*S. F. Emmons, 17th Ann. Rep. U.S.G.S., Part 2, p. 462. † Reed, Bull. Dept. Geol. U. of California, 4, pp. 189 and 192. Lindgren, 17th Ann. Rep. U.S.G.S., Part 2, p. 121. Beck, Nature of Ore Deposits, Weed's translation, II, 377. W. H. Emmons, unpublished manuscripts. L. J. W. Jones, Proc. Colo. Sci. Soc., VÏ (1897), 48. **R. C. Wells. been made of the waters of gold and silver mines. For them the writer is indebted to Professor W. H. Emmons who has very kindly placed at his disposal the data which he is about to publish concerning secondary enrichment. The figures given are in parts per million. 1. Geyser silver mine, Custer Co., Colo., 500-foot level. 2. Same, 2,000-foot level. 3. Comstock lode, Savage mine, 600-foot level. 4. Comstock lode, G. and C. shaft, 2,250-foot level. 5. Comstock lode, Comstock mine, vadose water. 6. Comstock lode, Gould and Curry mine, 1,700-foot level. 7. Comstock lode, Hale and Norcross mine. 8. Comstock lode, Gould and Curry mine, 1,800-foot level. 16. Creede, Colo., Solomon mine, 1,500-foot level. 17. Idaho Springs, Colo., Stanley mine. 18. Tonopah, Nev., Mizpah mine; water from a bore hole 2,316 feet deep. A study of these data shows clearly that the salts most abundant in mine waters are the sulphates, and further, that there is usually present an excess of acid radicles over basic; i.e., the solutions are acid in most cases. The preceding analyses are all of waters taken at some depth, hence must be assumed to be much more nearly neutral than the waters near the surface, since the tendency of descending waters is to become less acid through reaction with minerals. At Creede, as will be noted, the waters at the 1,500foot level have actually become alkaline through the solution of carbonate in excess. It is a significant fact that this level is below the zone of secondary enrichment.1 The amount of ferric and ferrous salts present in the waters, as shown by the analyses, is surprisingly small. This is probably to be explained by the fact that iron salts are very easily precipitated by carbonates, which are here present in large amount. On the whole, the analyses tend to confirm the hypothesis that the active agents in the secondary processes are acid sulphate waters. They would also suggest that carbonate solutions may be a factor 1 W. H. Emmons, unpublished manuscript. in the processes, but this supposition should be confirmed by analyses of waters from horizons somewhat nearer the surface than those from which the above samples were taken. EXPERIMENTAL WORK The experimental work described in this paper deals with the following questions: a) Solvent effect of sulphuric acid and ferric sulphate on argentite and its associated sulphides; both natural and chemically pure artificial minerals being used. b) The solvent effect exerted on metallic silver by the various reagents that may occur in ground-waters, such as sulphates, chlorides, nascent chlorine, sulphuric and hydrochloric acids. c) Solvent effect of ferric sulphate solutions on silver chloride. d) Effect of the presence of ferric sulphate on the solubility of silver sulphate. e) The equilibrium in dilute solutions between ferric, ferrous, and silver sulphates and native silver. f) The substitution of silver for antimony or arsenic in the previously formed sulphides of these elements. g) The reaction of metallic silver with precipitated sulphur. As these experiments deal with somewhat widely different subjects, each series will be described and discussed separately. All of them were made under conditions approximating those which obtain in ground-waters. The temperatures were uniformly room temperatures, about 22° C. Pressures were atmospheric pressures. Concentrations of solutions were small, usually decinormal; these, although somewhat greater than those that obtain in ground-waters, may be confidently assumed to cause differences only in the speed, and not in the nature, of the reactions which take place. The paper also includes a discussion, based on the experimental work of Barlow' and Schierholz,2 of the effect of chloride solutions on the solubility of silver chloride, and the precipitation of the silver from such solutions as sulphide. 1 Barlow, Jour. Am. Chem. Soc., XXVIII (1906), 1446. 2 Schierholz, Sitzungsberichte der Kaiserlichen Akademie der Wissenschaften zu Wien, 101, 2b (1890), 8. METHOD OF PROCEDURE The mineral to be tested was first powdered. In the preliminary experiments natural minerals were used, and only that portion which would pass through a 40-mesh sieve was used. It was thought, however, that the unequal sizing of the mass of mineral so obtained might affect the results materially. Therefore in the later experiments with chemically pure materials, only the fraction that passed an 80-mesh and was held by a 100-mesh sieve was used. Of this material a certain amount, usually 1.0000 gm., was weighed out, washed into a flask, and covered with 200 c.c. of the solution whose action was to be tested. The corked flask was then set away in the dark room at room temperature. After standing for a period of from one to three months, during which time the flask was shaken almost daily, the contents were analyzed. In some cases the analysis was of the liquid contents, after the solid residue had been removed by filtration; in others, where such analysis would prove difficult, it was thought sufficient to determine the materials in solution qualitatively, and the loss of weight of the solid residue. The latter was done by filtering into a weighed Gooch crucible, drying at 120° C., and weighing. This procedure was always adopted when the minerals used were the double sulphides pyrargyrite and polybasite. Where stibnite was the sulphide acted upon, the procedure differed, in that the solutions themselves were analyzed. This was done because of the ease of the analysis and for the greater accuracy thereby obtainable. In the absence of ferric sulphate the solutions were repeatedly evaporated to dryness in the presence of nitric acid. The resulting precipitate was then heated to a dull-red heat, whereby it was converted into the oxide Sb,O,, then cooled, and weighed. Where ferric sulphate was present, hydrogen sulphide was passed in until there was no further precipitate. The solution was then filtered, the filtrate being repeatedly passed through the filter until clear, after which the precipitate was dried and washed with carbon disulphide to remove excess of sulphur. The antimony trisulphide remaining on the filter was redissolved in concentrated hydrochloric acid, the resulting solution mixed with concentrated nitric acid, and cautiously evaporated on a water bath. As fast as the solution became colorless, more nitric acid was added, until further addition caused no change of color. The solution was then evaporated to dryness, the precipitate heated to a red heat, cooled, and weighed. Trial of this method with a weighed amount of pure antimony trisulphide showed it accurate to I per cent. I MAKING UP SOLUTIONS A ferric sulphate solution was made, containing 35 grams of Kalhbaum's C.P. powdered ferric sulphate per liter. This gave a solution nearly F/20 (actually 0.0535 F. An F/1 solution = solution containing 1 formula weight in grams per liter). The sulphuric acid solution used was roughly F/20. That used in the first series of experiments was 0.0635 F; that used in the second series, 0.0502 F. These solutions will hereafter be mentioned simply as "ferric sulphate solution" and "decinormal sulphuric acid," and in the tables will be designated as "F" and "A," for brevity. A proportionality, such as A:F::1:3, indicates the proportions of acid and ferric solutions. The total volume of all solutions, unless otherwise stated, is 200 c.c. ACTION OF SOLUTIONS ON MINERALS. SERIES I As may be seen from the tables, this series of tests included not only the silver minerals, but also those with which it is most closely associated, i.e., the sulphides of arsenic, antimony, and lead. Copper sulphide was not tested, since Vogt' had already proved its ready solubility in ferric solutions. The table also shows that both the dilute sulphuric acid and the dilute ferric sulphate solutions exert a powerful solvent action on all the minerals tested except argentite; that in each case the action is much more powerful when ferric sulphate is present; and that, except in the case of galena, an increase in the concentration of the ferric sulphate does not cause a corresponding increase in the solvent action. The absence of results in the case of the lump argentite is rather to be ascribed to the smallness of the surface exposed than to an actual absence of action. This will be shown in the table (Series 2). Vogt, Genesis of Ore Deposits (1896), p. 676, footnote. |