol. 26, p. 377, 1901.

" Hunt, T. S., Contributions to the chemistry of copper : Am. Jour. Sci., 2d ser., 

vol. 49, p. 154, 1870.
24 DEPOSITION OF NATIVE COPPER

I

The latter reaction is known to be displaced strongly to the right 

on cooling and seems to be the fundamental reaction that explains 

the deposition of copper by cooling alone.

Experiments 159 and 160, referring to equation (32) and also described elsewhere,48 show conclusively that copper is deposited in 

visible crystals when the neutralization necessary to bring about 

reaction (32) in chloride solutions is made to occur slowly, even without any cooling. When sulphates are used no copper has up to the 

present been observed to deposit before cooling begins.

Potassium fluoride appears to have the same effect as alkalinity 

in bringing about the deposition of copper in mixtures of cupric ' 

and ferrous salts, according to Miiller.49

Deposition by ferrous compounds would seem to afford a possible 

explanation for the origin of native copper wherever acid copper- \ 

bearing solutions entered rocks carrying ferrous compounds, where 

the solutions were neutralized to the point at which ferric oxide 

would be precipitated. This explanation is not applicable, how- \ 

ever, where ferric oxide appears to have been reduced and dissolved 

concomitant with the deposition of the copper, as is indicated in the < 

Michigan deposits.

Mellor 50 gives an extensive bibliography referring to the many 

ways in which metallic copper can be formed from its compounds \ 

by reduction with organic substances, metals, and different reagents, \ 

especially at high temperatures, but very few of these ways could ' 

have any application to geochemical processes. Field evidence indi- i' 

cates that carbonaceous matter has precipitated metallic copper in ' 

some natural occurrences, however.

Let us now consider certain ways in which copper may be deposited chemically when some of the compounds involved contain 

sulphur in other forms than sulphate. |

When a dilute solution containing cupric sulphate and an excess 

of sulphur dioxide is heated between .100° and 150° C. beautiful 

crystals of copper are deposited (experiments 46 and 47).

CuS04+SO2+2H2 0=Cu+2H2SO4 (33)

Cupric sulphite yields the same result as the sulphate. If, however, the temperature is raised to a point between 150° and 250° the . copper first formed is more or less converted into cuprous sulphide 

(experiment 61). We have already seen that cuprous sulphide is 

formed when sulphuric acid acts on copper above 200°. Copper,

48 Wells, R. C., Reaction between ferrous salts and cuprous salts: Am. Jour. Scl., 5th 

ser., vol. 8, pp. 428-433, 1924.

40 Mtlller, Erich, Ueber die reducierende und oxydierende Kraft von Eisensalzen: 

Zeitschr. Elektrochemie, Band 14, p. 76, 1908.

60 Mellor, J. W., Inorganic and theoretical chemistry, vol. 3, p. 10, Longmans, 1923.
j PRECIPITATION OF COPPER 25

when heated with a solution of sulphur dioxide at 200° to 230° 

(experiments 48, 59, and 60), is also converted into cuprous sulphide, probably according to the reaction

2Cu+3SO2+2H2O=Cu2S+2H2SO4 (34)

Now in this reaction the presence of ferric oxide decreases or 

, prevents the formation of the sulphide, doubtless because the sulphur dioxide is consumed in reducing the ferric oxide (experiment 

60). This experiment furnished the first concrete evidence found 

* that the presence of ferric oxide could explain the persistence of 

copper, whereas in its absence chalcocite would be formed. Suppose that copper were first formed by any reaction whatever; a subi sequent accession of sulphur dioxide would convert the copper into 

1 .sulphide except where protected by the presence of ferric oxide.

If in the action of sulphur dioxide on cupric sulphate (experi-

; ment 61), there had been an excess of cupric sulphate instead of

sulphur dioxide, the result would have been somewhat different.

Under these circumstances less sulphide of copper is formed in the

interval between 150° and 250°, and copper is the principal solid

product (experiment 124). The products depend on whether sul-

/ phur dioxide or cupric sulphate is in excess: an excess of sulphur

) dioxide forms some chalcocite; an excess of cupric sulphate forms

i chiefly copper, unless too much sulphuric acid is present (experi-

( ment 128).

' These somewhat confusing results are evidently explained by 

differences of temperature, the proportions of the substances taken, 

and reversibility of the reactions involved.

Thus, reaction (33), which proceeds to the right under the conj ditions mentioned, is the reverse of the one given by Pickering for 

i the generation of sulphur dioxide by the action of concentrated 

sulphuric acid on copper.

Again, the reaction written for experiment 58,

; Cu2S + 6H2S04?^5S02 + 2CuS04 + 6H20 (35)

is practically the reverse of that which explains the results of 

 experiment 61.

Finally, if we multiply equation (33) by 5 and add it to equation 

(35) we obtain

Cu2S + 3CuS04 + 4H2 0^5Cu + 4H2S04 (36)

which is the equation that has been found to explain most nearly 

the formation of copper from cuprous sulphide. Equation (36) is 

the reverse of equation (IT), which represents in part the action of 

copper on sulphuric acid. We shall return to this reaction presently.
26 DEPOSITION OF NATIVE COPPEE

HYPOTHETICAL REACTIONS

The hypothesis that native copper has been formed through oxidation of sulphur by ferric compounds, which has been advocated with 

confidence by the geologists of the Calumet & Hecla Consolidated 

Mining Co., is, in simplest terms, that the copper was brought to its 

present position in a solution from which it would have been deposited as sulphide in the absence of hematite, but owing to the 

presence of the ubiquitous hematite the sulphur was oxidized, so 

that metallic copper was deposited instead.

In its simplest form this conception might be expressed by the 

following reaction :

q.) + 5Fe(OH) 2 + 2Cu (37)

The thermal effect of this reaction at ordinary temperature is 

calculated to be an absorption of 38,700 calories, which is not encouraging for the reaction. Supposing the reaction to be entirely 

an ionic one, it may also be written

2Cu+ + S + 6F6++* + 80H- = 6Fe++ + 2Cu + SO" + 4H20 (liquid) (38)

It is possible to evaluate the difference in free energy between 

the factors and products of this reaction by the methods outlined 

by Lewis and Randall,51 and the result is the very favorable diminution of free energy of 253,914 calories when all of the substances are 

present in their standard states. This calculation is based on one 

value not found in Lewis and Randall :

CU+ = CU++ + E-; AF° = 3920

It must be noted at once, however, that the insolubility of the compounds involved greatly alters the situation, and according to the 

computations of the writer, the reaction can not be expected to take 

place under ordinary conditions, except in alkaline solutions and 

when the sulphate concentration is low. *

The details of the calculation are as follows :

If ° 253,914 it is calculated that log K298=186, and K, the 

equilibrium constant at ordinary temperature, is consequently 10186 . 

That is :

[Fe++]» [SO"]

[Cu+]2 [S ] [OH-] 8 [Fe+++]6 = iu ( '

61 Lewis, G. N., and Randall, Merle, Thermodynamics, p. 596, McGraw-Hill Book Co., 

1923.
OXIDATION OF SULPHUE 27

Now [Cu+]2 [S ], the solubility product of Cu2S, is 2.2 x 10~47 , 

according to Triimpler.52

Lamb 53 gives [Fe++] [OH-]2 = 9.8X 1Q-15 , and Miiller 54 gives [Fe+++ ] 

[OH-]3 = 1.1X10-36 .

On substituting in equation (39) the values of [Fe++] and [Fe+++] 

obtained from their solubility products, also 2.2 X 10~47 for [Cu+]2 [S ], 

and solving, we obtain]

K = 2.4X10 178[J^~~] (40)

If [S03 ] is ten-thousandth molal, corresponding to the solubility 

of barium sulphate, we have

For neutrality, K = 2.4 X10188

For 0.00001 N OH", K = 2.4xl0184

For 0.001 N OH-, K-2.4X10180

A comparison of these values with that for equilibrium, K=10186 , 

shows that it is only in the alkaline solutions that the reaction can 

be expected to occur, as already stated. Experiments with alkaline 

solutions, however, did not yie^d any conclusive results on account of 

experimental difficulties. The'conditions of ionization and concentration in "magmatic solutions" above 300° still remain largely 

speculative and might be considerably different from those assumed 

above.

The calculations were not completed until after a good many experiments had been made, which was just as well, perhaps, because it 

was found that a certain amount of copper could be formed, even 

in acid solutions, from the above-named substances by virtue of 

other reactions. Eeaction (37) is merely illustrative, and there are 

several other possibilities.

Instead of preparing solutions of cuprous sulphide of workable 

concentration in the laboratory it was necessary to use the solid. 

Until more is known about the character of the natural solutions 

carrying " copper sulphide " it seems necessary to postulate rather 

large quantities of solution, but this difficulty is clearly not insuperable, because in nature, in many parts of the world, copper sulphide 

has evidently been extensively transported in solution to form ore 

deposits. Obviously experimental conditions can seldom conform to 

natural conditions in all respects.

It appears possible to test the hypothesis under scrutiny by field 

evidence also. A good many localities where it seems to explain the

M Trtimpler, G., Zur Kenntnls des Elektromotorischen Verhaltens von Metallverbindungen mit Elektronenleitung: Zeltschr. physikal. Chemle, Band 99, p. 9, 1921.

M Lamb, A. B., The potential of iron calculated from equilibria measurements: Am. 

Chem. Soc. Jour., vol. 32, p. 1218, 1910.

** Miiller, Erich, Ueber die reduzierende und oxydierende Kraft von Eisensalzen: 

Zeltschr. Elektrochemle, Baud 14, p. 76, 1908.
28 DEPOSITION OF NATIVE COPPER

facts, in addition to those in the Lake Superior district, have already 

come to light. In studying the origin of native copper in the red 

sandstone of Corocoro Steinmann concluded that ferric oxide has 

inhibited the deposition of copper sulphide and favored the deposition of copper.

STEINMANN'S EXPERIMENT

As further support for the conclusion above cited an experiment 

was made by L. Gassner under Meigen's direction, given in Steinmann's article, in which precipitated copper sulphide and ferric 

hydroxide were allowed to act on each other in the presence of 

calcium carbonate.55

Gypsum was formed, and apparently a ferrous compound, but 

nothing is stated to show that metallic copper was one of the 

products.

This experiment has been repeated and studied very carefully 

under different conditions.

It has been found that when the reacting substances are carefully 

freed from soluble sulphates and when air is wholly excluded the 

quantity of sulphur oxidized in a couple of weeks, even at 92°, is 

extremely small (experiment 38). Nor were any better results obtained at 300° under the simple conditions recorded in experiments 

44 and 45. Experiment 50 resembles experiment 38&, with 23 days' 

heating.

The striking changes observed when the solutions had an acid 

reaction, however, as in experiment 38e?, were deemed worthy of 

further study. This phase of the problem is discussed in a later 

part of the paper.

ACTION OF SULPHIDE WATERS ON FERRIC OXIDE

Waters containing sulphides of the alkalies or alkaline earths convert ferric oxide into pyrite under certain conditions,56 at least when 

the sulphides are in excess. Just what the results would be when 

ferric oxide is in excess does not seem to have been investigated with 

care. Some suggestions may be drawn from the experiments described in this paper that involve cuprous sulphide. In the Lake Superior district the sulphides of iron seem to be even rarer than chalcocite, and any hypothesis for the copper that involves sulphur as one

65 Harry Rosenbusch Festschrift, p. 360, 1906.

56 Doelter, C., Synthese und Zusammensetzung des Pyrrhotin : Min. pet. Mitt, vol. 7, 

p. 535, 1885-86. Alien, E. T., and Crenshaw, J. L., Effect of temperature and acidity in 

the formation of marcasite (FeS2 ) and wurtzite (ZnS) a contribution to the genesis of 

instable forms; Microscopic study by H. E. Merwin : Am. Jour. Scl., 4th ser., vol. 38, p. 

427, 3914.
FERRIC OXIDE, COPPER SULPHIDE, AND ACID 29

of the constituents would seem to require that the sulphur must have 

been almost completely oxidized and removed at some stage of the 

process. As will be shown later the most plausible theory to fit these 

requirements is that the sulphur was oxidized to sulphate or other 

soluble compounds and thus removed from the field of action.

Acid solutions in the presence of an excess of ferric oxide oxidize 

even cuprous sulphide, and considerable study has been given to the 

possible relations involved in such combinations.

FERRIC OXIDE, COPPER SULPHIDE, AND ACID

Mixtures of ferric oxide and copper sulphide are attacked by 

acids with results that depend on a number of factors. Ferric oxide 

is apparently attacked first, and further action resembles in many 

respects that produced by soluble ferric compounds. The marked 

tendency of ferric salts to hydrolyze, however, especially at high 

temperatures, is a significant feature of the problem. More and 

more acid is required to maintain appreciable quantities of ferric 

, iron in solution as the temperature rises.

The composition of acidified ferric sulphate solutions that can remain in equilibrium with ferric oxide has been determined over a 

wide range of concentration and temperature by Posnjak and Merwin. 57 The figures given in Table 6 have been interpolated graphically from their data for the purpose of showing the content of 

ferric oxide of such solutions up to 4 per cent S03 .

TABLE 6. Content of ferric oxide of acid ferric sulphate solutions in equilibrium with Fe^Oa, expressed in per cent

Temperature (°C.)

50. 

75......... . _ . ............. ..

100.-. . 

140 

200... 

0.5 per 

cent 80s

«0.2

".I

«.03

a Trace.

 * Trace.

1.0 per 

cent SOs

<>0.5

«.3

".2

d .01

 'Trace.

1.5 per 

cent 80s

"0.9

°.8

'.5

<f.04

".02

2.0 per 

cent SOa

 1.3

= 1.3

».9l

".2

o.M

4.0 per 

cent SOj

02.6

o2.5

»2.1

 >1.2

"A

a Solid phase, FezOs.HsO.

* Solid phase, FezOa.HaOandSFeaCMSOs.Q^O (?). 

c Solid phase, 3Fe203 .4 SOa.9 HzO.

* Fe2 03 .

67 Posnjak, E., and Merwln, H. B., The system Fea03 S03 H20 : Am. Chem. Soc. Jour., 

vol. 44, p. 1965, 1922.
30 DEPOSITION OF NATIVE COPPER

The tendency toward hydrolysis is clearly shown by Table 6.58 

It is evident that a solution of ferric sulphate, of any such acidity 

as shown in the table, which is in equilibrium with ferric oxide at 

a high temperature, will dissolve much more of it on cooling. This 

fact offered a practical difficulty in the experimental work, as solutions in equilibrium with ferric oxide at a high temperature would 

acquire more ferric salt on cooling and thus perhaps dissolve a little 

copper if any had been formed. It seemed best not to attempt to 

work with ferric oxide in excess under such conditions.

The work of Posnjak and Merwin seems to show that the ferric 

salt in equilibrium with a solution saturated with ferric sulphate 

and containing considerable sulphuric acid at temperatures above 

200° C. would probably be the basic salt, Fe2O3 .2S03 .H20; and 

the maximum concentration of iron in the solution would be less 

than 3 per cent of Fe,03 for 10 to 50 per cent S03 . But quantities 

far less than this would oxidize sulphides.

BLEACHING ACCOMPANYING COPPER

The red rock of the Calumet conglomerate and the amygdaloidal 

lodes is noticeably bleached near the copper. Geologists have found 

that the bleached areas contain a lower percentage of total iron, 

especially cf ferric iron, than the near-by unbleached rock.

This is apparently a widespread phenomenon that requires explanation. Either the solutions that brought in the copper were able 

to reduce hematite or the copper has done it; in either event the 

ferrous compounds thus formed have been removed. That the deposition of copper may not invariably accomplish bleaching is suggested by a few occurrences of copper in contact with unbleached 

rock, and by the fact that in one of Stokes's experiments at 200° he 

obtained hematite and copper side by side.59 But it seemed desirable 

to obtain further experimental evidence on this point.

Solids, such as ferric oxide and copper, can of course affect each 

other only through the agency of solutions in contact with them 

that serve as media for the transportation of the reacting elements. 

But ferric oxide is one of the most insoluble oxides known, if we 

judge from its behavior in the fractional precipitation of hydroxides.60 Neutral and alkaline solutions do not dissolve enough of 

it to give a test with thiocyanates, and the writer has found that the 

solution must have an acidity as great as pH=5.5 permanently to 

dissolve enough to respond to the test.

68 See also Wells, R. C., The electrical conductivity of ferric sulphate: Am. Chem. Soc. 

Jour., vol. 31, p. 1027, 1909.

88 Econ. Geology, vol. 1, p. 648, 1906.

w Wells, R. C., The fractional precipitation of some ore-forming compounds at moderate 

temperatures: U. S. GeoL Survey Bull. 609, p. 25, 1915.
FERRIC OXIDE, COPPER SULPHIDE, AND ACID 31

While making the above tests the writer also noted the marked 

tendency of ferric oxide to go into a colloidal solution around pH= 

5.4 (experiment 11). Thomas and Frieden 61 have indicated the 

value pH:=5.0 for the same phenomenon.

Experiments on the modification of aqueous solutions by metallic 

copper were made by Fischer,02 who showed that metallic copper will 

come into equilibrium (equation 19) with a cupric sulphate solution under certain conditions in a few hours. It seemed desirable 

to investigate whether copper could affect ferric oxide in this way.

In neutral or alkaline solution no action could be detected (experiment 5). In the presence of carbon dioxide the results were inconclusive (experiment 39). With ferric sulphate, however, action 

began at once (experiment 12) and the ferric salt was largely reduced with accompanying solution of copper that is, acid solutions 

caused a reaction to occur rapidly.

Some experiments were also made to see if cuprous salts are able to 

reduce ferric oxide. In alkaline or even neutral solutions no significant action could be detected (experiments 2, 3, 4).

The reaction which might be expected here may be written

Fe2O3+2CuCl=FeO+2CuO+FeCl2 (41)

Not only did the experiments fail to show that such a reaction 

occurs to any significant extent, but essentially the reverse reaction 

was found to occur with ease by Hunt,63 who gives the equation

3CuO+2FeCl2=Fe2 03+2CuCl+CuCl2 (42)

Here the cupric compound is only partly reduced to cuprous 

salt, but the whole of the ferrous salt is oxidized to ferric. The 

progress of the reaction is facilitated by the presence of NaCl to 

dissolve the CuCl formed.

Biddle also states that from an emulsion of ferrous and cupric 

hydroxides, after long standing, crystals of cuprous oxide can be 

separated.04

In the presence of acids, on the other hand, noticeable reaction 

occurs (experiment 145), and the results are very similar to those

61 Thomas, A. W., and Frieden, A., Ferric salts as the " solution link " in the stability 

of ferric oxide hydrosol: Am. Chem. Soc. Jour., vol. 45, p. 2529, 1923.

^Fischer, Franz, Die chemiscne Uebertragbarkeit der Metallpotentialle und der chemische Lb'sungsdruck der Metalle: Zeitschr. physikal. Chemie, Band 52, p. 55, 1905. 

Earlier experiments are given by Abel, E., Uber das Gleichgewicht zwischen den verschiedenen Oxydationstufen desselben Metalles: Zeitschr. anorg. Chemie, Band 26, p. 361, 

1901.

63 Hunt, T. S., On, the general metallurgical method of Messrs. Whelpley and 

Storer: Am. Jour. Sci., 2d ser., vol. 43, p. 305, 1867 ; also somewhat more fully in Contributions to the chemistry of copper: Idem, vol. 49, p. 153, 1870.

64 Biddle, H. O., The reduction of copper by solutions of ferrous salts: Am. Chem. Jour., 

vol. 26, t>. 377, 1901.

42S80 25 3
32 DEPOSITION OF NATIVE COPPER

in which copper is used. The soluble ferric salt formed by solution 

of ferric oxide in the acid is rapidly reduced to ferrous salt with 

accompanying oxidation of the cuprous salt.

RESULTS WITH AN EXCESS OF ACID

In order to hasten the investigation a number of experiments 

were tried with ferric sulphate and ferric chloride in place of ferric 

oxide and acid.

The results obtained depended on whether the copper sulphide or 

the ferric salt was in excess.

With the ferric salt in excess, experiment 11, for example, shows 

that the cuprous copper is quickly oxidized even at ordinary temperature. '

Cu2S+2FeCl3=CuCl2+CuS+2FeCl2 (43)

But conjointly with this reaction some of the sulphide sulphur is 

also oxidized to free sulphur, the reaction for which would be

Cu2S+4FeCl3=4FeCl2 -f2CuCl2+S (44)

In experiment 101 sulphur was observed as a product soon after 

the components were mixed and warmed slightly. On heating to 

a higher temperature covellite became visible. The further possible 

reactions involved in this experiment are given with the discussion 

of the analytical results.

In experiment 102, in which Cu2S was decidedly in excess, cuprous 

chloride crystallized whenever the tube was cooled, indicating still 

another reaction:

Cu2S+FeCl3=CuCl-f-FeCl2+CuS (45)

Lastly, there is the possibility of the oxidation of free sulphur 

to sulphur dioxide or sulphuric acid or both. Many hints are found 

in the literature that sulphuric acid may be formed in experiments 

such as those under consideration, especially at 170° or higher. 

Stokes 65 found that at 100° ferric sulphate oxidizes about IT per 

cent of the sulphur of marcasite and about 60 per cent of that of' 

pyrite to sulphuric acid. He found that an excess of pyrite or marcasite reduces ferric salts essentially to the ferrous condition. The 

reaction occurs very slowly in the cold but fairly rapidly at the 

boiling temperature.

Experiment 95 shows that practically no sulphate is formed by 

the action of ferric chloride on cuprous sulphide at ordinary temperature in the course of a few weeks. In fact, Vogt 66 found that 

ferric chloride in acting on chalcocite produced only a trace of sulphate in several years at ordinary temperature.

05 Stokes, H. N., On pyrite and marcasite: U. S. Geol. Survey Bull. 186, p. 15, 1001. 

00 Posepny, Franz, The genesis of ore deposits, 2d ed., p. 677, 1902.
CUPEIC SULPHATE AND COVELLITE 33

. At higher temperatures, however, the production of sulphate 

begins to take place. Experiment 9, at 110°, can be explained only 

by assuming that some sulphuric acid was formed, starting with 

ferric oxide and cuprous sulphide in an acid solution, although all 

the sulphur was not oxidized in 42 hours.

The necessity of this action at high temperatures may be deduced 

from the fact that sulphur reacts with water thus:

4S+4H2 0:=3H2S+H2SO4 (10)

Experiment 100 shows that copper salts are converted to sulphides 

by sulphur in the presence of water at 200°.

Equation (43) therefore represents only an intermediate stage in 

the final result, as indicated more fully in experiment 101; and, on the 

other hand, in experiment 102, in which there was a large excess of 

cuprous sulphide, very little sulphuric acid was formed even at 200°, 

the oxidizing action of the ferric salt being used up in the easier or 

quicker reactions.

The oxidation of free sulphur by ferric sulphate was studied in 

experiments 18, 19, 20, and 21, which show that sulphur reduces 

ferric sulphate very slowly at ordinary temperature, increasing in 

rate up to 200°, at which temperature the action is complete in a few 

hours with concentrated solutions but either incomplete or slower 

with dilute solutions.

Wardlaw and Pinkard 67 state that sulphur is without effect on 

cupric chloride when boiled with it for four hours, except in the 

presence of concentrated hydrochloric acid, when the following reaction occurs :

6CuCl2+S+4H20=3Cu2Cl2+6HCl+H2S04 (46)

Experiment 100 shows that covellite is formed at 200° under certain conditions.

CUPRIC SULPHATE AND COVELLITE

It has been shown conclusively by Zies, Alien, and Merwin os that 

covellite and cupric sulphate react thus:

5CuS+3CuS04+4H2 0=4Cu2S+4H2SO4 (46a)

These authors further suggest that this final result is attained in 

two stages, one of which is

CuS+7CuSO4+4H2O=:4Cu2SO4+4H2SO4 (47)

"Wardlaw, William, and Pinkard, P. W., The oxidizing properties of sulphur dioxide, 

pt. 3: Chem. Soc. Jour., Trans., vol. 121, p., 218, 1922.

68 Zies, E. G., Alien, E. T., and Merwiu, H. E., Some reactions involved in secondary 

copper sulphide enrichment: Econ..Geology, vol. 11, p. 429, 1916.
34 DEPOSITION OF NATIVE COPPEE

and the other is the formation of the less soluble Cu2S at the expense 

of CuS by the cuprous sulphate formed by equation (47), thus:

CuS+CuaS04=Cu,S+CuSO4

As the chemical tendencies that start reaction (47) are of special 

interest in the present problem it is helpful to analyze them somewhat 

more fully. Any general tendencies thus established will be helpful 

in interpreting other similar reactions. Some of the constituents in 

the reaction for example, sulphate ions appear on both sides of 

the equation and are therefore probably of only incidental consequence. On the other hand, copper obviously changes its stage of 

oxidation, one atom of sulphur is oxidized from the sulphide stage to 

sulphate, and all the hydrogen of the water appears as sulphuric acid.

THEOBY OF THE REACTION

The ionic reactions that appear to be involved are :

8Cu++ + S = 8Cu+-fS++++++ (48)

4H20^8H+ + 40~ (49)

S-H-H-++ + 40 = SOr - (50)

Equation (48) is an oxidation-reduction reaction and expresses 

the simultaneous reduction of the cupric copper to cuprous and .the 

oxidation of the sulphide sulphur to sulphate sulphur. The next 

equation expresses the common possibility that the constituents of 

water participate in a reaction as ions. Equation (50) expresses the 

process whereby the somewhat hypothetical sulphur and oxygen 

ions shown assume the more stable form of a sulphate ion with no 

net change of electrification.

On cancellation and addition the three preceding equations give

8Cu++ + S~ + 4H20 = 8Cu+ + 8H+ +S04 (51)

which is the simpler ionic form of equation (46<z) . 

The above changes might also be expressed in the following way :

 =8Cu+ + 4S (52)

3H2S + H2S04 ' (10)

3H2S+H2S04 = 8H+ +3S +S04 (53)

which by cancellation and addition yield the same result as before. 

In this method of interpretation the sulphur is viewed as being 

oxidized first from the sulphide stage to free sulphur, which then 

reacts, with water according to equation (10), the equilibrium 

studied by Bichowsky.
CUPRIC ;SULPHATE AND CHALCOCITE 35

Equation (48) itself probably summarizes more than one reaction, 

as the sulphur is assumed to pass directly from its lowest to its 

highest valence. On account of the presence of cupric ions and 

hydrogen ions in excess it is obvious that the concentrations of S~ 

and O available for reactions (48) and (50) must be very small 

indeed. 

APPLICATION TO CHALCOCITE

The theory developed above, as indicated more particularly by 

equation (51), was found to be applicable to chalcocite as well as 

to covellite, especially at higher temperatures. Thus in experiment 

109, which employed cupric chloride and which was run for 1 day 

at 200°, cupric ions oxidized a small quantity of the sulphide ions 

of the chalcocite to sulphuric acid.- The equation is believed to be:

Cu2S+8CuCli+4H2 0=10CuCl+6HCl+HaS04 (54)

Many other experiments made by the writer show the same 

tendency. For example, in experiment 131 the action was continued 

for 2 days at 250° with complete oxidation of the sulphur.

The reverse reaction, evidence for which was sought experimentally, is clearly more difficult. Thus, in experiment 85 cuprous 

chloride was not oxidized by dilute sulphuric acid in the course of 

one day's heating at 300°. In experiment 134 cuprous sulphide was 

heated 4 days with dilute sulphuric acid with very slight consequences, if any. However, the action of more concentrated sulphuric acid even at 200°, as shown by experiment 58, appears to 

involve the oxidation of both cuprous and sulphide ions as well as 

the reduction of the sulphate, the sulphur from both sources reaching 

the intermediate form corresponding to S02 .

CUPRIC SULPHATE AND CHALCOCITE

RESULTS OF EXPERIMENTS

That some copper appears to be formed when cupric sulphate acts 

on cuprous sulphide at 200° is stated by Zies, Alien, and Merwin 09 

and has been verified by the writer (experiment 66), but as the 

quantity of copper was always relatively small a number of variations in the conditions were tried to see if the yield of copper could 

be increased. The quantity previously obtained might have been 

formed merely on cooling from the cuprous sulphate produced in 

the reaction.

Increasing acidity seemed to lessen the simultaneous formation of 

basic copper sulphate (experiments 67 and 62), but too much acid

89 Op. cit, p. 430.
36 DEPOSITION OF NATIVE COPPER

also seemed to lessen the yield of copper (experiment 64). No 

basic salt was obtained under certain conditions when the experiment 

was performed in a silica tube at 300° (experiment 122). The yield 

of copper improved as higher temperatures were employed (experiment 127) and the greater the proportion of cupric sulphate (experiment 125).

If we attempt to analyze the probable mechanism of this change, 

as was done for covellite, we find that reaction (51) is clearly involved. This reaction must produce 8 Cu+ for every S removed.

If, then, Cu is removed from solution, thus,

2Cu+ =Cu+Cu++ (24) 

the following equation would result:

3CuSO4+Cu2S+4H1O=5Cu+4H2S04 (36)

The same result is obtained without the use of the ionic form of 

expression, thus:

CuaS+80uSO4H-4H2 0=5Cu2SO4H-4H8S04 (5'5> 

5Cu2SO4 :=5Cu+5CuSO4 (55a)

which on addition give equation (36).

The numerous experiments made to establish this reaction (114, 

115, 116, 117, 122, 123, 125, 127, 128) are described in detail at the 

end of this paper and in the accompanying tables. The results 

obtained by analyzing the contents of the tubes after cooling satisfied equation (36) fairly well, but the question arose whether any

of the copper was formed at the high temperatures, In order to

throw light on this question a device was arranged for withdrawing 

the sealed tubes from the furnace for brief examination without 

copling them off appreciably (experiment 136). Although it was 

impossible to examine the tubes much above 250° or at very close 

range, copper could not be observed. On the other hand, its deposition from the resulting solutions on cooling was conspicuous and 

easily seen.

It seems probable, therefore, that the production of copper in the 

tubes is the result of two separate reactions:

Hot: Cu2S + 8CuS04 + 4H20 = 5Cu2S04 + 4H2S04 (55)

On cooling: 5Cu2S04^5Cu + 5CuS04 (55a)

or simply Cu2S04 ^Cu + CuSO4 . (19)

The net result is expressed by equation (36). The cupric sulphate 

is not entirely consumed when chalcocite is present in excess, so that 

obviously an equilibrium is involved, and it would seem that reac-
CUPRIC SULPHATE AND CHALCOCITB 37

tion (55) should be reversible. We know that when copper dissolves 

in sulphuric acid some chalcocite is formed, but experiment 85 gave 

no grounds for believing that cuprous salts will reduce sulphuric 

acid. Equation (19) is known to be reversible, but we have already 

concluded that no copper is produced at, say, 250°. It must consequently be assumed for the present that a small concentration of 

some other product not identified in the analyses, such as hydrogen 

or sulphur dioxide, presumably put a stop to the reaction.

The following possible intermediate reaction may be worthy of 

consideration:

Cu2S+4CuSO4=3Cu2SO4+2SO2 . (556)

Such a reaction might account for the somewhat variable results actualty obtained in verifying reaction (36) and might well be studied 

further. Sulphur dioxide was not identified as a promment constituent in any of the final solutions, but in some of them a slow increase of sulphate seemed to occur on standing after the tubes were 

opened, which may possibly have been due to small amounts of sulphur dioxide, or also possibly to traces of thionic acids which were 1 

not tested for. In experiment 132&, involving the action of copper 

on sulphuric acid, a very small amount of sulphur dioxide was found 

in the final solution.

If equation (55) is reversible, it might be possible to apply the 

mass law to it, which would give, with the usual omission of the 

solids and water, ionization being neglected,

(56) F [Cu2S04]5 [H2 S04 ]*

~ [CuSOJ8

This constant was evaluated roughly by making the assumption 

that the copper that was formed came from the cuprous sulphate 

present when hot, as indicated by equation (19), thus giving a measure of the concentration in the hot solution of cuprous sulphate and 

allowing the cupric concentration to be evaluated by difference. The 

results are shown in Table 7.

TABLE 7. Apparent equilibrium constant for reaction (55) 

[Concentrations in gram-molecules per 1,000 grams of water]

Experiment

114... 

116...... 

117.........-.......-... ..............

1156.... . - 

U5C. ............ .. .. . 

122...... 

123..... 

125....... . 

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

Temperature (°C)

250

250

250

300

300

300

300

300

350

[CuaSO«]

0.118

.134

.098

.21

.28

.36

.21

.82

.63

[HiSO<]

0.56

.52

.48

.59

.63

.31

1.53

.51

[CuSOi]

1.14

1.27

.52

1.09

.95

.75

.81

.28

, * 

8.8X10-'

4.6X10-'

9. 2X10-«

2.5X10-*

8. 1X10-* 5.1X10-4

1. 2X10-*

1.7X10-«

1.9X10
38 DEPOSITION OF NATIVE COPPER

The results in Table 7 indicate that reaction (55) proceeds more 

and more to the right as the temperature rises. By comparing the 

ratios of cuprous to cupric copper in Tables 7 and 1 it seems clear 

that the solutions referred to in Table 7 could not have been saturated with metallic copper at the temperature of heating, but only 

became so on cooling.

As the tubes were generally allowed to cool in the heater, which 

sometimes required several hours, there may have been a slight shift 

of the equilibrium represented by equation (55) from the state 

obtained at higher temperatures, but this shift was evidently not 

enough to prevent the solution becoming saturated with copper on 

cooling. The higher the temperature at which the mixture was 

heated the greater the proportion of cuprous copper, and accordingly the higher the temperature at which the solution became saturated with copper on cooling and the more metallic copper deposited.

It would appear that at some higher temperature, not yet realized experimentally, the solution might become saturated with 

metallic copper in virtue of reaction (55), and then reaction (36) 

would be realizable at that temperature. As previously stated, the 

net result of heating and cooling conforms fairly well with the requirements of reaction (36), in so far as it occurs, as shown by 

Table 12. In experiment 127 about 75 per cent of the cuprous sulphide taken was used up and more metallic copper was formed than 

the cuprous sulphide taken, owing to contribution from the cupric 

sulphate.

EFFECT OF SULPHURIC ACID

With considerable sulphuric acid present in experiment 64 no 

copper could be detected after 2 days' heating at 200°, whereas 

without initial acid in experiment 110 a little copper was formed. 

Experiments 122 and 123 were run at 300° one with no initial acid, 

the other with considerable acid. The copper formed with no initial 

acid (and only slightly more cupric sulphate) was practically twice 

the quantity formed in the other way. Lastly, as shown by experiment 138, sulphuric acid alone with copper reverses the reaction, 

dissolving copper to form cuprous sulphide, cuprous sulphate, and 

cupric sulphate until the concentration of the sulphuric acid is reduced to a definite relation to that of the cupric sulphate.

A few experiments were tried with substances present intended to 

use up the sulphuric acid as fast as formed, but no striking yields 

of copper were thus obtained (experiments 108 and 111).
THEORIES PROPOSED 39 

CUPROUS CHLORIDE AND CHALCOCITE

It appeared to be important to determine whether cuprous salts 

 can oxidize sulphides. No oxidation was obtained by means of 

cuprous chloride at 200° or 250° (experiments 69 and 126). This 

result seems to rule out the possibility of the following reaction, accounting for the formation of metallic copper:

8Cu+ + S~ = 8Cu + S++++++ (57)

In short, as far as the experiments have shown, cuprous salts do 

not oxidize the sulphur of chalcocite. The action is limited to the 

cupric salts.

THEORIES PROPOSED TO ACCOUNT FOR THE DEPOSITION OF NATIVE COPPER FROM ASCENDING 

SOLUTIONS

Owing to the number of conditions that are unknown because they 

have left little if any record, it is difficult to frame a picture complete in all details of the exact way in which native copper has 

formed in nature through changes that can be linked to ascending 

solutions. Certain suggestions can be offered, however, which are 

based on the studies recorded in this paper and supported by them. 

Obviously, these suggestions do not exhaust the possibilities, especially as neither alkaline solutions nor solutions above 300° C. have 

been studied to any great extent. The results are such as could be 

obtained with workable solutions and the apparatus available. They 

are sufficient, however, both to supplement and to modify previous 

theories.

So far as shown positively by experimental laboratory evidence, 

copper can be deposited by solutions assumed to be ascending that 

is, solutions being cooled, diluted, or reacting with surface rocks in 

the ways set forth below.

DEPOSITION BY COOLING

Hot rising solutions carrying cuprous sulphate are gradually 

cooled and as they cool will deposit metallic copper. The presence of chlorides, however, seems greatly to lessen if it does not 

entirely prevent this action.

DEPOSITION BY DILUTION

Solutions carrying considerable cupro.us chloride, by virtue of its 

large solubility in certain brines and other concentrated solutions, 

deposit a small fraction of the copper as free metal on being5 

gradually diluted. This mode of origin may account for some of

42980 25 4
40 DEPOSITION OF NATIVE COPPEE

the " mass " copper in the Lake Superior mines, where cross fissures 

have permitted the intermingling of the concentrated brines with 

more dilute solutions.

DEPOSITION BY FRACTIONAL OXIDATION OF COPPER SULPHIDE

The formation of metallic copper by oxidation of the sulphide, as 

practiced in smelting, has been alluded to early in this paper. Analogous reactions have been given in the literature based on oxidation 

by atmospheric oxygen, cupric salts, and ferric salts. Much of the 

work described in this paper concerns oxidation by ferric iron, or 

rather, inasmuch as that appears to involve the formation of ferrous 

and cupric salts, oxidation by cupric salts. The stages by which 

ferric oxide may be viewed as acting on a sulphide solution are as 

follows:

Hot acid solutions containing copper and sulphur, assumed for the 

sake of simplicity to be equivalent in their chemical activity to acid 

solutions carrying cuprous sulphide, meet ferric oxide. The acid 

is decreased by solution of some ferric oxide. The ferric salt thus 

formed exerts an oxidizing influence, which is, however, at once 

balanced by the reducing action of cuprous sulphide, with the production of ferrous sulphate and cupric sulphate, as indicated by the 

following reactions

Cu2S+5Fe2O3+llH2S04=10FeSO4+2CuSO4+llH2O (60)

If there is available a further quantity of cuprous sulphide in 

solution, its presence, the decrease of acidity, and the formation of 

water, ferrous sulphate, and cupric sulphate would all favor the 

reactions

Cu2S+8CuS04+4H2 0=5Cu2S04+4H2S04 (55) 

Cu2S+3CuS04+4H2 0=5Cu-{-4H2S04 (36)

The second of this pair of reactions may possibly occur at a very 

high temperature but has not yet been realized experimentally. 

However, the first reaction yields cuprous sulphate, which deposits 

copper on cooling. The consumption of cuprous sulphide in these 

ways would obviously leave less of it to be deposited as the solution 

cooled, whereas copper would still be deposited on cooling. Conditions favoring the deposition of copper would be an initial high 

temperature, the reduction of acidity to the point where no more 

ferric oxide would be attacked or the local exhaustion of the ferric 

oxide by solution and reduction as outlined, and lastly cooling and 

.dilution of the initially heated solutions. In this way the sulphur
THEORIES PROPOSED 41

would make its exit as ferrous sulphate. A reaction that embraces 

all the steps mentioned would be:

Cu2S+3Fe2 03+5H2S04r=2Cu+6FeS04+5H2 0 (5Ya)

A qualitative verification of the possibility of such a reaction 

in the presence of an excess of sulphuric acid is recorded in experiment 119.

This is essentially the theory advanced by the geologists of the 

Calumet & Hecla Consolidated Mining Co. to account for the deposition of native copper rather than copper sulphide in the Lake Superior district, where ferric oxide has apparently been the oxidizing 

agent. It appears that laboratory evidence supports this hypothesis 

under certain conditions, the principal experimental limitation being the necessity of working in acid solutions. As it is obvious that 

every natural condition could not be reproduced in the laboratory it 

would probably be unwise to claim that acidity is a necessary condition for the reaction, but it is the only key the writer has found to 

start the chain of reactions noted. In fact, all the theories here proposed permit the use of solutions of at least sufficient acidity to carry 

significant quantities of dissolved copper. It is only fair to state 

that on account of experimental difficulties solutions of mild acidity, 

such as those containing large quantities of carbon dioxide or hydrogen sulphide under great pressure, have not yet been thoroughly 

studied.

DEPOSITION BY NEUTRALIZATION AND REDUCTION

Deposition of metallic copper by the reduction of compounds of 

oxidized copper by hydrogen has been referred to above (p. 7) and 

need only be mentioned here. On the other hand, reduction of copper 

compounds by ferrous compounds is the theory that has long been 

most generally held to account for the formation of native copper in 

the Lake Superior district, without much regard to whether the 

cupriferous solutions ascended or descended. The necessary reactions for this .process, as elucidated by Biddle, Stokes, Fernekes, 

Lane, and others, involve the removal of ferric oxide from solution 

along with the copper, a feature which the present study has purposely wished to avoid, if possible, in accordance with the field 

evidence, which shows that the copper is generally found where the 

ferric oxide has been dissolved from the accompanying rocks. A 

brief discussion of the chemistry of this process accompanies equations (28) to (32).

This mode of forming native copper can obviously not be left out 

of consideration entirely, because of the common association of ferrous and copper compounds, and it might be viewed as an adjunct
42 DEPOSITION" OF NATIVE COPPER

to any theory in which such mixtures are involved. On it Smyth 

based his theory of the deposition of copper from ascending solutions, 

which the writer has ventured to supplement with the processes of 

neutralization and cooling. Neutralization is suggested in order that 

the initial solutions can be assumed to have been not very alkaline, 

so that appreciable concentrations of copper salts and ferrous salts 

could be present in them. If the initial solutions were alkaline in the 

presence of ferrous compounds it is difficult to see how significant 

quantities of copper could have been transported. Cooling is suggested as logical in ascending solutions. But as has been shown under the preceding heading cooling alone will cause the deposition 

of copper from acid solutions containing cuprous and ferrous sulphates. On the other hand, ferrous compounds in acid solutions 

will precipitate copper as the solutions are neutralized, even without 

cooling.

COMPARISON OF THE THEORIES

In applying the theories presented above to various districts the 

essential conditions of each district must be considered and the 

choice made accordingly. The first two theories are so general that 

they might apply almost anywhere, and they take no account of the 

chemical character of the gangue or the rocks penetrated by the 

solutions. The third theory seems to fit the Lake Superior district 

best in that the specific influence of ferric oxide in the rocks is takeu 

into account in determining the character of the copper mineral deposited. The fourth theory is to a certain extent the opposite ot 

the third in that oxidized copper compounds are assumed to be reduced to metal, whereas in the third theory sulphide solutions are 

assumed to have been fractionally oxidized.

Each of the theories assumes a different kind of solution: the deposition of copper is brought about in the first by cooling a cuprous 

sulphate solution, in the second by dilution of a cuprous chloride 

solution, in the third by oxidation of a cuprous sulphide solution, and 

in the fourth by reduction of any oxidized copper compound. Possibly some of these different processes may have gone on simultaneously or successively in some places.

EXPERIMENTS

Experiment 1. Ferric oxide and cuprous sulphide. Initial weights: Cu2S 

0.159 gram, Fe2O3 0.4SO gram, NaCl 1.0 gram, HSO 10 grams. These substances were placed in a tube which was drawn out, evacuated, sealed off, and 

heated in a steam bath at about 92° C. for 13 days. No visible metallic copper 

was formed.

Experiment 2. Ferric oxide and cuprous chloride. About 1 gram of CuCl 

was placed in a tube with 0.1 gram FesOa, 2.1 grams NaCl, and 6 cubic centi-
EXPERIMENTS 43

meters H2O. After evacuating and sealing the tube it was heated at 110° 

C. for 1 hour. The solution was then tested by indicators and found to be 

acid to methyl red but not to methyl orange, indicating a pH value of about 4. 

The iron found in the solution was negligible, not more than 0.005 milligram.

Experiment 3. The same substances were used in experiment 3 as were 

used in experiment 2 with the addition of 2 grams of metallic copper and 

were heated at 150° C. for 5 hours. On cooling, some of the cuprous chloride 

recrystallized, and some thick treelike growths of recrystallized copper had 

apparently formed on the ends of the copper wires. A few scattered flakes of 

copper were also visible at various points in the tube. Not more than a trace 

of iron could be detected in the solution, but a very small amount, equivalent 

to about 3 milligrams FeO, as nearly as could be determined, appeared to exist 

in the ferrous condition in the solid. The color of the solid remained red. 

Repetitions of this experiment did not always yield metallic copper, and the 

crystals of copper that were observed may have resulted from the etching of 

the wires. See experiments 10, 31, 69, 126, 163, 164, and 165 for similar data.

Experiment 4- The same substances as those used in experiment 3 except 

the Nad were heated at 200° C. for 17 hours. The solid remained red; no blue 

color appeared in the solution. The solid was extracted with ammonia, after 

which it was heated with hydrochloric acid; the solution then gave no test for 

ferrous iron with potassium ferrocyanide.

Experiment 5. Ferric oxide and copper. Some clean copper in the form of 

wire was heated with a suspension of ferric hydroxide in a sealed tube at 110° 

for five hours. No change was visible. Heated at 200° for three hours, neither 

copper nor iron could be detected in the solution.

Experiment 6. Several solutions of different hydrogen-ion concentrations, 

buffered by mixtures of acetic acid and ammonium acetate, were shaken 

with well-washed ferric hydroxide, with the following results after long 

standing:

Ferric hydroxide dissolved

PH

4.0

4.6

5.2

5.3

5.4

5.5

5.9

Reaction with KSCN

.....do.... ~

Weak -.. ---------

 ..do... 

Very faint..... _ . _ ..

Character of solution

Clear.

Do.

Slightly colloidal.

Do.

Strongly colloidal.

Slightly colloidal.

Clear.

Experiment 7. Ferric oxide and sulphur dioxide. About 0.04 gram Fe-Oa in 

the form of hydroxide was sealed in a tube with 0.75 gram of NaHSOs and 5 

cubic centimeters H20. Heated to 100° 4 hours; only 0.0023 gram of Fe was 

found in the solution. The solid was not tested.

Experiment 8. Ferric oxide and hydrogen sulphide. About 0.05 gram Fe2Os 

was sealed in a tube with H0S. An immediate darkening occurred. Heated 

15 hours at 100° ; some films of pyrite were formed; the rest of the material remained black FeS.

Experiment 9. Ferric oxide, cuprous sulphide, and sulphuric acid. Initial 

weights: Cu2 S 0.02 gram, Fe20s 0.08 gram, H2 SC>4 1.91 grams=1.03 cubic centimeters, and H2O 3.97 cubic centimeters. Tube evacuated, sealed, and heated 

3 hours at 110°. The solution was then pale green, and some ferric oxide had 

been dissolved. A little ferric oxide remained after 20 hours. After 42 hours 

the residue was inconsiderable but appeared to contain a few droplets of sul-
44 DEPOSITION OF NATIVE COPPER

phur and a few black specks. Found in solution: FeSO* 0.0889 gram, CuSO« 

0.0356 gram. Molecular ratios :

FeS04 formed CuS04 formed 

Cu2S taken ~7<1 ' Cu2S taken ~ 1>8

The CuSO« formed is approximately equivalent to the copper of the cuprous 

sulphide, butj the FeSCh, estimated by titration with KMnCh. is greater than 

that required by both of the following reactions :

(58) 

S +2Fe+++ =2Fe++ +S (59)

Apparently, therefore, some of the sulphur was also oxidized at the expense 

of the ferric salt. As the oxidation of S to 80s by Fe20a would yield 6FeO 

an equation may be written to express the reaction on the assumption that all 

of the sulphur is so oxidized.

Cu2S+5Fe203+llH2SO«=10FeS04-i-2CuS04-rllH2 0 (60)

Actually only TFeSO* was produced under the conditions described, as some 

of the sulphur was left in the free state.

Equation 60 requires the following weights of substances: Cu2 S 1.00 gram, 

Fe203 5.02 grams, HzSO* 6.77 grams.

Experiment 10. Initial weights: CuCl 1.5 grams, NaCl 2 grams, H2O 6 

cubic centimeters. Heated at 150° for 3 hours. At the conclusion of the experiment the tube was tilted while hot from its original position so that the 

solution would drain from the solid previous to cooling, and analysis showed 

that the solution must have contained at 150° about 19 per cent of CuCl and 

16 per cent of NaCl.

Experiment 11. Cuprous sulphide and ferric chloride. Weights taken: 

Cu2 S 0.10 gram, FeCl3 0.532 gram, H2O 50 cubic centimeters, in an atmosphere of C02, at room temperature. After 15 minutes' shaking 5 cubic centimeters was withdrawn and titrated, showing 33 per cent of the iron as ferrous. 

At 17 hours 44 per cent was ferrous. After 3 days 51 per cent was ferrous, 

cupric copper was obviously present in solution, and the solid had a blue tint 

suggesting CuS. The first reaction is probably

Cu2S+2FeCl3=CuCl2 +CuS+2FeCl2 (43)

But as the ferrous salt produced was greater in amount than that which 

corresponds to this reaction some of the sulphide must also have been oxidized 

in addition, possibly as follows :

CuS+2FeCls=2FeCl2 -fCuCl2+S' (44)

Experiment 12. Copper and i'erric sulphate. Initial weights: Copper 0.87 

gram, solution 3 cubic centimeters, containing 0.0765 gram Fe^SCMs and 0.153 

gram EkSCh. Action began immediately, giving a blue solution. Heated over 

night in a steam bath, a slight flocculent precipitate appeared which eventually 

turned dark. The blue color was estimated colorimetrically at about 0.3 normal 'CuSlX. After standing several months at room temperature the solution 

was titrated with KMnO4 , giving apparently 0.0498 gram Fe". Dissolved 

Cu 0.0284 gram. The following reactions were assumed :

(61) 

Cu+Fe2 (SO4) 3=CuSO4+2FeSC>4 (62)
EXPERIMENTS 45

By making a few simple corrections for the Cu' and Fe" content corresponding to equation (61) the following values were obtained for equation (62) : 

Cu"=0.0275 gram, Fe" calculated from Cu" 0.0498 gram, Fe" found 0.0482, 

in good agreement with the written equation. The Cu' content was calculated 

from Cu" by Abel's equation.70 As the ferric sulphate was not all reduced it is 

assumed that an equilibrium was obtained in which the ferrous salt greatly predominated over the ferric.

Experiment IS. Ferric oxide, cuprous sulphide, and sulphuric acid. Initial 

weights: Cu2 S 0.02 gram, FesOa 0.08 gram, H2S04 5 cubic centimeters of 0.5 

normal acid. Considerable reaction in 6 hours at 178°. Some blue grains 

appeared and some sulphur. Solution pale blue. Heated 17 hours more at 

141°; most of the solids dissolved. Tube stood in drawer for 7 months; the 

insoluble flakes were brilliantly colored, purple and lavender, and some ferric 

oxide remained.

Experiment 14. Repetition of Stokes's experiment.71 Initial weights: 

FeS04.7H20 0.5 gram, CuS04.5H20 0.5 gram, KHCOs 0.05 gram in a separate 

tube, dissolved later, H2O 5.0 cubic centimeters. Heated on steam bath 6 

hours; no copper. Heated at 200° for 3 hours. Some copper was visible on 

cooling, also a black mineral, probably hematite, and a red substance, which 

may have been either Fe20a or Cu20.

Experiment 15. Initial weights: Cu2S 0.02 gram, Fe208 0.08 gram, H2S04 

0.06 gram, FeSO4.7H20 0.5 gram. Heated at 186° for 17 hours; no copper. 

Heated at 200° for 20 hours; a few specks of copper around the rim. Heated 

at 200° for 2 days more; considerable copper around the rim. The results of 

further experiments similar to this one are summarized in Tables 8 and 12, 

on pages 66-68.

Experiment 16. Initial weights: Cuprous sulphide 0.08 gram, ferric oxide 

0.02 gram, H2S04 0.01 grarn=0.80 cubic centimeter of 0.256 normal acid. Some 

CuS04 was formed before the tube was sealed. Heated at 250° for 2.5 hours. 

Both FezOn and Cu2 S remained. Green and yellow films were conspicuous on 

the walls of the glass tube, and at one point a few small bright particles of 

metallic copper. The solution had become alkaline. This experiment shows 

that ferric sulphate oxidizes cuprous copper faster than it does sulphur, especially at low temperatures.

Experiment 17. Initial weights: Chalcoclte 0.16 gram, hematite 0.1G gram, 

NaCl 1.0 gram, KHCOs 0.05 gram, H2 0 5.0 gram. Heated at 183° for 17 hours; 

no visible effects; solution strongly alkaline.

Experiment 17a. Initial weights: Cuprous sulphide 0.08 gram, ferric oxide 

0.02 gram, NaHSO4 0.03 gram dissolved after sealing tube, H20 10 grams. 

Heated at 250° for 1 hour; solution was very pale blue, containing a trace of 

Cu but no Fe, and was practically neutral (pH about 7.7). A little cuprite as 

a stain above the solution.

Experiment 18. Ferric sulphate and sulphur. Taken: 2 cubic centimeters of 

a solution containing 0.383 gram FeaCSCUK 0.026 gram H2S04, 0.03 gram S, 

and 1.0 gram H20. Initial iron as ferric sulphate 0.0430 gram. Heated 5% 

months at 92° in a sealed tube. Found 0.0020 gram iron as ferrous sulphate.

Experiment 19. Same as experiment 18 except for dilution to 17 cubic centimeters. Found 0.0025 gram iron as ferrous sulphate.

Experiment 20. Taken: Fe2 (SO4 ) 3 0.383 gram, H2S04 0.026 gram in 2 cubic 

centimeters of solution, S 0.03 gram. Heated at 200° for 17 hours. The ferric

TOAbel, B., tiber das Gleichgewicht zwiscnen den verschledenen Oxydationstufen dessclben Metalles: Zeitschr. anorg. Chemie, Band 26, p. 397, 1901.

71 Stokes,, H. N., On pyrite and marcasite: U. S. Geol. Survey Bull. 186, p. 44, 1901.
46 DEPOSITION OF NATIVE COPPER

color disappeared. Some sulphur remained. A little Fe2 03 and a little FeS* 

formed. The solution was still acid and pale green. Found FeSO4 with KMnOo 

equivalent to 0.359 gram Fe^SOOa. The solution showed only a trace of ferric 

iron with KSCN.

Experiment 21. Same as experiment 20 with 10 cubic centimeters H2 0. Results similar to those of experiment 20 except that more sulphur remainedr 

partly in drops and partly coating the glass beneath the solution. The coating 

was brownish on the inside. Titration showed FeSO4 equivalent to 0.140 gram 

Fe2 (S04 ) 3 , or about 37 per cent of that taken.

Experiment 22. Sulphuric acid and sulphur. Initial weights: H2S04 0.065 

gram diluted to 5 cubic centimeters, S 0.0103 gram, Na2 S04 0.0912 gram. Heated 

at 190° for 18 hours; the solution smelled of H2 S but gave only 0.0012 grain> 

H2 S on analysis.

Experiment 23. Cupric sulphate and sulphur. Initial weights: CuSO* 

0.1596 gram, NaHSO* 0.24 gram, S 0.0480 gram, H2 0 5 cubic centimeters. 

Heated at 210° for 19 hours; much Cu2 S formed and most of the sulphur was 

used up. Only 0.0176 gram copper remained in solution.

Experiment 24- Ferric oxide and sulphur. Initial weights: About 0.215- 

gram Fe(OH) 3 and 0.0215 gram sulphur. Heated at 210° for 5 hours; some 

pyrite formed but was a dirty yellow, not bright. The solution was alkaline 

and free from iron; apparently contained no S02 , but on acidifying precipitated sulphur and consumed KMnO4 equivalent to 0.019 gram FeO. Residueseemed to show some ferrous salt present.

Experiment 25. The purpose of experiment 25 (Steinmann's experiment) Ts 

was to allow precipitated ferric hydroxide and copper sulphide to act on each 

other in the presence of calcium carbonate. After heating at 92" for a week 

the solution gave positive tests for calcium and S04 with (NH4 ) 2C2O4 and 

BaCl2 , but as the ferric hydroxide used was not free from sulphate results are 

not conclusive. The residue extracted with HCl gave a test for ferrous iron, 

but the inconclusiveness of this test is pointed out in experiment 3Sd. The test 

is described further in experiments 36 and 38a.

Experiment 26. Ferric oxide and sodium sulphide. The addition of sodium 

sulphide to a suspension of ferric oxide seemed to produce very little effect at 

first, but on heating a black colloidal solution resulted.

Experiment 27. Ferrous sulphide and ferric oxide. Certain solutions were 

mixed in quantities to yield FeS and Fe(OH) 3 in the proportion of one molecule of each. The mixture was heated at 130° for 16 hours. A white deposit 

appeared and a colloidal greenish solution. Red color was absent. After long: 

standing, the solution was clear, the sediment was greenish, but the coating on 

the tube was white. The tube was considerably attacked. After opening the 

tube the solution was found to contain an excess of Na2S but no SO/'. The 

white deposit was not sulphur but probably residue from the glass. Thiosulphate could not be detected. The green sediment, which did not seem to dissolve in dilute HCl, gave an apparent content of 7.6 per cent of ferrous ironr 

17.5 per cent ferric iron, and a little sulphur. . Possibly it could be termed; 

greenalite.

Experiment 28. The same mixture used in experiment 27 with some CuSO* 

in addition, on heating at 280° for 4 hours yielded a clear solution with a dark 

precipitate. Small specks in the precipitate suggested covellite, pyrite, bornite,. 

and sulphur, but most of the solid was black and amorphous. The tube was not 

much attacked.

"Harry Rosenbusch Festschrift, p. 360.
EXPERIMENTS 47

Experiment 29. Similar to experiment 9. Weights : Cu2 S 0.159 gram, FeiOa 

0.32 gram, H2S04 0.29 gram, H2 0 7 cubic centimeters. Heated at 172° for 

22 hours ; much Fe2 03 was still undissolved ; the solution was blue. Heated 

at about 175° for several days ; the Fe^Oa gradually disappeared, but not theblue color. After several days' heating at 230° the remaining sulphide assumed a blue tint suggesting covellite, and apparently a few specks of copper 

were deposited on cooling.

Experiment SO. Experiment 30 resembled experiment 25 but was inconclusive on account of traces of sulphate in the ferric oxide used.

Experiment 31. Continuation of experiment 10. Weights: CuCl2 .2H2 0 3.5 

.grams, NaCl 3.5 grams, Cu 3.5 grams, H2O 6 cubic centimeters. Heated 3 hours 

at 250° and tilted hot. Much CuCl crystallized. Apparently Cu2 0 was formed. 

Some crystals of copper were visible on the wire near the lower end of the 

tube. An approximate determination of the CuCl presumably dissolved in the 

hot solution gave about 780 grams per liter.

Experiment 32. Action of hematite on chalcocite in fused sodium bisulphate.

Weights : Hematite 0.32 gram, chalcocite 0.16 gram, Na2 S2C>7 1.0 gram. Heated 

to 340° in a glass tube in C02 . SOa was evolved, a little sulphur sublimed, and 

apparently some CuS remained, with much Fe^Os. Considerable CuSOu was 

formed. In the water extract KMnO* gave an equivalent of 0.022 gram FeO.

Experiment SS. Effect of S02 on cupric hydroxide. Five cubic centimeters 

of a saturated solution of SO2 was added to 5 cubic centimeters of a suspension of cupric hydroxide in a tube with COz atmosphere. Red crystals soon. 

"formed, probably CuSO3.Cu2 SO3.2H2O. Heated to 125°; metallic copper deposited. The experiment was repeated several times. The solution was more 

or less blue.

Experiment 84. Experiment 34 was a continuation of experiment 32; thesame weights of the same substances were taken in a sealed tube with C02 . 

Heated at 300° for 1 hour; the red color disappeared. The residue was dark 

but not black. Some magnetite was formed. The solution was found to contain 0.125 gram FeO and 0.0495 gram copper.

Experiment 85. Chalcocite and sulphuric acid. Initial weights : Butte 

chalcocite 0.259 gram, HzSCX (concentrated) 1.0 gram. Heated at 132° for 

36 hours; no significant changes noted. Heated at 250° for 4 hours; somesmall crystals, needles and plates, appeared. On cooling and opening much 

SO2 under pressure was noted ; the residue contained a blue product, probably 

covellite. Some of the solution was sealed up and heated to 150°, with the 

production of a few centigrams of covellite. Pickering's equation for this 

reaction is

Further study is recorded in experiment 58, in which no sulphide of copper 

remained.

Experiment 86. A blank test similar to experiment 25, but omitting CU2S, 

after heating one week revealed the presence of some sulphate in the ferric 

hydroxide being used. For the continuation see experiment 38a.

Experiment 87. Experiment 37 was a continuation of experiment 35, with 

no further conclusions.

Experiment 88a. Steinmanu's experiment, continued from experiment 25. 

Some ferric hydroxide was prepared from sulphate-free reagents and washed 

until it began to become colloidal. Some of the suspension, 20 cubic centimeters=0.15 gram Fe20a, was mixed with 0.05 gram CaCOs. After long digestion at 90° the solution gave no test for sulphate.
48 DEPOSITION OP NATIVE COPPER

Experiment S8&. A mixture similar to that used in experiment 38o., with 

about 0.15 gram of precipitated copper sulphide, which had been washed until 

nearly colloidal, was placed in a tube, which was sealed after the air had been 

thoroughly removed by a vacuum pump. After 17 days' heating at 92° the 

solution gave a trace of S04 equal to only 0.2 milligram.

Experiment S8c. In experiment 38c the same procedure was employed as in 

experiment 386, using pure Cu2S, with practically the same result.

Experiment S8d. The purpose of experiment 38d was to produce considerable carbon dioxide in the tube in addition to the substances named in experiment 38& by introducing hydrochloric acid in a small tube which could be 

overturned later. Weights taken: Fft.0a 0.60 gram, Cu2 S 0.30 gram, CaCOs 0.6S . gram, HCI 0.24 gram, H20 6.0 cubic centimeters, volume of tube 24 cubic centimeters. Heated 4 days at 145°. Results: Iridescent particles which may have 

been copper, Cu20, or Fe2 0a; covellite; anhydrite. The results are largely due 

to the fact that a strong acid was injected into the mixture. Simple tests 

show that such an acid dissolves some of the ferric oxide, forming, say, ferric 

chloride, which is then quickly reduced by sulphides to ferrous chloride. A 

mixture of ferric oxide and copper sulphide extracted with warm dilute hydrochloric or sulphuric acid yields a solution containing soluble ferrous and 

copper compounds. Hence Gassner's conclusion that a ferrous compound was 

formed in the Steinmann experiment is not justified by the evidence, although 

it may be correct, as he extracted the residue with acid.

Experiment 89. A suspension of ferric oxide was sealed in a tube containing a quantity of clean copper wire in the presence of carbon dioxide. After 

being heated 12 days at 92° the residue was tested for ferrous iron by extracting with dilute hydrochloric acid and adding potassium ferrocyanide. A faint 

but distinct indication of ferrous iron was obtained. No ferrous iron could be 

detected in the solution, however. No copper could be detected in the solution. 

Experiment 40. Precipitation of copper by ferrous sulphate. Weights: Cu2S 

0.1 gram, CuSO4.5H2O 0.25 gram, FeSOj.THaO 0.25 gram, Fe2O3 0.1 gram, CaCO3 

0.5 gram, HCI (in a separate container tipped after sealing tube) 0.36 gram, 

HzO 5 cubic centimeters. A little C02 escaped before the tube was sealed. 

Heated at 205° for 2 days; the solution was colorless; much.CaSCh and abundant copper formed. Fe2Os remained.

Experiment 41. Similar to experiment 40. Weights: Cu2 S 0.1 gram, 

 uS04.5H2 0 0.25 gram, Fe2 03 0.1 gram, FeCO3 0.5 gram, HCI 0.30 gram, H2O 6 

cubic centimeters. Heated at 150° for 2 days; the solid was a brilliant mixture of blue, yellow, and red; the solution blue. Heated at 190° for 1 day 

more; no Cu was noted but apparently CuS.

Experiment 4%- Fe2O3 and Cu2 S. Initial weights: Fe2 0a 0.33 gram, CusS 

0.11 gram, H2O 10 cubic centimeters. Tube filled with CO2. Heated at 195° 

for 18 hours; the solution gave no test for copper or iron but a slight test for 

SO*". The solid was extracted with oxalic, acid, which gave a solution showing a little ferrous iron but no copper. Extracting with HCI gave both copper 

and ferrous iron in solution.

Experiment 48. Hematite and Cu2S. Initial weights: Hematite.0.48 gram, 

Cu2S 0.16 gram, H20 5 cubic centimeters. Tube filled with COa. Heated at 

210° for 18 hours. The solution was strongly alkaline. It gave no tests for 

iron or copper but showed S'O/'. The insoluble matter extracted with oxalic 

acid gave a test for copper but:none for ferrous iron. Hydrochloric acid extracted copper but no ferrous iron. The last two experiments require verification by means of blank tests to insure that the reagents were reliable and 

that no oxidation by the atmosphere could have occurred.
EXPERIMENTS 49

Experiment 44. Ferric oxide and cuprous sulphide at 300° C. Initial 

weights: Fe2O8 0.24 gram, Cu2S (prepared by heating copper sulphide in hydrogen at 600° for an hour and found by analysis to be Cu2S) 0.0796 gram, H2 0 5 

grams, and C02 to displace air. The experiment was first tried in glass tubes, 

which were heated in an iron bomb containing water to equalize the pressure 

within the tubes, but the glass tubes were so badly attacked and decomposed 

that the results were useless. Silica tubes were then substituted for the glass. 

At first these were placed in larger glass tubes containing water, as the first 

tubes available were rather thin. However, the alkaline solution formed from 

the glass also attacked the silica. Later strong silica tubes were obtained and 

heated directly in the steel bomb with water to equalize the pressure. Heated 

in silica tubes at 300° for 2 hours; no visible change. Heated at 300° for 3 

days; no visible change, except that some of the silica formed a colloidal solution and sonie was deposited at the bottom of the solution. Cu2 S remained 

unchanged. No sulphate was formed.

Experiment 45. Initial weights: FeA 0.080 gram, Cn2S 0.1592 gram, H20 

and CO2 as in experiment 44. Results same as those in experiment 44.

Experiment 46. Action of sulphur dioxide on cupric sulphate solutions at 

160° C. Initial weights: CuSO4.5H20 0.79 gram=0.2 gram Cu, SO2 .aq. 2.9 

<mbic centirneters=0.19 gram S02 , H20 1.0, air about 5 cubic centimeters. 

Warmed to 100°; spangles of copper became visible. Heated to 160° for 1 

hour; S02 used up 0.14 gram, Cu precipitated 0.125 gram.

Experiment 47. Similar to experiment 46. Initial weights: CuS(X5H2 0 

0.79 gram, S02 0.30 grain in 4.75 cubic centimeters. Heated for 1 hour at 160°. 

Copper estimated still in solution from color, 0.04 gram. Heated at 200° a 

moment, then 19 hours at 160°. Solution still blue, but the copper had a black 

film on it.

Experiment 48. Action of sulphur dioxide on copper. Weights: Copper 

2.9530 grams, S02 5 cubic centimeters=0.305 gram. Heated at 145° for 21 

hours; a few brownish floes suggesting a sulphide. Heated for 21 hours more; 

same result. Heated at 217° for 42 hours; a lot of black sediment. S02 remaining, 0.020 gram; used, 0.285 gram; SO* formed=0.125 gram H2S04. Copper remaining, 2.8434 grams; changed to Cu2S, 0.1096 gram.

Experiment 49- Action of dilute sulphuric acid on copper. A little copper 

was heated with 10 cubic centimeters 0.5 normal sulphuric acid in a sealed tube 

free from air to 200° for 42 hours. No color could be .detected in the solution, 

and the copper remained bright. Analysis gave 0.08 milligram in the solution, 

 equivalent to 8 milligrams per liter. See experiment 138 for results at 300°.

Experiment 50. Resembled 386. Initial weights: Fe20s about 0.15 gram 

in the form of a suspension, CaCOa 0.05 gram, H2 0 25 cubic centimeters. Air 

boiled out. Cu2 S added, 0.15 gram. C02 was passed for 1 hour, and tube was 

sealed without interrupting stream of CO2. Heated at 92° for 23 days. Practically no sulphate formed. Residue warmed with dilute HC1, filtered off; 

the solution showed ferrous iron and cupric salt but no sulphate.

Experiment 51. Experiment 51 was similar to the preceding experiment. 

Initial weights: FezOs 0.08 gram, Cu2 S 0.1592 gram, CaCOa 1.27 grams. HCl 

(in inner tube) 0.92 gram. Volume of tube 26.5 cubic centimeters. Heated at 

about 150° for 17 days; considerable black and some purple precipitate formed. 

No further significant changes were noted during the course of 27 days further 

heating. No metallic copper was observed.

Experiment 52. Action of copper on hydrochloric acid. Taken: 10 cubic 

centimeters 0.5 normal acid, 3.0 grams copper wire, in tube with atmosphere 

of C02. Heated 2 days at 150°, then 17 hours at 165°. Some of the copper
50 DEPOSITION OF NATIVE COPPER

was dissolved and reprecipitated on one end of the wires. The solution, after 

cooling, gave the equivalent of 0.70 gram copper per liter.

Experiment 52a. An apparatus, shown in Figure 1, was arranged by which 

a concentrated solution of cuprous chloride, dissolved in strong calcium chloride, could be diluted with pure water, all in an atmosphere of CO*. Only a 

few drops were mixed every few days by allowing some of the mixture to 

escape while first water and then cuprous chloride entered the long tube in 

which mixing occurred. During the course of six months a noteworthy deposit 

of crystals of copper formed where the first dilution occurred. Cuprous 

chloride was also deposited. At the second point of dilution these phenomena 

were less conspicuous. A long copper wire had been placed in the tube to see

Cu wire 

FIGURE 1. Apparatus used in experiment 52a

If any electrolytic transfer would occur, but the results appeared to be due 

entirely to the dilution of the concentrated solution.

Experiment 58. Copper and mine water from 81st level, Red Jacket shaft,. 

Calumet & Hecla mine, depth 4,900 feet, collected August, 1921. Some metallic 

copper was heated with 20 cubic centimeters of this water at 92° in a glass 

tube for 3 months. Copper found in solution 0.2 milligram, equivalent to 0.01 

gram per liter.

Experiment 54- Experiment 54 was similar to the preceding one, with an 

atmosphere of COz. Heated at 150° for 40 days; 15 cubic centimeters of the? 

solution gave 1.8 milligrams of copper equivalent to 0.12 gram per liter.
EXPERIMENTS 51

Experiment 55. Solubility of copper in calcium chloride brine (equal parts 

 aC!2 and H2O). Heated at 175° for 20 hours and filtered hot; the solution 

gave the equivalent of 0.07 gram of copper per liter.

Experiment 56. Similar to preceding. Initial weights: CaCl2.H20 15.0 

grams, HaO 12.0 grams, Cu 3.0 grams, F&Oa 0.2 gram, Si02 0.2 gram. Heated 

at 200° about 2 days; the solution gave the equivalent of 0.19 gram of copper 

per liter.

Experiment 57. Solubility of copper in the presence of acid and Fe2Oa. 

Initial weights: FeaOa 0.2 gram, HC1 0.1 gram, H2 0 15.0 cubic centimeters; 

copper in excess. Heated at 90° for 3 days; the solution gave 0.042 gram of 

iron in the ferrous state and 0.017 gram copper=1.3 gram per liter.

Experiment 58. Repetition of experiment 35. Initial weights: Cu2S 0.250 

gram, H^SO* (boiled) 1.50 grams. Heated in an evacuated sealed tube at 200° 

for 4 days; the Cu3S disappeared. Volume of tube 9.3 cubic centimeters. 

Pale-blue crystals, probably a sulphate of copper, perhaps CuSO<.3H2O, formed. 

Opened under iodine solution, the S02 formed was estimated by titration to be

0.32 gram after allowing for the iodine liberated from the KI by the CuSO<. 

The following equation is suggested:

Cu2S+6H2SO4=5SO,+2CuS04+6H2O (35)

This demands 0.50 gram SOz calculated from the weight of Cua S taken, 

 which is somewhat larger than the S02 found. A little S02 may have escaped 

when the tube was opened, however, as the pressure was considerable, and 

some iodine solution had to be added while the S02 was being absorbed, 

whereas the iodine solution should have been in excess. .

Experiment 59. Similar to experiment 48. Initial weights: Cu 1.999 grains, 

S02 2 cubic centimeters=0.l830 gram, H2O 3.0 cubic centimeters. Heated 4 

days at 225°; considerable Cu2S formed. S02 remaining, 0.029 gram; SO*" 

formed, 0.0770 gram. Apparently 0.0009 gram of sulphur was also formed.

Experiment 60. Similar to preceding, with the addition of 0.46 gram Fe=O.i. 

Heated for 4 days at 225°; only a little Cu2S formed. No SO2 remained. 

FeSO4, 0.214 gram. Sulphur, none. Solution neutral.

Experiment 61. Sulphur dioxide and cupric sulphate. Initial weights: 

CuSO4.5H20 0.100 gram, S02 5 cubic centimeters=0.158 gram, a little C02*. Volume of tube 9.6 cubic centimeters. Heated at 225° for 4 hours, the copper first 

deposited became coated with Cu2 S'. Heated at 225° for 5 days, much Cu2S was 

formed and in one end of the tube droplets of sulphur. Sulphur found, 0.0084 

gram. No H2S or copper. Trace of SO2. Reaction acid. SO4 found, 0.2082 

gram. Two reactions probably occur:

5S02 +2CuS04 +6H2(M6H2S04 +Cu2S 

and

8S02 +2CuS04 +8H20=8H2S04 +Cu2S+S

Experiment 62. Action of cupric sulphate on cuprous sulphide. Initial 

weights: Cu2S 0.080 gram, CuS04.5H2O 0.500 gram, H2SC>4 0.200 gram, H2O 5 

<rubic centimeters. Heated at 230° for 1 day; showed numerous flakes of 

copper on cooling. The second day a very small quantity of an insoluble white 

salt appeared which was observed in several similar experiments and is believed to be formed by attack of the glass.

Third day, no appreciable change. The results of analyses of the solid 

residue and final solution obtained are given in Tables 8 and 9. These determinations were made by Prof. Glenn V. Brown, of Bucknell University, to
52 DEPOSITION OF NATIVE COPPER

whom the writer wishes to express his appreciation for the attention given by 

him to possible sources of error in the analytical work. Professor Brown also 

analyzed the products obtained in experiments 69, 101, 102, 104, 108, 110, 114, 

115, 116, 117, 121, 123, 125, 127, 128, 132&.

Experiment 68. Qualitative experiment. A mixture of CuS with a solution 

containing FeSCX, CuCl2, and HC1 was heated at 130° for 15 hours. Cu-S and 

Cud were formed. No further change was noted after heating at 200° for 

two days.

Experiment 64. Experiment 64 resembled experiment 62 except in the use 

of 0.69 gram of HaSCX. Two days' heating at 200° produced no metallic copper. 

Four days' further heating some months later, at 300°, yielded considerable 

copper on cooling.

Experiment 65. Similar to experiment 62, but with CaClz present. (See 

Table 8.) Heated at 200 9 for 1 day; no copper formed, but the color of the 

solution disappeared, and CuCl may have, been precipitated in the CaS04 

present.

Experiment 66. Action of cupric sulphate on cuprous sulphide, continuing 

experiment 62. Initial weights: Cu=S 0.080 gram, CuSO4.5H2O 0.250 gram, 

HzO 5.0 cubic centimeters. Heating at 150° two hours caused the precipitation of considerable basic salt. Heated at 200°; no copper was observed the 

first day, but some appeared the second day, also possibly Cu2O. The third 

day there was more Cu2O, and all the other constituents were still present. 

The same result the next day, possibly CuO in addition. Some sulphide was 

still present. The tube was accidentally broken at this stage, and contents 

lost.

Experiment 67. Similar to experiment 66, except with 0.500 gram CuSO<.- 

5H30, and 0.100 gram H2 S04 in a total volume of 5 cubic centimeters. Heated 

at 200°. A little copper and basic salt appeared the first day on cooling; somewhat more copper the second day. After 9 days there was no further visible 

change. The crystals of copper were easily observable with a lens.

Experiment 68. Copper introduced with the other reagents. Initial weights: 

Cu2 S 0.08 gram, CuS04.5H2O 0.50 gram, H2 S04 0.23 gram, Cu 0.066 gram, H2O

about 5.0 centimeters. Heated at 200° for 1 day; considerable copper appeared

as beautiful crystals, on cooling, on the surface of the solution. Heated 24 

hours "more; even more copper was recrystallized and the wire was about half 

gone.

Experiment 69. Will cuprous salts oxidize cuprous sulphide? Initial 

weights: Cu2S 0.08 gram, CuCl 0.20 gram, Cu 0.10 gram, H20 5.0 cubic centimeters. Heated 6 days at 200°; a little Cu20 was formed and a little copper 

was apparently dissolved and reprecipitated on the glass and throughout the 

solution. No S<V' was formed. For the results at 250° see experiment 126.

Experiment 70. Action of H2S04 on Ag2S. Initial weights: Ag2 S 0.50 gram, 

H2 S04 (concentrated) 2.0 cubic centimeters, H2S04(1-1) 0.4 cubic centimeters. 

Heated at 92°, some gas was formed. Heated at 200° for 20 hours; some Ag2 S 

still remained. The tube showed the presence of SOa on opening. The solution 

gave a precipitate of Ag2SO4 on dilution. Found in the solution: S02 0.11 gram, 

Ag 0.1772 gram. Possible reaction :

Ag2S+4H2S04=Ag2S04+4SO2+4H2 0 (65)

Experiment 71. Uncompleted.

Experiment 72. Will FeSO4 act on Cu2S to produce copper? Weights: 

FeSO4.7H3O 0.50 gram, Cu2S 0.02 gram, H2 0 4.78 cubic centimeters. Heated
EXPERIMENTS 53

at 200° 6 days; no significant change was noticed. Some months later the tube 

was heated 3 days more at 250° without any results.

Experiment 73. Action of dry cupric chloride on cuprous sulphide. Weights 

taken: CuCU 0.270 gram, Cu2S 0.159 gram. Sealed in a vacuum and heated to 

500° C. for 20 minutes. The mixture melted. Some sulphur sublimed into the 

ends of the tube. The melt was dark in .the center and brown around the edges. 

Heated again 30 minutes. The melt was clear to brownish black in transmitted 

light. The sulphur found by extraction with CS2 was 0.0070 gram. According 

to this result only part of the materials reacted, and on dissolving the residue 

in ammonia much Cu2S appeared which had not been attacked. The ratio 

of copper to chlorine in the soluble part showed that about 36 per cent of the 

copper was in the form of cuprous chloride. These findings support the view 

that as far as the materials reacted the reaction was

2CuCl2+Cu2S=4CuCl+S (66)

Experiment 14- Action of sulphur dioxide and hydrogen on copper at 900° C. 

About 10 grams of copper in the form of a bundle of small wires was heated 

gradually to 900° during the course of 80 minutes in a stream of gas consisting 

of about 2 volumes hydrogen to 1 volume sulphur dioxide. Water was formed, 

but no hydrogen sulphide could be detected in the outgoing gas. After cooling 

and opening the tube the copper was seen to be converted into beautiful crystals of the sulphide for about a .seventh of its length. The crystals were black, 

blue, brass-colored, lavender, and pale green, very well crystallized, and had 

apparently been very nearly if not wholly melted. A little bunch of crystals 

terminated the end of each wire.

Experiment 75. Will S02 convert Cu2 S to copper? Weights: Cu2 S 0.02 

gram, SOa as dry gas 0.10 gram. Volume of tube 9 cubic centimeters. Heated 

at 92°, 150°, 200° for 1 day each; no copper was formed.

Experiment 76. Will S0a reduce Fe2 03 ? Weights: Fe2 03 0.020 gram, SOZ 

as dry gas 0.100 gram. Heated a day at 92°; no change was visible. Heated 

a day at 150°; the color turned to buff. Heated a day at 200°; no further 

change. On opening several months later S0a was still present under .pressure, 

but the Fe203 was not all reduced. In fact, on washing out the S02 with 

H2O, dissolving in HaSCX, and titrating with KMnCh only 0.0020 gram FeO 

was found.

Experiment 77. SO2 and Cu2S with Fe2O3 . Weights: Cu^S 0.020 gram, Fe2O8 

0.04 gram, SOz 0.024 gram, H2O 5.0 cubic centimeters. Heated at different temperatures, finally at 250° for 6 days; no copper was observed.

Experiment 78. Similar to experiment 77. Weights: Cu2 S 0.02 gram, 

FfeOs 0.13 grain, S0a 0.04 gram, FeSC>4.7H2O 0.5 gram, H2O 5.0 cubic centimeters. Heated at 200° 2 days; no copper was observed. The solid remained red.

Experiment 79. Action of liquid sulphur dioxide on cuprous sulphide. Half 

a gram of S02 was condensed into a tube, previously evacuated, containing 

0.02 gram CU2S, and sealed off. No action could be observed on standing 

or after the tube was heated 24 hours at 65° C.

Experiment 80. Action of liquid dioxide on ferric oxide. Ferric oxide was. 

also treated as described in experiment 79. The red color turned slightly 

paler, somewhat buff, but by no means disappeared.

Experiment 81. Attempt to increase the action of Cu2S by increasing its 

solubility. Initial weights: Cu2 S 0.02 gram, KCN 0.05 gram, Fe^Os 0.04 gram, 

H2O 5.0 cubic centimeters, SiO2 0.02 gram. No change resulted after heating at
54 DEPOSITION OF NATIVE COPPER

 92° for a day. On heating one day at 200° a dirty black precipitate formed, 

the Fe2O8 was no longer visible, and after long standing a few extremely faint 

Indications of metallic copper were observed. On heating to 250° a year 

later the tube exploded. Old tubes are apparently considerably weaker than 

new ones.

Experiment 82. Behavior of colloidal Cu2 S. Some CuCl was precipitated

 with H3 S. After washing by decantation about 8 times with water containing 

H2 S a colloidal solution resulted. Fe2O3 was precipitated and washed until it

 was more or less colloidal. (See experiment 6.) Equal volumes of these 

two colloidal solutions were mixed and sealed in an atmosphere of CO2 , 

and the mixture was heated at 92° for 15 days. A green sediment formed, 

and the solution cleared completely. This result can probably be ascribed 

to the action of the C02 and also to the coagulating effect of one colloid 

on the other.

Experiment 88. Initial weights: Fe2 03 0.0028 grams, colloidal Cu2S 15 cubic 

centimeters. Heated in CO2 at 200° for 1 day; the red oxide disappeared 

and a few black floes were formed. On further heating the tube broke.

Experiment 84. Action of colloidal Cu2 S on Fe2O3 . A long tube in the 

form of a siphon 6 millimeters in diameter was filled with hematite and 

Fe2Os, and through it was caused to pass very slowly a colloidal solution of 

Cu2S, which was protected by an atmosphere of CO=. In the course of some 

months certain stains that suggested chalcopyrite, bornite, covellite, and, it 

is barely possible, copper resulted, but the amounts were extremely small 

and could not be readily observed. The most noticeable effect was that the 

colloidal Cu2 S was precipitated and retained by the hematite, so that the 

outflowing solution was clear.

Experiment 85. Action of cuprous salts on sulphuric acid. Initial weights: 

CuCl 0.5119 gram, H2 S04 5 cubic centimeters of 0.5 normal acid. Heated at 

300° for 1 day; the resulting solution was not noticeably blue, and therefore 

no significant quantity of cuprous copper could have been oxidized. More or 

less cuprite formed on the inner glass tube in which the CuCl had been introduced, showing the marked attack that the solution made on the poorer grade 

of glass. From the fact that no cupric salt was formed it may be concluded 

that no sulphuric acid was reduced.

Experiment 86. Similar to experiment 84. Fifteen grams of commercial

 cuprous chloride was placed in a liter bottle with about 500 grams of copper 

shot. The bottle was filled with distilled water previously freed from air 

and shaken about an hour, in order to bring the copper salt into equilibrium 

with copper, thus making it chiefly cuprous chloride. The precipitate which 

formed was allowed to settle out. The resulting solution was decanted into 

another bottle and saturated with hydrogen sulphide. A colloidal suspension

 of cuprous sulphide resulted which remained suspended indefinitely. When 

the excess hydrogen sulphide was displaced by passing a stream of carbon 

dioxide through the solution for an hour, the stability of the suspension was 

lessened to such an extent that it showed a zone of settling in the course of a 

few days. When some of the suspension containing carbon dioxide was 

allowed to flow slowly through a tube filled with coarsely powdered hematite, 

the resulting solution was entirely clear. It contained principally ferrous 

bicarbonate, not more than a trace of copper,, and no sulphide or sulphate. 

Indistinct tarnishings on the hematite seemed to resemble chalcopyrite in 

places and metallic copper in other places, but, as in experiment 84, these 

products could not be positively identified. The experiment was terminated 

by stoppage of the tube caused chiefly by coagulation of the suspension.
EXPERIMENTS 55

Experiment 87. The experiment was started with 2.0 grams of CuCl, and 

procedure was as in experiment 86. The sulphide was precipitated by exactly 

200 cubic centimeters of H2S gas. The Cu2S that formed flocculated to such 

an extent that the solution was not used further.

Experiment 88. Not completed.

Experiment 89. Fifteen grams of commercial CuCl was treated as in experiment 86 until the solution was saturated with H2S. This solution remained in 

complete suspension for 4 days, when some suspended ferric oxide was mixed 

\A-ith it. The suspended particles began to flocculate at once, and the solution 

became gradually clear. A few shiny yellow scales resembling chalcopyrite 

appeared in the precipitate, but these were not separated or positively 

identified.

Experiment 90. Some copper wire was heated in a current of dry S0a at 

500° for half an hour. Scarcely any tarnish appeared on the copper. When 

the S02 was moist, however, the copper was well blackened in half an hour.

Experiment 91 Action of air on Cu3S at 500°. A weighed quantity, 0.1000

gram, of Cu2S gained 0.0469 gram after heating 3 hours at 500°. The 

product, based on a determination of its content of sulphur, was CuO.CuS04. 

Sulphur found, 13.2 per cent as demanded by the written formula. Soluble 

in dilute HC1.

Experiment 92. Carbon monoxide and Cu2S at 500°. Apparently no action.

Experiment 93. Attempt to prepare colloidal cupric sulphide. A solution 

containing 1 gram of CuCl2.2H2 0 in 500 cubic centimeters was mixed with a 

solution of 0.187 gram of Na2S in 500 cubic centimeters. After 2 days a 

considerable quantity of black copper sulphide had precipitated, and the 

solution above was greenish yellow.

Experiment 9Jt. Similar to experiment 93. Three solutions were prepared, 

each containing 1 gram of CuCl2.2H2O in a liter. Hydrogen sulphide gas was 

added to each, 38, 76, and 152 cubic centimeters, respectively. Precipitates 

began to coagulate in all the bottles very soon after the gas was introduced 

and settled out completely in the course of a few days. The results were 

not comparable with those in experiment 86.

Experiment 95. Action of ferric chloride on cuprous sulphide at ordinary 

temperature. Initial weights: Cu2 S 0.500 gram, FeCls 0.0500 gram, HC1 0.005 

gram, H20 50 cubic centimeters. The air was pumped out of the solution and 

replaced with carbon dioxide. The first day the yellow color was replaced 

by a blue-green color. The addition of KSCN gave a red color and a white 

precipitate, possibly CuSCN. No sulphate had formed. At the end of a week 

the iron in solution was about 89 per cent in the ferrous state, as shown by 

titration with KMnO<. After 52 days a similar result was obtained, and still 

practically no sulphur had gone to the sulphate condition.

Experiment 96. Ferric oxide and cuprous sulphide. Similar to experiment 

9. Weights: Fe208 0.192 gram, Cu2S 0.063 gram, H2S04 0.222 gram, H2O 0.208 

cubic centimeter. Heated at various temperatures up to 300°; no copper was 

observed. The quantity of H2O was too small to be thoroughly in contact with 

all the substances.

Experiment 97. Action of H2SO4 on copper. Acid of specific gravity 1.42 

developed no visible color on heating a week with copper at 90° C.

Experiment 98. Action of FeCla on Cu2S. Weights: CiuS 0.100 gram, FeCL 

0.022 gram, H2O 5.0 cubic centimeters. Heated at 200° for 17 hours; some 

Fe208 was deposited in thin flakes, ferrous salt was produced equivalent to 

0.0067 gram FeO, also a little SCV', and a possible trace of copper. The 

weakness of this experiment is that the Fe^Os becomes more soluble on cooling,
56 DEPOSITION OF NATIVE COPPER,

and the ferric salt thus formed would tend to dissolve any copper precipitated, 

if the Cu2 S did not reduce the ferric salt fast enough.

Experiment 99. Initial weights: Cu2 S' 0.100 gram, FeCl3 0.022 gram, HC1 

0.175 gram, H2O about 5 cubic centimeters. Heated at 200° for two days; beautiful black crystals formed, probably recrystallized Cu2S, not more than traces 

of SO/' or sulphur. Ferrous iron in solution equivalent to about 0.006 gram 

FeO. Cupric and ferric salts present. For repetition .see experiment 101.

Experiment 100. Sulphur and cupric chloride. Initial weights: CuCl2.2H2O 

0.85 gram, S 0.64 gram, H2O 5.0 cubic centimeters. On heating at 200° for 17 

hours the blue color of the solution had disappeared. After three days' heating 

there was abundant copper sulphide, blue, lavender, purple, and black.

Experiment 101. Ferric chloride and cuprous sulphide. Initial weights: 

Cu2 S 0.240 gram, FeCU 0.325 gram, HC1 0.200 gram, H2 0 to a volume of 5 

cubic centimeters. The yellow color began to change to green immediately 

on mixing, and on warming to 90° sulphur appeared in the course of an hour. 

Heated at 200° for two days; the solution was still slightly yellow, the solid 

somewhat purple and blue, suggesting covellite. After the solution was heated 

four days more there was no further visible change except an increase in the 

quantity of covellite. Analysis then showed in the precipitate Cu 0.0946 gram, 

S 0.0343 gram, and, in the solution, SO/' 0.0378 gram. The results can not 

be reconciled with any simple single reaction. It is believed that all the following reactions occurred either simultaneously or in succession.

CiuS+4FeCl«==4FeCli+2CuCla+S (44)

Ou.S+2Fe01tf=CuCl,+2Fe01s+CuS (43)

4S+4H2 0=3H2S+H2 SO4 (10)

CuCl2+H2S=CuS+2HCl (67)

Or, by combining (44), (10), and (67) to eliminate sulphur one obtains

4Cu1S+16FeCU'+4H3 0=16FeCU+5CuCU+H,SO«+3CuS+6HCl (68) 

Analysis of the products seems to show that about half of the precipitate 

was unchanged Cu2 S, so that on this basis 0.175 of the original Cu2 S reacted. 

The requirements of equation (68) are then as follows:

Quantities required for equation (68) and quantities found

Cuas CuS-.... ...--._-___.._.-_...._....._._._.._._.._.__._.___....._......_....

S as CuS..... .--...._.__.._...-.____.._.___....._.______.._......._._..-.._.....

S0<" in solution ___ ........ _ _ ...--. ...... ___ ..... _ ..... _ ....... ...

Calculated

Oram 

0.053

.026

.026

Found

Gram 

0.043

.021

.038

These results indicate that another reaction must have taken place to a small 

extent, forming H2 S04 without the simultaneous production of CuS. This reaction was probably some such reaction as the following:

Cus S+10FeCl3 +4H,0=2CuCl,+10FeCl.+H»SO«+6HCl (69) 

Adding equations (68) and (69) we obtain (quite arbitrarily)

6CiuS+26FeCU+8H,O=26FeCh+7CuCl«+2HiSO«+3CiiS+12H01 (70)

Assuming that 0.175 gram of Cu2 S reacted we have by equation (70) the 

quantities shown in the following table:
EXPERIMENTS 57 

Quantities required for equation (70) and quantities found

 Cu as CuS... ____ ......... ___ . ____________________ . __ .

S as CuS.......-.-.-.--..-----..-------.----.--.-.-.. ......... _ ..............

S(V' in solution _ . ______ . _____________ . __________

Calculated

Gram 

0.042

.021

.042

Found

Grain 

0.043

.021

.038

Apparently the net result is approximately a combination of equations (68) 

and (69).

Experiment 102. Action of Cu2 S in excess on FeCL. Initial weights: Cu2S 

0.300 gram, FeCls 0.150 gram=1.41 cubic centimeters of stock solution, HC1 

0.195=1 cubic centimeter of stock solution, H2O in addition 2.5 cubic centimeters. Heated at 200° for 1 day. Some covellite was formed, also cuprous 

chloride, which separated on cooling. The solution was still very pale yellow. 

Probable reaction:

Cu2S+FeCl3=FeCl2+CuCl+CuS (45)

Heated 5 days longer at 200°. The covellite became gradually less noticeable. 

The yellow element in the color of the solution did not seem to disappear entirely. Final result: Total sulphur in precipitate, 0.0592 gram (taken as Cu2 S 

0.0603 gram) ; SO*" in solution 0.0033 gram.

Experiment 103. Initial weights: Cu2 S 0.16 gram, CuCl 0.50 gram, CuCU 

0.12 gram, FeCl2 0.033 gram, HC1 0.39 gram, H2 0 about 4.0 cubic centimeters. 

Heated at 150° for 1 day and at 200° for 1 day; no visible changes. When 

heated at 250° the tube exploded.

Experiment 104- Copper and sulphuric acid. Initial weights: Cu 0.5 gram, 

H2S04 0.50 gram, Hi>0 2.85 cubic centimeters. Volume of tube 12 cubic centimeters. Heated at 250° for 3 days. Some of the copper recrystallized on 

cooling; in fact, only a little of the original wire remained. Cu2S' and CuSO4 

were formed. Heated 2 days more; the copper mostly recrystallized on cooling. Analysis of products is given in Tables 9 and 10. From later work discussed on pages 15 and 37 it is inferred that theoretically all the copper taken 

in this experiment should have dissolved at 250° that isj if the sulphuric acid 

was not too greatly reduced by the reaction to be effective.

Experiment 105. Effect of Cu2O. Initial weights: ChbS 0.16 gram, Cu2 0 

0.429 gram, H2 0 5.0 cubic centimeters. Heated at different temperatures, 

finally 2 days at 250°; no visible copper.

Experiment 106. Same as experiment 105 with 0.05 Fe2Oa in addition. 

Heated finally 4 days at 250°; no visible copper.

Experiment 101. Cu2S and Fe2 (S04 )s. Weights: Cu2 S 0.30 gram, Fea (So4 ) 8 

0.192 gram, H2SO4 0.513 gram, H2O 5.0 cubic centimeters. Heated at 150° for 

2 days; no visible change except that the solution turned blue. Heated at 250° 

for 2 days; no copper deposited on cooling.

Experiment 108. Cuprous sulphide, cupric sulphate, and cuprous oxide. 

Initial weights: Cu2S' 0.08 gram, CuS(X5H2O 0.38 gram, Cu2O 0.28 gram, H2O 

5 cubic centimeters. Heated at 200° for 1 day; a little copper formed, solution 

still blue. Heated another day; a little bright-green basic salt formed. The 

precipitate was now analyzed. Sulphur found 0.0151 gram, originally taken 

as cuprous sulphide 0.0160 gram, which shows that a little of the sulphur of 

the cuprous sulphide had been oxidized.

The experiment was repeated, using Cu2S 0.106 gram, CuS04 .5H2O 0.50 gram: 

otherwise as above. Heated to 300° for 4 days. A little copper and a green
58 DEPOSITION OF NATIVE COPPEB

basic salt were formed. Both Cu2 S and Cu20 remained. The solution was 

moderately blue. Two years later no basic salt was visible; CuzS, Cu, and! 

Cu2O remained.

Experiment 109. Cupric chloride and cuprous sulphide. Initial weights: 

Cu2S 0.08 gram, CuCl2 .2H2O 0.26 gram, H2O 5.0 cubic centimeters. Heated at 

200° for 1 day; found in solution as sulphate 0.0027 gram=0.0080 SO4 ". About 

17 pet cent of the sulphur in the chalcocite was oxidized by heating at 200°" 

for 1 day. Cuprous chloride was suspected as another product of the reaction.. 

(See further experiment 131.)

Experiment 110. Oxidation of chalcocite by cupric sulphate. Initial weights t 

Cu2 S' 0.08 gram, CuS04.5H2O 0.37.5 gram, H2O 5.0 cubic centimeters. Heated 

at 200° for 1 day; a very little copper formed, also a green basic salt. Heated:

I day more. Analyzed. Found 0.0160 gram of sulphur in precipitate, identical 

with sulphur taken as Cu2 S. Practically no acid in the solution. The reaction 

had scarcely started.

Experiment 111. Cuprous sulphide, cupric sulphate, and sodium acetate. 

Initial weights: Cu2S 0.08 gram, CuS04.5H2O 0.50 gram, NaC2H3O2 .3H2O 

0.72 gram. Heated for 2 days at 200°; a little copper was formed and a large 

quantity of a red salt which was not determined. The blue color of the solution 

practically disappeared.

Experiment 112 Oxidation of Cu2 S at 200°. (See Table 8.) Four days' 

heating at 200° gave a bare trace of copper on cooling.

Experiment 112. Oxidation of Cu2S at 200°. (See Table 8.) Four days' 

end. Weights as given in Table 8; the solution was dilute. Heated for 6 

days with one end in furnace supposed to yield 250°, the other end water-cooled. 

No copper. The acidity was unchanged. Evidently the reaction did not progress appreciably, probably because 250° was not attained.

Experiment 11%. Similar to experiment 62, but more CuS04 was taken. 

Weights in Table 8. Heated at 250° for two days. A measurable quantity of 

copper deposited on cooling. Results of analyses in Tables 9 and 12.

Experiment 115a. Tube broke.

Experiments llSb and 115c. Heated at 300° for four days and six days respectively. Weights and analyses shown in Tables 8-12. Experiments at temperatures above 250° were made by inclosing the tubas in an iron bomb with, 

water inside to equalize the pressure on the tubes.

Experiment 116. Similar to experiments 115b and 115c; heated 56 days at 

250°. (See Tables 8-12.)

Experiment 117. Effect of FeSO*. (See Tables 8-12.)

Experiment 118. Run in a silica tube. Initial weights: Cu2S 0.32 gram, 

CuSO4 .5H2O 1.500 grams, H2S04 0.10 gram, H2 0 4.96 centimeters. Heated

II days at 350°; a little copper appeared on cooling. Not analyzed.

Experiment 119. Ferric oxide, cuprous sulphide, etc. (See Tables 8-12.) 

Heated at 300° for 6 days. A little copper was deposited on cooling, and the 

solution was slightly blue.

Experiment 120a. Action of sulphuric acid on Jena glass tubes. Exactly 

5 cubic centimeters of 0.5 normal sulphuric acid was placed in a Jena glass 

tube, 22 centimeters long, 1 centimeter in internal diameter, and heated in a 

horizontal position at 200° C. for 20 days. Original HS SO4 0.1225 gram, remaining 0.0932 gram, neutralized 0.0293 gram. Some crystals formed on the 

inside of the tube which qualitative tests indicated to consist of SiOs, S'04 ", Ca, 

and possibly K. There was not enough material to analyze. E. S. Larsen 

examined it microscopically and thought that three separate things were present ; some basal hexagonal tablets, uniaxial (-f-) w=1.5S and c about 1.60, possibly an alunite; some skeletal prisms with n about 1.47; and some needles
EXPERIMENTS 59

with high index of refraction. Some of the crystals formed little radiating 

clusters, and some grew along the glass. The solution was found to contain 

Si02 0.0192 gram, A1203 0.0010 gram, CaO 0.0055 gram, MgO trace, KaO trace, 

Na2O not determined.

Experiment 1201). Similar to preceding, in a tube 16 centimeters long, 

heated 11 days at 300°. Original H2 SO4 0.1225 gram, remaining 0.1125 gram, 

neutralized 0.0100 gram, which is less than found in 120a. No crystals formed 

in the tube. The solution gave SiOa 0.0067 gram, CaO 0.0010 gram, MgO trace, 

K2 0 0.0036 gram, Na2O not determined.

Experiment 121. Ferric sulphate and cupric sulphide. Resembled experiment 107. Initial weights and results shown in Table 8. Heated 'for 18 days 

at 200°; solution slightly blue but no copper. Heated 11 days more at 300°; 

some copper appeared on cooling. The sulphur of the solids was more than 

enough to form Cu2 S, so that either sulphur or CuS may also have been 

present.

Experiment 122. Resembled experiment 118. A silica tube of volume 

about 15 cubic centimeters was used. (See Tables 8-12.) No acid was used; 

that formed in the reaction prevented any formation of basic salt, even at 300°.

Experiment 128. Similar to experiment 122 but with H2 SO4 present, also 

less CuS04. The copper deposited on cooling was only about half that formed 

in experiment 123. (See Tables 8-12.)

Experiment 124. Action of cupric sulphate on sulphur dioxide. Initial 

weights: CuS()<.5H20 2.00 grams, S02 0.222 gram, H2S04 0.49 gram, H2O 10.0 

cubic centimeters. Heated at 250° for 1 day; apparently nothing but copper 

was formed. Heated 4 days more; in addition to the copper a trace of fine 

dark crystals, apparently either CuS or Cu2S, was formed.

Experiment 125. Resembled experiment 114 but with more CuSO<. Some 

of the CuSO4 crystallized out on cooling. Details in Tables 8-12.

Experiment 126. Resembled experiment 69. Initial weights: Cu2 S 0.10 

gram, CuCl 0.20 gram, Cu 0.10 gram, HCI 0.20 gram, H20 5.0 cubic centimeters. 

Heated 2 days at 200°; the CuCl recrystallized on cooling. Heated 5 days at 

250°; a very little copper was recrystallized and deposited as a red tint on 

some of the CuCl on cooling, especially at a certain zone near the end of the 

tube. No sulphate was formed.

Experiment 127. Similar to experiment 122 but heated at 350°. Details 

in Tables 8-12. Silica tube.

Experiment 118. Resembled experiment 49. Initial weights: Cu 0.50 gram, 

H2 SO« 0.50 gram, H2O 2.85 cubic centimeters. Heated 4 days at 350°; considerable Cu2S was formed. Some copper was recrystallized, but the results 

resembled in most respects those obtained by starting with Cu2S and CuS04. 

The acid was not all used up, but the copper may not have been present in 

sufficient amount. Details in Tables 8-12.

Experiment 129. Neutralizing action of rock from the Leucite Hills, Wyo., 

supposed to be wyomingite, on sulphuric acid. Initial weights: Wyomingite 

1.0 gram, HCI 0.5 gram, H2 0 5.5 cubic centimeters. Heated 2 days at 250°; 

HCI then 0.150 gram; neutralized 0.350 gram. Potassium in solution as K20, 

0.120 gram=12 per cent of rock. The alkalies were apparently all dissolved. 

In fact, a little potassium may have dissolved from the tube, or the rock may 

have been orendite with approximately 12 per cent K2O.

Experiment ISO. Neutralizing action of Jena combustion tubing, powdered 

and selected between 60 and 200 mesh. Quantity taken, 1.0 gram; IkSO* 

0.100 gram, H2O 6.96 cubic centimeters. Heated 2 days at 250°. H2SO4 was 

then 0.0684 gram; neutralized 0.032 gram.
60 DEPOSITION OF NATIVE COPPER

Experiment 131. Oxidizing action of cupric salts. Initial weights: Cu2 S 

0.16 gram, CuCl2 0.96 gram, H2O 5.60 cubic centimeters. Heated 2 days at 

250°; the sulphur of the Cu2S was completely oxidized to sulphate. A considerable quantity of CuCl crystallized out on cooling. The solution was a 

pale willow-yellow.

Experiment 182a. Resembled experiment 49. Initial weights: Cu 0.50 gram, 

H2S04 0.1225 gram, H2O 4.96 cubic centimeters. Heated 4 days at 250°. The 

resulting solution was pale blue. A small quantity of copper was recrystallized 

and a trace of Cu2 S formed. The recrystallized copper, separated as well as 

possible, after correction for a trace of sulphide, was found to be 0.0050 gram. 

Copper in solution 0.0069 gram. The solution was boiled and a very small 

amount of S02 , 0.008 gram, found by collecting in KMnO*.

Experiment 1321). A duplicate of the tube analyzed in 132a was heated 4 

days more at 300°. Found: In the solid, copper 0.473 gram, sulphur 0.0065 

gram, equivalent to 0.0323 gram Cu2S; in the solution, Cu 0.0225 gram, SO/' 

0.1005 gram, acidity b^ titration equivalent to 0.078 gram H2SO4. The results of 

calculation of the possible reaction, given in Table 12, do not conform very 

closely to the requirements of equation (17), which is the reverse of equation 

(36) expressing the action of cupric sulphate on chalcocite.

Experiment 133. Similar to preceding experiment. Initial weights: Cu 0.127 

gram, H-SO4 0.392 gram, H2O 0.392 cubic centimeter. Volume of tube 14 cubic 

centimeters. Heated at 300° for 4 days; the copper was entirely recrystallized 

with the production of Cu2S and CuS'O*. In fact, the solution nearly all crystallized as CuS04.5H2O. The products were not analyzed.

Experiment ISJf. Similar to preceding experiment. Initial weights: Cu2S0.160 gram, H2SO* 0.196 gram, H2O 7.93 cubic centimeters in all. Heated at 

300° for 4 days; there was no visible change. On opening the tube no SO2 

could be detected by the smell, but on boiling the solution, absorbing the gasin water, and titrating with KMnO<, a possible trace of S02 was indicated.

Experiment 135a. Effect of the presence of CuCl on reaction (36). Initial 

weights: Cu2 S 0.16 gram, CuSO* 1.21 grams, CuCl 0.40 gram, introduced dry in 

separate inner tube, H2 0 5.0 gram. Heated at 300° for 4 days. The tube was 

considerably attacked. A bright-green salt was formed and possibly CifcO but 

no copper.

Experiment 1351). Initial weights: Cu2 S 0.32 gram, CuSO< 1.21 gram, CuCl 

0.1996 gram, H2 SO« 0.1225 gram, H2O 4.96 cubic centimeters. Heated at 300° 

for 2 days. Some of the green salt formed, also very good crystals of Cu2O, 

especially on the poorer glass of the inner tube, but only one .or two small 

flakes of copper.

Experiment 136. To see if the copper is formed by reaction (36) at the high 

temperature. Initial weights: Cu2S 0.32 gram, CuSO4.5H2O 3.00 grams, H2O 

5.0 cubic centimeters. Heated at 250° for 15 hours and inspected while hot. 

The color was a dark green with a large quantity of basic salt, ,but no copper 

was visible. When heated 2 days more the quantity of basic salt had decreased, 

but no copper appeared. When heated 3 days more the basic salt was mostly 

gone; When heated 2 days more the solution was green, but no copper 

appeared. The tube was now allowed to cool in a sloping position. Crystals 

of copper appeared in the solution, which became blue. A little CuSO<.5H2O 

crystallized, also a little green basic salt and a few octahedra of cuprite.

Experiment 137. Similar to experiment 49. Initial weights: Copper 1.0003 

grams, HzS'O* 0.3691 gram, H20 4.88 cubic centimeters. Heated at 200° for 2 

days; no visible color developed, but a trace of Cu2 S formed. When heated 

at 250° for 7 hours more a pale-yellowish color developed, with considerable 

Cu2S. When heated 17 hours more the solution was faintly blue, and a few
EXPERIMENTS 61

specks of recrystallized copper appeared in the end of the tube, which was 

probably the coldest part; this shows bow easily the copper migrates from 

warm to cold regions. When heated 1 day more the solution was a pale 

greenish blue. On cooling the solution deposited a lot of copper and became 

blue. Most of the Cu-S was well crystallized, showing lathlike and feathery 

forms.

Experiment 188. Similar to experiments 49 and 132 but heated at 300°. 

Initial weights: Copper 1.00 gram, H2 SO« 0.3691 gram, HSO 4.88 cubic centimeters. Heated at 300° for 2 days and cooled. Tube now contained Cu2 S, recrystallized copper, and CuSO*, estimated colorimetrically at 0.086 gram copper. 

Analysis: Copper recrystallized. corrected for Cu2S, 0.0794 gram. Total copper 

in solid 0.9054 gram; total sulphur 0.0150 gram. Solution: Copper 0.0896 

gram, SO*" 0.3180 gram. The calculations based on these determinations are 

given in Tables 1 and 11.

Experiment 189. Similar to experiment 138 but heated at 350° in silica

tube. Duplicates, one in silica, one in glass. The glass tube was much attacked 

but held. Weights: Copper 1.000 gram, EbSO* 0.3691 gram, H2O 4.880 cubic 

centimeters. Heated at 350° for 2 days. The contents of the silica tube were 

analyzed. Solid: Recrystallized copper, corrected for Cu2S, 0.0996 gram, total 

copper 0.8704 gram, total sulphur 0.0212 gram. Solution: Copper 0.1246 gram, 

SO*" 0.3005 gram. For calculations see Table 1.

Experiment 140. Pyrite and copper in- the presence of water. Uncompleted.

Experiment l^la. Effect of high alkalinity. Weights: Fe20s 0.768 gram, 

CibS 0.064 gram, NaOH 5 cubic centimeters of solution of specific gravity 1.19. 

Heated at 200° for 3 hours; no visible changes occurred.

Experiment l^ll). Similar to experiment 141o. Weights: Fe2 08 0.764 gram, 

Cu2S 0.064 gram, NaOH 5 cubic centimeters of solution 15.5 normal. Heated at 

300° for 1 day. The tube cracked and was very much attacked, so that the 

results were not satisfactory. However, there were no evidences of copper 

or ferrous sulphate. Only on extracting with dilute acid was ferrous iron 

found, and this is explicable by the evidence of experiment 38.

Experiment 142. Pyrrhotite and ferric oxide. Weights: FeS 0.088 gram, 

Fe2Oa 0.64 gram, H20 5.0 cubic centimeters. Heated at 300° for 1 day; no soluble SO/' or Fe" was formed. Hydrochloric acid of course gave some Fe", 

but no SO*" in the solid.

Experiment 143. Ferric oxide and sodium sulphide. Initial weights: FeCla 

equivalent to 0.897 gram Fe2O3 , NaOH 1.88 gram=2.17 cubic centimeters of 15.5 

normal NaOH, Na2S 0.78 gram=l cubic centimeter of 2 normal Na2S. Heated 

at 300° for 1 day. Some FeS' was formed at once. The tube was much attacked. 

The inside of the tube was more or less covered by a greenish coating. No 

soluble SO*" or Fe" found; in fact, the solution was faintly alkaline. The 

solution gave a white precipitate on acidifying with H2 SO* and consumed a 

small amount of KMnO* equivalent to 0.0034 gram FeO. No pyrite could be 

detected. The dry powder was slightly magnetic.

Experiment 144- Ferric oxide and sulphur. Compare experiment 24. 

Weights: Fe2O8 0.500 gram, sulphur 0.05 gram, H20 5.0 cubic centimeters. 

Heated at 300° for 1 day; the solution contained considerable SO*" (not acid, 

hence probably K2SO* from the glass) and a little Fe". The solid appeared to 

contain some sulphide, perhaps FeS or FeS2 in the form of a slimy dull-gray 

powder not rapidly soluble in acid but of indistinct character under the microscope. Hydrochloric acid dissolved some ferrous compound, at least 0.0034 

gram FeO. No residual sulphur could be detected.
62 DEPOSITION OF NATIVE COPPER

Experiment 145. Ferric oxide and cuprous chloride in acid solution. 

Weights: Fe2 (S04)3 and H2SO«=0.0765 gram FesOs and 0.1535 gram HjSO* 

in 1 cubic centimeter of solution, H2O 2.0 cubic centimeters, CuCl 0.3015 gram 

(in an inner capsule). The substances were not mixed until the air was very 

thoroughly pumped out and the tube sealed. After mixing the solution soon developed a blue color, estimated colorimetrically to be equivalent to 0.056 gram of 

cupric copper. Theory for complete reduction of the ferric salt 0.061 gram. 

Heated at 125° for 2 days; some Fe2 0a and possibly other hydrolytic products 

were deposited, and the solution became somewhat paler.

Experiment 146. A blank test on the CuCl, treated as in the preceding experiment but without any ferric salt, gave no blue color. After heating at 125° 

for 2 days and cooling a very little copper deposited on cooling. No blue color 

was visible.

Experiment llfl- Gradual neutralization of a ferrous-cupric solution likely 

to deposit copper. Weights: H2O 50 cubic centimeters, FeS'04.7H2 0 0.027S 

gram, CuSO4.5H2O 0.250 gram, NaC2H3O2.3H2 0 1.09 grams. The material 

was sealed in a vacuum, the solids being in a separate tube. Mixed. A fine 

yellow precipitate soon began to form and in 15 minutes had flocculated. The 

solution was then somewhat greenish, and the precipitate gradually became the 

color of ferric hydroxide, verging slightly toward yellow. Heated at 100° 

for 4 hours and then cooled; the precipitate was decidedly red, although no 

particles of copper could be identified by any metallic appearance. The solution was nearly colorless, but after filtering off the precipitate in air it gave 

only 0.0050 gram iron and 0.0050 gram copper. When the solid was immersed 

in silver sulphate solution metallic silver was precipitated, but as cuprous 

oxide gives the same result this does not prove that metallic copper existed in 

the precipitate, and there is no evidence that it was present. The fact that 

ferric hydroxide appeared to be formed and that the precipitate reduced silver 

sulphate solution seems to show that the precipitate was a mixture of ferric and 

cuprous oxides or hydroxides.

Experiment 148. Similar to experiment 147, but with twice as much ferrous 

sulphate and sodium acetate. The precipitate did not coagulate as quickly as 

that in experiment 147. The flask was heated very gradually up to 100° over

the course of a week. The results were similar to those obtained in experiment 

147 except that the color of the precipitate was much more yellow, hardly at 

all red, resembling freshly precipitated cuprous oxide. The solution gave 0.0062 

gram iron and 0.0018 gram copper.

Experiments 149-153. Solubility of CuCl in solutions of CaCl2 at 25° C. 

The determinations were made in an atmosphere of nitrogen in the apparatus 

shown in Figure 2. It will readily be seen that water or calcium chloride 

solution could be freed from air and introduced into the tube or the saturated 

solution removed from it by connecting as desired, turning the proper stopcocks, and allowing nitrogen from a cylinder of the compressed gas to pass as 

slowly or fast as desired. The precipitated CuCl was first well washed. The 

tube containing the mixture was then shaken from time to time and kept at 

constant temperature in the thermostat accurate to 0.1°. After settling, the 

clear solutions were decanted into a graduated cylinder and weighed. They 

were then analyzed by separating the copper as sulphide and determining it by 

electrolysis. The calcium was precipitated as oxalate and weighed as lime. 

The results are shown in Table 13..

.Experiment 155. This experiment was similar to experiment 14. Initial 

weights: FeS<X7H2O 0.5 gram, CuSO4.5H2O 0.5 gram, HaO 5.0 cubic centimeters. The tube was observed from time to time during 24 hours while it
EXPERIMENTS 63

was being heated at 200°. No copper appeared, but some red substance, thought 

possibly to be cuprite, was formed. Copper appeared on cooling, however, but 

redissolved on heating. The experiment shows that the copper obtained in 

Stokes's experiment was probably precipitated by cooling and not by reaction 

at a high temperature.

Experiment 157. This experiment was performed to neutralize the acidity 

gradually and thus to increase the quantity of Cu2 SO« formed. Initial weights: 

FeS04.7H2 0 0.5 gram, CuS04.5H20 0.5 gram, wyomingite 0.2 gram, H2O 

5.0 cubic centimeters. After heating the tube at 200° for two days no copper 

was visible in the hot solution. The tube was then tilted and cooled slowly. 

When cold a considerable quantity of copper had separated from the solution, 

and no significant quantity of it could be seen in the solids, showing that the 

copper is not formed while hot, at least for a certain initial period.

Experiment 158. Initial weights: FeSCh.TI^O 0.5 gram, CuSO4.5H»O 

0.5 gram, Cu2 0 O.S gram, M2 0 5.0 cubic centimeters. The tube after heating

FIGURE 2. Apparatus used in experiments 149^153

at 200° for two days showed no copper while hot, but after tilting the tube 

and cooling as in the previous experiment copper was observed in the solution 

and only a very few specks in the solid. Some feathery crystals were also 

present in the solution, but not definitely shown to be CiuSCh. There was not 

as much hematite in the solid as in the preceding experiment.

Experiments 159a~-159g. Reaction of CuCl on FeO at 100°. A ferrous 

chloride solution was made by immersing an excess of iron in hydrochloric 

acid, the solution being protected by an atmosphere of nitrogen. This solution 

was withdrawn quantitatively as desired and precipitated by an exact equivalent of caustic soda solution, also in an atmosphere of nitrogen. In most of 

the experiments saturated sodium chloride was used as a diluent to the extent 

of half the volume of the main solution. Similarly, cuprous chloride dissolved 

in a concentrated solution of sodium chlpride was added to the suspension of 

ferrous oxide. All reactions were done at about 100° C. in a closed flask, the 

contents of which could be boiled in a slow stream of nitrogen. In experiment 

159c the sodium hydroxide was added to the mixed ferrous and cuprous salts

42980 25 5
64 DEPOSITION OE NATIVE COPPER

instead of first precipitating the ferrous oxide. The data obtained are given 

in Table 14, in which the lines entered as " theory " are computed from the 

weights of FeO or CuCl taken, depending on which was in excess, according 

to equations (71), (72), and (73).

3FeO+2CuCl=2Cu+Fe203+FeCl2 (71)

FeO+2CuCl=Cu2O+FeCl2 (72)

4FeO+2CuCl=2Cu+Fe3O4+FeCl2 (73)

The precipitate formed in these experiments, after drying on filter paper in 

air, was strongly magnetic. This suggests that magnetite was one of the products, as indicated by equation (73).

The data in Table 14 show that equation (72) may be left out of consideration. Equation (71) evidently expresses the principal change, especially when 

CuCl is in large excess. At the same time, however, some magnetite appears 

to be formed, as expressed by equation (73), and this equation appears to be 

still better verified as the excess of CuCl is decreased, as would be expected.

In experiment 159f, a single blank test with cuprous chloride solution and 

metallic copper, a little metallic copper appeared to be dissolved in the usual 

procedure at 100°. Also, while equations (71) and (73) serve to represent 

the essential changes, there is a little too much iron dissolved, the cause of 

which is not clear, although it may be caused in part by hydrolysis of the 

cuprous chloride.

The results of the experiments are in harmony with the following facts: 

(1) Cuprous oxide is of the same order of solubility as ferrous oxide. (2) 

Ferric oxide is much less soluble. (3; The equilibrium

Cu+Fe+++<=*Cu+ + Fe++ (74)

lies naturally toward the right but is displaced toward the left in alkaline 

or slightly acid solutions by the relatively much smaller solubility of ferric 

oxide.

According to the writer's computations from available electromotive-force 

measurements, the equilibrium constant for equation (74) at ordinary temperature is

[CylJ£f+] = 5,000 (75)

If the solubility products of Fe(OH) 2 , Fe(OH) 8, and CuOH at 25° are 

taken as 19.8X10"15, 1.1X10"88 and 1.6X10'9 respectively, it is easily calculated 

that in solution as alkaline as Fe(OH) 2 itself ferrous oxide should readily 

convert cuprous salts into metallic copper. As a matter of fact, however, ferrous chloride is formed, and this is even more favorable for the precipitation of 

copper. The final solution in experiment 1590, for example, contained ferrous 

chloride, 0.0042 molal, and cuprous chloride dissolved in sodium chloride, 

0.0281 molal. The unknown ionic concentrations are [Cu+] and [Fe+++ ], the 

first depending on the existence of a complex ion and the second depending on 

fOH~] that is, on the hydrogen ion concentration or pH value, and possibly 

on the existence of magnetite as a solid phase.

According to Bodlander and Storbec^k 78 [Cu+ ] will be highest in the least saline 

solutions. Evidently, however, [Fe++ ] can be varied more widely than [Cu*].

78 Bodlander, G., and Storbeck, 0., BeltrSge zur Eenntnis der Cuproverblndungen: 

Zeitschr. anorg. Chemie, vol. 31, pp. 1-41, 458-476, 1902.
EXPERIMENTS 65

Experiments 160a^l60j. Action of FeCU on CuCl in solutions of regulated 

pH, or hydrogen ion concentration. The pH value was fixed by different proportions of acetic acid and sodium acetate. When ferrous chloride and cuprous 

chloride were introducted into such solutions, boiling in an atmosphere of 

nitrogen, no copper was precipitated if the pH value was initially less than 

obout 5.0. Under certain conditions, shown more fully in Table 15, with a pH 

value of 5;2 a slight film, probably containing some copper, formed on the walls 

of the flask in the course of 2 hours boiling. The experiment was repeated 

with less sodium chloride present, but no crystals of copper could be detected. 

With a pH value of 5.38 and somewhat more sodium chloride, the deposit 

showed some small particles of ferric oxide, each surrounding a nucleus of 

minute crystals of copper visible with a lens.

Repetition of the experiment with more ferrous chloride and less cuprous 

chloride, in a flask that had a slight film remaining from the previous experiment, gave a deposit that showed under a lens or a microscope numerous crystals of copper, some spicules, some thick crystals showing triangular faces,

most of them largely covered with iron oxide, which also adhered to the glass 

in the form of minute lumps and floes without definite form. This iron oxide 

was not magnetic. The reaction that occurred here may be written

2FeCU+2CuCl+6NaCOOCH,+3HaO=2Cu+Fe»08+6HOOOCH8+6NaCl (76)

These experiments show that the use of ferrous hydroxide is not absolutely 

necessary for the deposition of metallic copper and also that copper is a recognizable product when the reaction is made to occur very slowly, and further 

that the deposition of copper is not dependent on cooling in the chloride solutions, as it seems to be in the sulphate solutions (experiments 155-158).

Experiments 161 and 162. Unfinished.

Experiment 163. This experiment was a repetition of experiment 3. The 

tube was heated at 150° for 25 hours. There were no signs of recrystallized 

copper on cooling.

Experiment 164- Similar to experiment 31. Weights: CuCl 4.5 grams, 

NaCl 3.5 grams, Cu 2.0 grams, H2O 6.0 cubic centimeters. This tube was 

sealed with some air in it, then heated to about 80°. Soon after, one of the 

wires was seen to be much attacked near one end so that it had the appearance of a bunch of crystals at that point and some crystals or fragments of 

copper were loose in the tube. It would therefore have been impossible to 

note any recrystallization of copper in this tube. When the tube was heated 

at 150° for 15 hours a little green basic salt formed, and much of the CuCl 

recrystallized on cooling. When the tube was heated at 175° for 5 hours 

and tilted while hot the green basic salt remained with the copper and had 

therefore not dissolved in the hot solution, but most of the CuCl had dissolved and recrystallized on cooling.

Experiment 165. Similar to experiment 164, but 4.0 cubic centimeters of 

H2O and 2 cubic centimeters of concentrated HC1 were taken instead of 6 

cubic centimeters of H2O. When the tube was heated at 150° for 15 hours no 

recrystallized copper was observed on cooling. The solution was practically 

colorless. The tube was heated at 175° for 5 hours and tilted hot; about 95 

per cent of the CuCl recrystallized, showing a solubility of approximately 700 

grams of CuCl per liter in this solution at 175° C. No basic salt was formed. 

The CuCl was beautifully crystallized.
66 DEPOSITION OF NATIVE COPPEE

TABLE 8. Substances taken and conditions of experiment 

[All weights in grams]

Experiment

No.

0 IK Itf 62

»64

"65

 66

 67

» 72

'77

'78

'101

"102

 108

'109

"110

olio iifi »113

< 114

1156

115C

116

117

119

120o

1206

121

122

123

125

127

128

1320

1326

138

139

Temperstare

/Ofl \

\ ^->->

200

230

200

200

200

200

200

200

200

200

200

200

200

200

200

200

250

300

300

250

250

300

200

300

300

300

300

300

350

350

250

300

300

350

Duration of

experiment

(days)

3

2

1

6

Qy

6

3

3

6

6

1

4

6

2

4

10

56

8

6

16

11

11

11

11

4

4

4

4

4+

2

2

CuS04

0.320

.320

.320

.160

.320

.243

».26

 £ri 94

.24

.24

1.919

1.919

1.919

1.919

.640

None.

None.

None.

None.

1.215

.92

3.84

1.21 *.50

*.50

*.50

*1.00

*1.00

Cu2S

0.02

.08

.08

..08

.08

.08

.02

. 09 vfi

.02

.24

.30

.08

.08

.08

.08

.08

.1600

.1600

.1600

.1600

.1278

.1000

None.

None.

.1000

.3200

.32

.500

.16

None.

None.

None.

None.

None.

HjS04

0.06

.23

.69

None.

None.

.10

d.024

d.040

'.201

M95

".28

.23

.23

.340

.340

.340

.340

.340

.444

.1225

.1225

.742

None.

.70

.340

None.

.500

.1225

.1225

.3691

.3691

FeS04

0.50

"4.0

.273

.50

None.

None.

None.

None.

.546

None.

None.

None.

None.

None.

None.

None.

None.

None.

None.

None.

None.

None.

Fea03

0.08

.04

.130

'.325

'. 150

None.

None.

None.

None.

None.

.192

None.

None.

M92

None.

None.

None.

None.

None.

None.

None.

None.

None.

HnO

5.0

5.35

5.69

5.2

5.09

5.26

4.8

5.0

5.23

2.0

3.5

5.14

5.08

5. 13

20.3

30.3

8.25

8.25

8.25

8.25

6.06

5.00

5.35

5.35

5.21

5.69

5.26

5.00

5.69

2.85

4'. 96

4.96

4.88

4.88

 A trace of copper formed. 

6 No copper observed.

 CaCl2 .

 *S02.

«H01.

/ FeCl3 . 

' CU20. 

» CuCh. 2HjO.

 ' Measurable copper formed in most of the succeeding experiments. 

' Fej(SO*) 3 .

 Metallic copper.
EXPERIMENTS 67

TABLE 9. Analysis of solid residues

[All weights in grams]

Experiment 

No.

62

101

102

104 

108

110

112

114 

1156 

115c 

116 

117 

119 

121

122 

123 

125 

127 

128 

132a 

1326 

138

Temperature 

(°0.)

220

200

200

250 

200

200

200

250 

200 

200 

250 

250 

300 

300

300 

300 

300 350' 

350 

250 

300 

300

Duration of experiment 

(days)

6

6

6

5 

2

2

4

2 

4 

10 

56 

8 

6 

11

11 

11 

4 

. 4 

4 

4 

4 

2

Total 

copper

0. 0659

.0946

.1461

.1422 

.1977 

.2120 

.1695 

.1372 

.0530 

.0500

.3405 

.276 

.5033 

. 2590 ' 

.2920 

0.0067 

.473 

.9054

Sulphur

0. 0343

.0592

.022 

.0151

.0160

.0211 

.0225 

.0157 

.0260 

.0255 

.0101 

.0143

.054 

.0515 

.626 

.0087 

.0347 

.0004 

.0065 

.0147

Copper equivalent to 

sulphur

0. 1360

.2347

.0872 

.0600

.0838 

.0894 

.0623 

.1031 

.1014 

.0402

' .214 

.208 

.248 

.035 

.1376 

.0017 

.0258 

.0582

Metallic 

copper

A little.

0.0589

.0584 

.1083 

.1497 

.0664 

.0358 

.0128

Present.

.126 

.068 

.255 

.224 

".1544 

«.0050 

.0447 

.847

« 0.3456 reacted.

* Copper recrystallized with a little sulphide.

-«.Copper recrystallized.

TABLE 10. Analysis of solutions

[All weights in grams]

Experiment 

No.

101 

102 

104

109

113

114 

1156 

115C 

116 

117 

120o

1206

191

122 

123 

125 

»127 

»128 

132a

1326 

138

Copper

"o.'isos"

.6622 

.6867 

.6475 

.7330 

.2387

.0165

.4008 

.3425 

1.391 

.329 

.1810 

.0069

.0225 

.0896

Total S0<"

0.0378 

.0033 

.4092

1.436 

1.498 

1.526 

1.512

SSflQ

.7704 

1.2878 

2.604 

.775 

.4013 

.1138

.1005 

.3180

HsSO* 

by titration

"6.~2l64~~

-.0080

.23

.3629 

.4682 

.5134 

.3617 .2862 ' 

.0932

.1125

.729

.144 

.777 

.509 

.263 

.1051

.078

* 

HzSO« 

by calculation

"6."0034"

.450 

.478 

.560 

.417

.174 

.786 

.50 

.29 

.129

.070

' SOi" only. »Silica tube.
68 DEPOSITION OF NATIVE COPPER

TABLE 11. Calculation of weights that reacted, based <m preceding data'

[All weights in grams]

Experiment 

No.

102

108

114 

1156 

115c 

116 

117

122 

123 

125 

127 

128 

132a 

1326 

138

CuSO« 

in final 

solution0

1.665 

1.725 

1.626 

1.841 

.5998

1.005 

.860 

3.50 

.773 

».455 

.0169 

' .0565 

.2251

CuSO4 

that 

reacted

0.255 

.195 

.293 

.078 

.0402

.210 

.099 

.30 

.442 

».455 

» . 0158 

».0565

CuzS in 

residue

0. 2941

.0750

.1048 

.1118 

.0780 

.1291 

.1267

.268 

.260 

.311 

.044 

> . 172 

» . 0021 

'.0323

Cu2S 

that 

reacted

0.0059

.0050

.0552 

.0482 

.0820 

.0309 

.0333

.052 

.060 

.189 

.116 

>.172 

'.0021 

'.0323

S<V' 

combined 

with 

copper in 

solution

0.996 

1.030 

.975 

1.103

.604 

.515 

2.10 

.495 

.274 

.0100 

.0340 

.1335

Balance 

of S04"

0.0033

.441 

.468 

.550 

.409

.167 

.773 

.51 

.28 

.127 

.104 

.068 

.1825

Equivalent weightofHjSOi.

a 0034

.450 

.478 

.560 

.417

.170 

.789 

.52 

.29 

.129 

.106 

.070 

.1862

HsS04 

formed

0.0034

.110 

.138 

.220 

.077

.170 

.089 

.180 

.29 

".38 

'.017 

«.053 

«.183

Including CuiSO*. »Formed. That reacted.

TABLE 12. Molecular ratios of reacting substances

Milligram-molecules

Experiment No.

A 

1U. .................................................. .

1151). ..................................................

115c... ... 

116. ............................................... .... in. ................................... ................

122.. ................. . . 

123 _ .. __ . . .. 

125....... .. . . . , 

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

128

1326....................................................

Temperature 

(°0.)

250

300

300

250

250

300

300

300

350

350

300

CuSO<

1.60

1.22

1.84

.489

.252

1.31

.62

1.88

2.77

2.85

.35

CuzS

0.347

.30

.515

.194

.209

.33

.38

1.19

.729

1.08

.20

Cu

0.920

1.70

2.36

1.05

.564

1.98

1.06

4.01

3.52

5.44

.76

HjSOj

1.12:

1.40

2.25

.79"

1.73-

.91

1.84

2.95

3.88-

.54-

Molecular ratios

114....................................................

1156........ ....... . . 

115c . ._ ..........................................

116 ___ ... _ . _ ........ __ .... __ ........ . ...

117......................... ...........................

122.. ........................................... ... 

123......................... ................. ..........

125....................................................

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

128. . .

1326..-. _ . . ..

250

300

300

250

250

300

300

300

350

350

300

4.6

4.1

3.57

2.5

1

4.0

1.6

1.6

3.8

2.6

1.7

3

1

1

1

1

1

1

1

1

1

1

1

1

2.7

5.7

4.58

6.4

3

6.0

2.8

3.4

4.83

5.02

3.5

5

3.2

4.6

4.38- 4. or

5.3

2.4

1.5

4.06-

3.6

2.7

4
EXPERIMENTS 69

'TABLE 13. Solubility of cuprous chloride in calcium chloride solutions at 25° C. 

[Concentrations in per cent by weight]

Experiment

No.

153

152

149

151

150

Specific 

gravity of

solution

1.04

1.12

1.17

1.30

1,48

CaCls

3.33

9.75

13.48

19.84

25.55

CuCl

0.438

1.934

3.698

8.503

19.71

TABLE 14. Reaction of CuCl on FeO in aqueous solution at 100" 0. 

[All weights in grams]

Experiment 159<J __ ......................... ..... 

Experiment 1590..... ______ . _ . _____ -.

'Theory for equation (72) _ - ___ - ___ ....

'Theory for equation (73) __ ________ - ___ --

Experiment 159c .....................................

Theory for equation (73) ____ ... ___ -. __ -.

Experiment 159e __________________ _.

Experiment 159/. __________________ -

Volume 

(cubic centimeters)

139 

196 

200

144

200

187

Fe

taken 

as FeO

0.141 

.141 

.141 

.141 

.141 

.141 

.141 

.141 

.049 

.141 

.282 

.211 

.070 

.281 

 NoneCu

taken 

as CuCl

0.324 

.444 

.444 

.107 

.321 

.080 

.111 

.107 

.111 

.080 

.160

.160 

.444

Cu 

precipitated

0.092 

.090 

.087 

.107 

.321 

.080 

.081 

.107 

.111 

.080 

.155

.160 

k .009

Pe dissolved

0.074 

.047 

.047 

.047 

.141 

.035 

.040 

.047 

.049 

.035 

.080

.070

1 Metallic copper taken in excess in place of FeO. *Cu dissolved.

TABLE 15. Reaction of CuCl and FeCli at 100° m solutions of regulated pH

value

[Final volume of solution, 200 cubic centimeters]

Experiment 

No.

160o...... 

3606.......

aeoc......

160d.__... 

160e.......

160/_._.__.

160?......

160A......

16W.......

16Q/.......

NaCl 

(grams)

2.20 

1.97

1.50

2.08 

1.08

.43

.43

.43

1.30

2.17

Normal 

CHsCOOH 

(cubic centimeters)

20.0 

16.0

8.0

,5.0 

3.2

5.0

8.0

10.0

12.0

5.0

Normal 

CHsCOONa 

(cubic centimeters)

20.0 

24.0

32.0

35.0 

36.8

35.0

32.0

30.0

28.0

35.0

PH

at 25° C.

4.5 

4.7

5.20

5.38 

5.62

5.45

5.23

5.08

4.93

5.38

0.504 M 

FeClj 

(cubic centimeters)

1.7 

1.7

1.7

2.0 

2.0

5.0

5.0

5.0

5.0

5.0

0.926 M 

CuCl 

(cubic centimeters)

4.8 

4.8

4.8

5.0 

5.0

2.0

2.0

2.0

2.0

2.0

Results

No precipitate. 

Do.

Very slight precipitate

after some time. 

A faint cloud appeared 

in 2 minutes, a considerable film on 

glass in 2 hours with 

small crystals of Cu.

once and was too 

rapid for the formation of many recognizable crystals of 

Cu. 

Precipitation fairly

rapid.

crystals of Cu seen. 

Do.

on walls of flask in 

2 hours.
\
INDEX

Page 

Acids, action of, on mixtures of ferric oxide

and copper sulphide... . 29-33 

attack of, on glass and minerals. 10-11,58-59 

Acknowledgments for aid_ . , 3,51-52 

"Bleaching of rocks near copper._ 30-33 

Brown, Qlenn V., acknowledgment to. 51-52 

Butler, B. S., theory advanced by. 1,26

 Calcium chloride, action of, on copper 16-17

 Caustic soda, action of, on copper.. . 17

Chalcocite, action of cupric sulphate on 35-38

action of ferric salts on... 32

behavior of cuprous chloride with 39

copper obtained from, by action of steam. 6

'Chalcocite and hematite, experiments on... 45,47

 Colloidal solutions, experiments with..... 54-55 

^Combinations of copper at depth '....... 3-12

 Covellite, action of cupric sulphate on 33-35 

Cupric chloride, action of, on cuprous sulphide at 500°.................... 53

 Cupric sulphate, action of, on cuprous sulphide..- - 51

action of, on sulphur dioxide.. 59 

decomposition of, by heat... 5 

solubility of... * 18

 Cuprous chloride, melting point of, lowered

by other chlorides.. .... 3-4 

solubility of, in calcium chloride solutions- . 19,62,69

 Cuprous chloride and ferrous chloride, reaction of_ 69 

'Cuprous sulphide, action of air on, at 500°.. 55 

dissolved by concentrated sulphuric acid. 11 

reduction of, by hydrogen. 4 

by steam... .. 6

 Cuprous sulphide and ferric chloride action

of.. . - 44,55,57

.Deposition by cooling... . . . 39,60 

by dilution... -- - 22,39-40,50 

by fractional oxidation of copper sul- , 

phide.. 40-41,51

by neutralization and reduction... 41-42,62-65 

comparison of theories of__ 42 

See also Precipitation.

"Diffusion of copper in aqueous solutions. 20-21

.Euteetics, melting points of.__ 4

^Experiments, analyses of solid residues from. 67

analyses of solutions from.... 67

descriptions of______________ 42-65

molecular ratios of substances reacting in. 68

substances taken and conditions for . 66

weights of substances that reacted in 68

"'Ferric hydroxide, solubility of, shown by ordinary tests... 43

"Ferric oxide, action of sulphide waters on... 28-29 

solubility of, in ferric sulphate solutions.. 29-30

Ferric oxide and cuprous sulphide, experiment with, at 300°-...-... .. 49

Terric salts, action of cuprous sulphide on... 32 

ferric sulphate and copper, action of. 17,44-45

Terrous chloride, action of cuprous chloride

on..__......................... 65-69

^Ferrous oxide, action of cuprous chloride

on 63-64,69

42980 25 6

Page 

Glass, attack of acids on_______ 10-11,58-59

Hematite and Chalcocite, experiments on._ 45,47 

Hot solutions, reactions in__________ 60 

Hydrochloric acid, action of, on copper. __ 16 

obtained from Calumet conglomerate

rock . . 17-18 

Hydrogen, reducing action of..._______ 7-9

Iron oxides, equilibrium of, with carbon

oxides _______ 12 

equilibrium of, with water_______ 11-12

Jena glass, action of sulphuric acid on____ 58-59

Magma, pressure in, increased by crystallization.. . . 

Magmatic gases, sulphur compounds in..... 

Mine water, experiments with...............

Nickel-iron segregations, copper in.

50

4

Oxidation, connection of reduction with__ 12 

Oxidation of copper solutions by ferric compounds... __ 26-33

Precipitation of copper by carbonaceous

matter............._______ 24

by dilution of cuprous chloride solutions. 22 

by electro-chemical action. _______ 22 

by ferrous compounds ..._____ 22-24 

by sulphur dioxide....._________ 24-25

I from cuprous sulphate solutions..___ 21

See also Deposition. 

Pyrrhotite and ferric oxide, behavior of___ 61

Reactions, reversible____________ 6-12 

Reduction, connection of oxidation with__ 12

Scope of the report. _____________ 2 

Silver sulphide, action of sulphuric acid on.. 52 

Smelting, reactions that take place in....__ 3 

Solubility of copper, experiments on ____ 51 

Solubility of copper and some copper compounds______________ 12-19 

Source of Lake Superior copper..._____ 2 

Steinmann's experiment________ 28,46,47,61 

Stokes's experiment, repetition of______ 45 

Studies on the origin of native copper____ 1-2 

Sulphur, oxidation of, to sulphuric acid___ 32-33 

and cupric sulphate, action of___ __ 46 

and ferric oxide, action of ________ 46,61 

and ferric sulphate, action of______ 45-46 

and sulphuric acid, action of.. ____ 46 

Sulphur dioxide, action of, on copper.___ 5,65 

action of, on cupric sulphate______ 51 

on ferric oxide. ___..'. 43,53 

and hydrogen, action of, on copper at 900°. 53 

Sulphuric acid, action of, on copper___ 12-16,57 

behavior of, with cuprous chloride_.... 54

Thiosulphates, formation and occurrence of.. 8

Transportation of copper..__....___... 19-21

Traps, water content of___ 5

Water, copper in the absence of. .... 3-5

copper in the presence of._..___ 5-12

Wyomingite, action of, on sulphuric acid.... 69

71
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