Geochemical Journal, Vol. 21, pp. 291 to 305, 1987

Ag2S solubility in sulfide solutions up to 250°C

ASAHIKO SUGAKI', STEVEN D. SCOTT', KENICHIRO HAYASHI', and ARASHI KITAKAZE1

Institute of Mineralogy, Petrology and Economic Geology, Faculty of Science, Tohoku University, Sendai 980, Japan' and Department of Geology, University of Toronto, Toronto M5S 1A1, Canada'

(Received December 2, 1987: Accepted February 22, 1988)

In order to estimatepossible silver sulfide complexes in ore solutions, the solubilityof Ag2Swas measuredbetween 25° and 250°Cin NaOH-H2S-H20solutions of 0.0 to 4.1m NaHSconcentration. Solu bilitychanges as a functionof temperature,total reducedsulfur concentration ES (mHS + mHS-)and pH. A maximumsolubility of 2140ppmAg was obtained at 250°Cin 4.1m NaHSconcentration under PH 2Sof 29.1 atm. From these solubility data, reactions to form silver sulfide complexes are estimated as follows: Ag2S + H2S = Ag2S(H2S), Ag2S + H2S + HS = Ag2S(H2S) (HS)-, Ag2S + H2S + 2HS = Ag2S (H2S) (HS)22 and Ag2S + 2HS = Ag2S (HS)22-. The equilibrium constants for these reactions are given in Table 2. Assuming that such silver sulfide complexes as above are in an ore solution , Ag2S (argentite or acanthite) is precipitated in response to changes of temperature, pH and total sulfur concentration (ES). Decrease of pH and ES is more effective than that of temperature. Silver sulfide complexes are more important than silver chloride complexes in ore solution of high ES, neutral to slightly alkaline pH and geologically reasonable chloride concentrations.

INTRODUCTION to the experimental study by Seward (1973), Thermochemical data for aqueous metal bisulfide complexes are significant for the trans species at elevated temperatures and pressures port of gold in ore solutions under epithermal are indispensable for understanding transport of conditions, even if chloride complexes have the metals in hydrothermal solution. With these ability to transport fairly large amount of gold data, the physico-chemical conditions that are at higher temperatures (Henley, 1973). In light necessary if ore solutions are to transport of the above, there arises the possibility that enough metal to form ore deposits can be evalu silver is also transported as sulfide complexes ated. Furthermore, processes of ore deposition under epithermal conditions. The stability of such as cooling, pressure change, reaction with chloride complexes of silver up to 350°C have wallrocks, mixing of solutions (including dilu been determined by Seward (1976), but the tion) and boiling (Skinner and Barton, 1973) thermodynamic data for silver sulfide complexes can be estimated quantitatively. are scarce, particularly at temperatures more Chloride is the important complexing ligand than 100°C. Also there is disagreement about because it is usually the most abundant anion in the stoichiometry and stability of silver sulfide ore solutions, although numerous other candi complexes reported at 25°C. Cloke (1963) dates as complexing ligands have been proposed measured the solubility of acanthite at 25°C in (Barnes, 1979; Barnes and Czamanske, 1967). weakly to strong alkaline sodium sulfide solu Bisulfide ion is also known to be strong metal tions, and determined that silver was dissolved complexing agents at relatively low temperatures as polysulfide complexes such as Ag(S4)2'-, such as in epithermal environments. Silver is an AgS5S43 and Ag(HS)S42-. Anderson (1962) important ore metal in epithermal deposits, summarized the data of Treadwell and Hepen commonly occurring with electrum. According strick (1949) and Schwarzenbach et al. (1958)

291 292 A. Sugaki et alt which showed that silver was dissolved as the evacuated glass tube method (Sugaki and Ag2S(H2S)2 in weakly acid solution saturated Shima, 1965; Scott, 1974) at 400°C from with H2S at 25°C. From the results of radio elemental silver (99.999% pure) and sulfur chemical measurements at 25°C in the presence (99.99% pure). of excess sulfide (0.02m), Schwarzenbach and The procedures in the solubility experiments Widmer (1966) suggested that Ag2S dissolved were as follows. Ag2S (massive blocks) of 10 as AgHS at pH<4, Ag(HS)2 at pH 5-9 and to 25g and the required calculated amount Ag2S(HS)22 at pH> 9. The solubility of Ag2S of NaOH were loaded in the reaction vessel. was also measured between 100° and 180°C After attaching the valve assemblage, the vessel in H2S-saturated neutral to acid solution by and valve assemblage were evaculated with Ol'shanskii et al. (1959) and Melent'yev et al. a rotary vacuum pump to between 5 and 7 (1970) using a radioactive tracer technique, but X 10-3 torr. H2S gas was bled into the vessel they could not determine stoichiometry of the from a small H2S cylinder connected to the valve silver complexes. Recently, Gammons and assemblage. The amount of H2S gas transfered Barnes (1987) measured solubility of Ag2S in was controlled by warming the cylinder with near-neutral solutions buffered by H2S(aq) and water and by fixing the elapsed time during, HS in the temperature range 25° to 300°C. which the valve was open to the reaction vessel. They estimated that silver was dissolved to form The amount of H2S gas loaded was measured by sulfide complex as Ag(HS)2-. weighing the small cylinder before and after the In this study, we have measured the solu H2S gas injection. After the valve assemblage bility of Ag2S (acanthite, which converts to was again evacuated, H20 was loaded from a argentite above 177'C) in NaOH-H2S-H2O separatory funnel into the reaction vessel using solutions from 25° to 250°C, and have deter a hydraulic pump. Distilled and deionized water mined the stoichiometries and stabilities of the was used, the was boiled for a half hour just silver sulfide complexes. before the experiment in order to remove dis solved gases. The amounts of NaOH, H2S and H20 in the reaction vessel for each experiment SoLUrnm EXPERIMENTS are listed in Table 1. Samples of the solution in Experimental method the reaction vessel were extracted under run Our solubility experiments were carried out conditions by means of the sample tube method at Pennsylvania State University (500-series described by Barnes (1963, 1971), Crerar and experiments) and Tohoku University (100-series experiments) using Barnes volumetric hydro Table 1. Amount of NaOH, H2S and H20 in the reaction thermal systems with 1.11 chrome or teflon vessel for each experiment. lined rocking autoclaves. A detailed description NaOH H20 H20 Run No. of the apparatus and reaction vessels are in (g) (g) (ml) Barnes (1963, 1971, 1981). 102 83.997 74.969 700 Temperature was controlled with solid state 103 4.000 24.734 660 104 1.000 21.827 590 proportional controllers and was measured by 105 0.529 22.815 610 potentiometer within an accuracy of ± 0.1 ° C 106 0.138 25.104 670 107 0.119 24.742 660 using chromel-alumel thermocouples which were 108 19.972 26.646 745 inserted into wells in the top and bottom ends 109 27.050 25.860 733 of the reaction vessel. The temperature gradient 501 0.000 44.575 600 502 104.770 105.840 592 between two ends was maintained at less than 503 104.800 88.330 594 1°C. Total pressure was measured by Heise 504 24.000 62.990 589 505 24.070 22.170 583 Bourdon gauges (600kg/cm2 and 4000psi) 506 24.050 20.360 578 within ± 0.1 kg/cm2. Ag2S was synthesized by Ag2S solubility in sulfide solutions up to 250°C 293

250 0 CD0 0 C/0 CD CC) pH = neutral H2S(aq)/ HS

200 o 00 0 0 coo 00 0 U 0

a)

0 150 0 Cm CD o 0 lm 0 0 CE) L a) CL E a) I 100 0 CD 00 CD 0 m CE)

50

0 0 0000 00

2 3 4 5 6 7 8 9 I0 II

pH

Fig. 1. Temperature pH conditions of solubility experiments. Predominance regions of H2S(aq) and HS and the neutral pH curve are also shown.

Barnes (1976) and Giordano and Barnes (1979) sulfur. The solution was then acidified with or directly by glass syringe connected with HNO3, and was analyzed for Ag by atomic capillary tube to the valve assemblage. Sampled absorption spectrophotometry. XRF was also solutions were passed through a teflon filter used to know approximate concentration of Ag (pore size, 2µm) attached to valve assemblage. before atomic absorption analyses. Generally, two or three samples were extracted at each temperature after a reaction period of Experimental results more than 50h. Each sampling temperature at The measured Ag2S solubilities are shown in 25° or 50°C intervals was approached from both appendix 2 together with estimated activities of lower and higher temperatures. According to H2S and HS-, total concentrations of Na and S Anderson (1962) and Giordano and Barnes and pH calculated by the method given below (1979), reaction rates for forming sulfide com and in the appendix 1. The experiments are plexes are very high, so equilibrium should have arranged in the order of increasing pH in the been achieved in our experiments as was, in fact, table of appendix 2. The measured concentra demonstrated by the reversals. The temper tion of Ag in the solutions range from 0.03 to ature and pH taken out samples in the experi 2140ppm. Maximum solubility was obtained in ments are in Fig. 1. run No. 502 at 250°C with 4.1 m NaHS and Sampled solutions were treated with excess 29.4atm PH2s. In general, the Ag2S solubility NaOH solution followed by 30% H202 in order increases in proportion to temperature, particu to convert reduced sulfur species to sulfate. The larly in the runs at high sulfur concentration. NaOH solution was used to increase solution pH and thereby avoid the precipitation of elemental 294 Sugaki et al.

sodium, and the charge balance for the system. INTERPRETATION OF THE DATA Equilibrium constants used in the calculations General methods are also summarized in the appendix 1. There Activities of H2S(aq) and HS as well as pH is a large disagreement among dissociation con were calculated at each temperature from the stants reported for HS-. We have used the values amount of NaOH, H2S, H20 that was loaded given by Ellis and Giggenbach (1971) because into the vessel, measured partial pressure of H2S these agree with recent spectroscopic measure and ionic equilibria summarized in the appendix ments by Meyer et al. (1983). 1. In solutions of the NaOH-H2S-H2O system Activity coefficients were computed using used in this study, the following reaction occurs an extended Debye-Huckel equation (Helgeson, to form sodium bisulfide (Barnes et al., 1967; 1969), Romberger and Barnes, 1970; Crerar and Barnes, -logy=AZ211/2/(I+aBI12)+bI (5) 1976): NaOH + H2S = NaHS + H20 (1) where y is activity coefficient, Z is charge of the ion, I is ionic strength, A, B and b are Debye The partial pressure of H2S in the table of Huckel coefficients of Helgeson et al. (1981), appendix 2 was estimated from the measured and a is the ion size parameter as summarized by total pressure as, Truesdell (1984). Activities of neutral species such as H2S, NaHS and NaOH were taken to PH2S= Ptotal PH2o (2) be same as that of aqueous CO2 as given by Helgeson (1969). Values from Helgeson (1969) PH2o in equilibrium with an NaHS aqueous solu were also used for activity of H2O. tion at elevated temperature was estimated from the equation given by Haas (1976) for NaCl Estimation of the stoichiometry and stability of aqueous solutions making the assumption that silver sulfide complexes the colligative properties of HS and Cl are In the present experiments, concentration of similar. Activity of H2S(aq) was calculated from S2 is very low even in the runs carried out at the relation, high pH. Thus only H2S and HS are considered to participate in complexing reactions. The dis H2S(gas) = H2S(aq) (3) solution reaction for Ag2S is written generally as,

and the equilibrium constant, Kgas, for equation Ag2S+ xH2S+ yHS = Ag2S(H2S)x(HS)y (3) which is, (6) and the equilibrium constant is

Kgas= aH2S(aq)/OH2S * PH2S (4) acomplex K = x y (7) where a and 0 denote activity and fugacity coef aH2S aHS ficient, respectively. The fugacity coefficient of In the case of reaction (6), possible silver sulfide H2S at each temperature and pressure was esti complexes are Ag2S(H2S), Ag2S(H2S)(HS)-, mated from reduced temperatures (Tr = T/Tj Ag2S(H2S)(HS)22 and Ag2S(HS)22-. Ag2S(H2S)2 and pressure (Pr = P/Pc) according to the method (HS)22 was also considered, but it was omitted of Newton (1935) and Ryzhenko and Volkov finally because its concentration is very low. (1971). Total concentration of sulfur in solu Other possible complexes such as Ag2S(H2S)2 or tion was calculated from the initial concentra Ag2S(HS)33 are not important under the present tion of NaHS and aH2S(aq)as estimated above. experimental conditions although they may Activities of H2S(aq) and HS and pH were become more significant at higher H2S(aq) or calculated from mass balances for sulfur and HS activities. In reaction (6), the total con Ag2S solubility in sulfide solutions up to 250°C 295 centration of silver is represented by summing non-linear 4 dimensional least squares methods all of the complexes, as together with measured Ag solubilities and calculated activities and activity coefficients. mAg(total)= 2K1 * aH2S/yAg2s (H2S) Activity coefficients of silver sulfide complexes were calculated using the extended Debye + 2K2, aH 2S' aHS-I'YAg2S(H2S)(HS) Hiickel equation (5). The d value for complexes + 2K3 ' a 2/ used in these calculations is estimated to be 5.0 H2S, aHS 7Ag2S(H2S)(HS)2 X 10-$ cm for monovalent and 6.0 X 10-s cm for + 2K4' aHS?/yAg 2s(HS)2 (8) divalent species. The calculated equilibrium constants are given in Table 2 and are plotted in where m denote molal concentration, and the Fig. 2 as a function of temperature. equilibrium constants are In the present study sulfide complexes of silver are considered to be di-nuclear complexes, K1 = aAg2s(H2s)/aH 2S (9) but mono-nuclear complexes should be also important. The dissolution reaction to form K2 = aAg2s(H2S)(HS)/a H2S aHS (10) mono-nuclear complexe is written generally as, K3 = aAg 2s(H2s)(HS)22-/aH2S • aHS? (1 1) Ag2S + xH2S + yHS K4 = aAg2s(HS)22/aHS? (12) Y = 2Ag(H2S) X-1 (HS) y+2 2 (13) values for K1, K2, K3 and K4 were obtained 2 2 from equation (8) at each temperature, using and equilibrium constant is expressed as,

-3 (C)

(b) -4 t (d)

Y -5 rn 0 (a) -6 _t - a, Age S H2S = A92S(H2S) -7 b, A92 S H2S + HS = Ag2S(H2S)(HS) C, Age S H2S + 2HS = A92S(H2S)(HS)2 -8 d, Ag2S +2HS = Ag2S(HS)2

100 200 300

Temperature, °C

Fig. 2. Equilibrium constants for silver sulfide complexes. The error bars represent one standard deviation. 296 Sugaki et al.

Table 2. Equilibrium constants for silver sulfide complexes

log K Reaction 25°C 100°C 15 0°C 200°C 25 0°C (a) Ag2S + H2S = Ag2S (H2S) -6.4±0.2 -6.1±0.2 -5.9±0.2 -5.8±0.3 -5.7±0.2 (b) Ag2S + H2S + HS = Ag2S (H2S) (HS) -3.9±0.2 -4.1±0.2 -3.7±0.2 -3.7±0.3 -4.1±0.5 (c) Ag2S + H2S + 2HS = Ag2S (H2S) (HS)22 -4.3±0.4 -3.6±0.2 -3.4±0.2 -3.3±0.2 -3.2±0.2 (d) Ag2S + 2HS = Ag2S (HS)22 -5.2±0.2 -4.8±0.2 -4.7±0.2 -5.0±0.4 -5.2±0.4

2 K = acomplex (14) plexes, we could not determine strictly whether aH2sx aHS_v silver dissolves as mono-nuclear complexes or Schwarzenbach and Widmer (1966) reported di-nuclear complexes. However, observed solu that Ag2S dissolved as AgHS at pH< 4, Ag(HS)2 bility is not explained without di-nuclear com at pH 5-9 and Ag2S(HS)22 at alkaline condi plexes such as Ag2S(H2S)(HS) and Ag2S(HS)22-. tions at 25'C. The dissolution reactions to form Furthermore, the calculated total concentration AgHS and Ag(HS)2 are written as, of silver approaches closer to the observed solubility when considering that all silver dis Ag2S + H2S = 2AgHS (15) solves as di-nuclear complexes than mono Ag2S + H2S + 2HS = 2Ag(HS)2 (16) nuclear species. Therefore complexing reactions for Ag2S are considered as given in equation (6) The left side of these reactions are same as the as a matter of convenience. equations to form sulfide complexes of Ag2S (H2S) and Ag2S(H2S)(HS)2-. In the method Concentration of silver sulfide complexes adopted in this study to estimate silver com The concentration of silver sulfide com

-3 200 °C 102 IS= 1.0 M E E CL Total Ag CL -4 C 0 I0 a C

C 0 0 a) 9~ 4 U -5 C C 100 a) 0 U 99 U ~s C ~ti ~s y 0 rn U 0 -6 10-I CO

-7 Oil Ipt~ Q 10-2

3 4 5 6 7 8 9 I0 pH

Fig. 3. Concentration of silver sulfide complexes as a function of pH at 200°C and ES=1.0m. Measured solubilities of Ag2S are normalized to ES=1.0m. Error bars represent one standard deviation.

9 Ag2S solubility in sulfide solutions up to 250°C 297

plexes was calculated at 200°, 100° and 25°C function of pH and for a total reduced sulfur using the equilibrium constants KI to K4. The concentration ES (mH2S +mHS-) =1.0m. Solubil results are shown as curves in Figs. 3, 4 and 5 as a ities determined in our study are also plotted in

-3 100 °C zS=I.O m 102 E 0. a C -4 Total Ag 0 I0 c I 0 L 4 0 C 4 as -5 C U °' 100 C 0 V C V 99 0 U 0 -6 N 10-1 19mss, /1 -7 C~ s Q' 10-2

3 4 5 6 7 8 9 10 pH

Fig. 4. Concentration of silver sulfide complexes as a function of pH at 100°C and ES=1.Om.

-3 25 °C 102 ES= 1.0 m E 0. 0. c -4 10 0 0 } 0 } 0 L Tota I Ag } c c -5 c 100 c 0 0 U U v CD 99 0 -6 10-' 4 sJ' ,y sJ ~tisi -7 10-2 Q

3 4 5 6 7 8 9 10 I I pH

Fig. 5. Concentration of silver sulfide complexes as a function of pH at 25°C and ES=1.Om.

0J 298 Sugaki et al.

these figures. However since the total concentra less than 5.1 at 200'C, 4.3 at 100'C and 4.7 tion of sulfur varied from 0.28 to 4.36m at at 25°C, Ag2S(H2S) is the dominant species. 25°C, the solubility data have been normalized to Ag2S(HS)22 dominates at high pH greater than ES=1.0m. In Figs. 3, 4, and 5, calculated total 8.3 at 200T, 7.3 at 100'C and 7.8 at 25'C. concentration of silver is shown as a heavy line. The intermediate region in pH is dominated by Measured solubility values (normalized to IS the complexes Ag2S(H2S)(HS) and Ag2S(H2S) =1 .0m) plot on or very close to this line. Gene (HS)22 but, at 25°C and ES=1.0m, the con rally speaking, sulfide complexes such as Ag2S centration of Ag2S(H2S)(HS)22 does not exceed (H2S), Ag2S(H2S)(HS)-, Ag2S(H2S)(HS)22 and that of Ag2S(H2S)(HS)-, as shown in Fig. 5. Ag2S(HS)22 become progressively dominant in The total reduced sulfur concentration (ES) this order with increasing pH. At low pH, i.e., also has an affect on the concentration of silver

100 b

80 a

C

60 a: A92S(H2S)

40 25 °C b: A9zS(H2S)(HS) pH= 7.0 C: Ag2S(H2S)(HS)2 20 d: A92S(HS)2 d 0

100 N a) C C x a b a) 80 a c. E b 0 60

a)

N 40 1000C 150 °C pH =6.I pH= 5.8 c a> U 20 a) d a d 0

100

80 a a b C

60 b 40 200°C 2500C pH =5.6 pH= 5.5 20 C d 0 -4 -3 -2 1 0 1 -4 -3 -2 -I 0 I

log Es' M

Fig. 6. The percentage distribution of sulfide complexes of silver as a function of total sulfur concentration from 25° to 250°C and at neutral pH. Ag2S solubility in sulfide solutions up to 250°C 299

sulfide complexes. The relative concentrations (H2S) becomes more important. of silver sulfide complexes is shown as a func tion of IS in Fig. 6 for neutral pH at temper APPLICATIONS atures of 25°, 100°, 150°, 200° and 250°C. As IS of the solution increases, sulfide com Deposition of Ag2S plexes such as Ag2S(H2S), Ag2S(H2S)(HS) and As a means of evaluating the contribution of Ag2S(H2S)(HS)22 become progressively pre silver sulfide complexes in ore solutions, the pre dominant in order. Ag2S(HS)22 is not impor cipitation mechanisms of Ag2S is considered. tant at neutral pH, as is shown in Fig. 6, although The relative importance of sulfide and chloride it does become dominant at higher pH. With complexing is then discussed. Deposition of increasing of temperature, the dominance region Ag2S is in response to changes in temperature, of each complexes shifts to higher IS. pH and total reduced sulfur concentration. According to Ohmoto (1972) and Barnes Figure 7 shows Ag2S solubility as a function of (1979), the total reduced sulfur concentration pH at 250°, 100° and 25°C for ES=1.Om and in ore solutions is typically between 0.1 and 0.1 m. It is shown in Figs. 3, 4 and 5 that the 0.0005m. At 25°C, Ag2S(H2S)(HS) con maximum solubility is observed at pH around stitutes about 90% of the silver sulfide com equal activities of H2S(aq) and HS as given in plexes at neutral pH and for a total concentra Fig. 1. The maximum solubilities of Ag2S for tion of reduced sulfur of 0. 1 m. It also is the ES=1.Om decreases from 38.3ppm at 250°C to dominant species up to 150°C, but its propor 20.1 ppm of 100° C. This means that about half tion decreases to about 60%. Ag2S(H2S) be of the dissolved silver is precipitated as Ag2S in comes dominant at higher temperatures, and this 150° temperature interval. However the almost all'of the silver is dissolved as Ag2S(H2S) effect of temperature on deposition is quite at 250°C. Where IS is less than 0.1m, Ag2S different depending on the pH of the solution.

-3 102

E 0 C . 0 0. -4 ES =1.0m 0 .1 1%. I0 C 0 C 0 U C C 0 -5 a) 0 100 U / C 0 Q 0

-6 Q 0 Es =0.I M 10-1 .00 1000

-7 2500C 10-2 Neutral pH 100 °C -- 250 100 25°C 25 °C "'•""

-8 3 4 5 6 7 8 9 10

pH Fig. 7. Ag2S solubility as a function of temperature, pH and total reduced sulfur concentration. 300 Sugaki et al.

For ES=1.Om, if the temperature falls from changed by proton metasomatic reactions 2500 to 100°C at a constant pH of 7.5, about between solution and wall rock such as forming 80% of the silver will precipitate, but at pH= 6.5, sericite from K-feldspar or kaolinite from this same drop of temperature causes an increase sericite, or their inverse reactions. Also where in solubility (Fig. 7). The temperature effect on an ascending hydrothermal solution meets solubility at lower total sulfur concentration is meteoric water, the pH of the solution changes rather small but about 50 to 70% of dissolved by mixing or oxidation (e.g., conversion of silver will precipitate in response to a drop of H2S to H2SO4, Seward, 1973), and Ag2S is temperature from 250° to 25°C at pH< 6 in precipitated. a solution with IS =0.1 m. A decrease of the total reduced sulfur con Since the maxima in silver solubility at all centration of an ore forming fluid will also cause temperatures approximately coincide with the precipitation of Ag2S. When IS decreases by pH range of the majority of hydrothermal ore one order of magnitude, the solubility of Ag2S solutions, any process which results in a change decreases by roughly two orders of magnitude as of pH to either more acid or more alkaline con shown in Figs. 7 and 8. Therefore, the IS effect ditions would cause the precipitation of Ag2S. on ore deposition seems to be large. IS in ore For example, at ES=1.Om, change of pH from solutions can be decreased in several ways. If an 7.5 to one unit acid and alkaline at 250°C causes ore solution is mixed with oxygenated meteoric deposition of about 90 and 70% of dissolved water, IS of the solution will decrease because silver, respectively. The pH of the solution is of dilution and also oxidation of reduced sulfur

200 °C A, -2 1CI =1.0m 103

-3 102 Sulfide h c complexes 0 dominant o -4 10 c X a L c 0 0 L -5 Chloride c 100 c 0 complexes a, U U dominant C pH. 4 0 Q -6 10-1 0

Q M 0 5 -7 10-2

6 -8 10-3 7

-9 10-4 8

-4 -3 -2 -1 0 I log (Total reduced sulfur, m ) Fig. 8. Ag2S solubility at 200°C in a solution containing ECl -=1.0m. Ag2S solubility in sulfide solutions up to 250°C 301

species to sulfate. Boiling of the solution should for transporting silver in ore forming solutions also be important mechanism to decrease total of most epithermal deposits, as is the case for reduced sulfur content. According to Drum gold (Seward, 1973). mond and Ohmoto (1985), dissolved gas such as H2S will rapidly concentrate into the vapor SUMMARY phase during boiling, and this separation results in a decrease of ES and a slight increase of pH The solubility of Ag2S from 25° to 250°C in the residual solution. The residual solution in the NaOH-H2S-H20 system is a function of will also have an enhanced chloride concentra temperature, pH and total reduced sulfur con tion, depending on how much H20 is lost to the centration (ES). From calculations of ionic vapor, which could have an effect on chloride vs. equilibria in the solutions, we have determined sulfide complexing. that Ag2S dissolves to form sulfide complexes such as Ag2S(H2S), Ag2S(H2S)(HS)-, Ag2S(H2S) Comparison with chloride complexes (HS)22 and Ag2S(HS)22-. The complex-forming The relative importance of silver chloride reactions and their equilibrium constants at 25°, and sulfide complexes is compared for 200°C 100°, 150', 200° and 250°C are given in Table 2. and ECl-=1.Om in Fig. 8 using the Seward's The concentration of each silver sulfide complex (1976) formation constants for silver chloride changes with temperature, pH and ES as shown complexes, and Helgeson's (1969) Ksp value for in Figs. 3, 4, 5 and 6. Silver sulfide complexes Ag2S. In this figure, contours with negative predominate over chloride complexes for slopes in the low ES region are dominated by chloride and sulfur concentrations, temperature chloride complexes, and those with positive and pH that are estimated for most epithermal slopes are due to the contribution of sulfide fluids. In such cases, decrease of temperature, complexes. change of pH away from neutral and decrease in According to Hedenquist and Henley (1985), ES cause precipitation of Ag2S. Among them, ore solutions which produce most epithermal the pH and ES effects are the important. deposits have salinities of less than 10 equivalent wt% NaCl. In the case of gold-silver deposits, Acknowlegment-The initial experiments were per the majority have salinities of less than 3 equi formed by A. Sugaki and S. D. Scott at the Pennsylvania valent wt% NaCl. The concentration of Cl of State University in the laboratory of Prof. H. L. Barnes who is heartily thanked for the use of his facilities, 1.0m in Fig. 8 corresponds approximately to 5.5 funding and his valuable advice. Our work has also equivalent wt% NaCl. As mentioned previously, benefitted from discussions with Prof. M. Ichikuni of the the total reduced sulfur concentration in most Tokyo Institute of Technology. This research was ore solutions is estimated to be in the range of supported financially by grants from The Ministry of between 10-3 and 10-2m (Barnes, 1979). For Education, Science and Culture of Japan to A. Sugaki, this ES, both chloride and sulfide complexes are the Society for the Promotion of Science of Japan to A. S. and S. D. S the Advanced Research Project Agency important under slightly acid condition (pH=4 to H. L. Barnes and the Natural Sciences and Engineer and 5 in Fig. 8 relative to a neutral pH at 200°C ing Research Council of Canada to S. D. S. of 5.6). Sulfide complexes become dominant at neutral to slightly alkaline pH. However, chloride concentrations in ore solutions respon REFERENCES sible for epithermal gold-silver deposits are Anderson, G. M. (1962) The solubility of PbS in H2S generally less than 1.Om and decrease in Cl water solutions. Econ. Geol. 57, 809-828. concentration results in a decrease of silver Barnes, H. L. (1963) Ore solution chemistry I. Experi solubility. Therefore considering the relative mental determination of mineral solubilities. Econ. concentrations of chloride and total reduced Geol. 58, 1054-1060. sulfur, sulfide complexing must be significant Barnes, H. L. (1971) Investigations in hydrothermal 302 Sugaki et al.

sulfide solutions. Research Techniques for High Helgeson, H. C. (1969) Thermodynamics of hydro Pressure and High Temperature 317-335. Edited by thermal systems at elevated temperatures and pres Ulmer, G. G. Springer-Verlag: Berlin. sures. Am. J. Sci. 267, 729-804. Barnes, H. L. (1979) Solubilities of ore minerals. Geo Helgeson, H. C., Kirkham, D. H. and Flowers, G. C. chemistry of Hydrothermal Ore Deposits, 2nd Ed. (1981) Theoretical prediction of the behavior of 404-460. Edited by Barnes, H. L. Wiley Inter aqueous electrolytes at high pressures and temper science: New York. atures: Calculation of activity coefficients, osmotic Barnes, H. L. (1981) Measuring thermodynamically coefficients, and apparent molal and standard and interpretable solubilities at high pressures and temper relative partial molal properties to 600°C and 5kb. atures. Chemistry and Geochemistry of Solutions at Am. J. Sci. 281, 1249-1516. High Temperatures and Pressures. Physics and Henley, R. W. (1973) Solubility of gold in hydro Chemistry of the Earth 13/14, 113-132. Edited thermal chloride solutions. Chem. Geol. 11, 73-87. by Rickard, D. and Wickman, F. Pergamon Press: Melent'yev, B. N., Ivanenko, V. V. and Pamfilova, L. A. Oxford. (1970) Solubility of some ore-forming sulfides Barnes, H. L. and Czamanske, G. K. (1967) Solubilities under hydrothermal conditions. Geochem. Internat. and transport of ore minerals. Geochemistry of 7,416-460. Hydrothermal Ore Deposits 334-381. Edited by Meyer, B., Ward, K., Koshlap, K. and Peter, L. (1983) Barnes, H. L. Holt, Rinehart and Winston: New York. Second dissociation constant of hydrogen sulfide. Barnes, H. L., Romberger, S. B. and Stemprok, M. Inorg. Chem. 22, 2345-2346. (1967) Ore solution chemistry II. Solubility of Naumov, G. B., Ryzhenko, B. N. and Khodakovsky, I. L. cinnabar in sulfide solutions. Econ. Geol. 62, 957 (1974) Handbook of Thermodynamic Data. Nat. -982 . Tech. Info. Service PB226-727, U. S. Dept. of Com Cloke, P. L. (1963) The geologic role of polysulfides merce: Springfield. Part II, The solubility of acanthite and covellite in Newton, R. H. (1935) Activity coefficient of gases. sodium polysulfide solutions. Geochim. Cosmochim. Ind. Engng. Chem. 27,302-306. Acta 27, 1299-1319. Ohmoto, H. (1972) Systematics of sulfur and carbon Crerar, D. A. and Barnes, H. L. (1976) Ore solution isotopes in hydrothermal ore deposits. Econ. Geol. chemistry V. Solubility of chalcopyrite and chal 67, 551-578. cocite assemblages in hydrothermal solution at 200° Ol'shanskii, Ya. I., Ivanenko, V. V. and Khromov, A. V. to 350°C. Econ. Geol. 71, 772-794. (1959) The solubility of silver sulfide in aqueous Drummond, S. E. and Ohomoto, H. (1985) Chemical solutions saturated with hydrogen sulfide. Dokl. evolution and mineral deposition in boiling hydro Akad. Nauk SSSR 124, 410-413. thermal systems. Econ. Geol. 80, 126-147. Romberger, S. B. and Barnes, H. L. (1970) Ore solution Ellis, A. J. and Giggenbach, W. F. (1971) Hydrogen chemistry III. Solubility of CuS in sulfide solutions. sulfide ionization and sulfur hydrolysis in high tem Econ. Geol. 65, 901-919. perature solution. Geochim. Cosmochim. Acta 35, Ryzhenko, B. N. and Volkov, V. P. (1971) Fugacity 247-260. coefficient of some gases in broad range of temper Gammons, C. H. and Barnes, H. L. (1987) Stability of atures and pressures. Geochem. Internat. 8, 468-481. silver bisulfide complexes under hydrothermal condi Schwarzenbach, G., Giibeli, O. and Ziist, H. (1958) tions: Applications to metal transport in geothermal Thiocomplexe des Silbers and die Loslichkeit von systems (abstr.). Geol. Soc. Am. Abstr. Programs Silbersulfid. Chimia 12, 84-86. 671. Schwarzenbach, G. and Widmer, M. (1966) Die Loslich Giordano, T. M. and Barnes, H. L. (1979) Ore solution keit von Metallsulfiden II. Silbersulfid(I). Helv. Chim. chemistry VI. PbS solubility in bisulfide solutions to Acta 19, 111-123. 300°C. Econ. Geol. 74, 1637-1646. Scott, S. D. (1974) Experimental methods in sulfide Haas, J. L. Jr. (1976) Physical properties of the coexist synthesis. Sulfide Mineralogy. Miner. Soc. Am. Short ing phases and thermochemical properties of the H2O Course Notes 1, Sl-S38. Edited by Ribbe, P. H. Min. component in boiling NaCl solutions. Preliminary Soc. Am.: Washington D. C. steam tables for NaCl solutions. U.S. Geol. Surv. Seward, T. M. (1973) Thiocomplexes of gold and the Bull. 1421-A, Al-A73. transport of gold in hydrothermal ore solutions. Hedenquist, J. W. and Henley, R. W. (1985) The Geochim. Cosmochim. Acta 37, 379-399. importance of CO2 on freezing point measurements Seward, T. M. (1976) The stability of chloride com of fluid inclusions: Evidence from active geothermal plexes of silver in hydrothermal solutions up to systems and implications for epithermal ore deposi 350°C. Geochim. Cosmochim. Acta 40, 1329-1341. tion. Econ. Geol. 80, 1379-1406. Skinner, B. J. and Barton, P. B. Jr. (1973) Genesis of Ag2S solubility in sulfide solutions up to 250°C 303

mineral deposits. Ann. Rev. Earth Plant. Sci. 1, 1872-1879. 183-211. Truesdell, A. H. (1984) Introduction to chemical Sweeton, F. H., Mesmer, R. E. and Baes, C. F. Jr. (1974) calculations. Fluid-mineral equilibria in hydro Acidity measurements at elevated temperatures VII. thermal systems. Reviews in Economic Geology 1, Dissociation of water. J. Sol. Chem. 3, 191-213. 1-8. Edited by Robertson, J. M. Society of Econo Sugaki, A. and Shima, H. (1965) Synthetic sulfide mic Geologists: El Paso. minerals (I). Yamaguchi Daigaku Kougakubu Ken Truesdell, A. H. and Singers, W. (1974) The calcula kyuu Houkoku 15, 15-31. tion of aquifer chemistry in hot-water geothermal Treadwell, W. D. and Hepenstrick, H. (1949) Uber die systems. J. Res. U. S. Geol. Surv. 2, 271-278. Loslichkeit von Silbersulfid. Hely. Chim. Acta 32 ,

APPENDIX 1 The charge balance for this system is also written,

Calculation of activities of H2S and HS and pH in the mNa++mH+ = mOH-+mHS-+2mS2 (A-6) system NaOH-H2S-H20 At elevated temperatures, aqueous species included and activity of Na+ is expressed as, in the system NaOH-H2S-H20 are H20, H+, OH-, Na+, H2S(aq), HS-, S2-, NaOH(aq) and NaHS(aq). From the KH2O • aH2O total mass balance on sulfur, aNa+ yNa+ ' + mHS aH+ , yOH

MS(total) = mH2S +mHS + mS2 + mNaHS (A-1) 2KHS • mHS * ^fHS aw + (A-7) aH+ • yS2 yH+ here m denotes molal concentration. Equilibrium con stants for associated species are written as, Also, from the mass balance on sodium, total concentra KH2S = aW • aHS / aH2S (A-2) tion of sodium can be calculated as,

KHS = aW * aS2 / aHS (A-3) mNa(total) = mNa+ + mNOOH + mNaHS

KNaHS = aNa+, aHS / aNaHS (A-4) aNa+, aOH = mN a+ + KNaOH NNaOH Using the relation ai = mi yi (yi: activity coefficient) , mHS can be written as aHS' + aNa+ (A-8) _ aH+. yHSm KNaHS * YNaHS HS MS(total) / K + 1 H2S yH2S Here KH2O and KNaOHdenote + aNa+ 7HS + KHS • yHS KN _ (A-5) KH aHS * yNaHS aW , ys2 2O = aW * aOH / aH2O (A-9)

Table A-1. Selected equilibrium constants used in the calculations.

log K Reactions 25°C 100°C 15 0°C 200°C 2500C H2S (gas) = H2S (aq) -1 .00 -1.43 -1 .54 -1 .56 -1.53 a H20 = H++ OH -13 .99 -12.23 -11 .59 -11 .21 -11 .08 b -7 112S (aq) = H++ HS .28 -6 .38 -6 .62 -7 .10 -7 .70 a HS =H++S2 -17.00 -16 .10 -15 .40 -15.10 -14 .80 c NaOH (aq) = Nat + OH 1.60 0.55 0.45 0.15 -0 .25 d NaHS (aq) = Na+ + HS 1.60 0.55 0.45 0.15 -0 .25 d

a Naumov et al. (1974) b Sweeton et al. (1974) c Ellis and Giggenbach (1971) d Truesdell and Singers (1974), Data for NaCI is used. 304 Sugaki et al.

KNaOH = aNa+* aOH / aNaOH (A-10) given in Table A-1.

Putting an approximate value of aH+ (=10-pH) into equation (A-5), mHS can be calculated. Using mHS Ar iiNDlx 2 and aNa+calculated from (A-7), the total concentration of Na can be obtained from equation (A-8). The calcu The measured Ag2S solubilities are shown in Table lation was repeated using a microcomputer until the A-2 together with estimated activities of H2S and HS-, calculated total Na agreed with the amount of loaded total concentrations of Na and S and pH calculated by NaOH. Equilibrium constants used in the calculation are the method given in appendix 1.

Table A-2. Ag2S solubilities in NaOH-IJ2SH2O solutions. Experimental conditions and calculated pH and activities of H2S (aq) and HS are also shown.

Sample Temp. ENa aHS rH fH aH2S ES pH EAg Standard 2S 2S deviation No. (,C) (m) (atm) (atm) (m) (ppm) (ppm) 501A 25 0.000 0.0002 9.5 8.6 0.86 0.85 3.68 0.10 0.02 501C 100 0.000 0.0005 19.8 18.1 0.67 0.73 3.26 0.10 0.02 501D 150 0.000 0.0004 23.9 22.4 0.64 0.74 3.38 0.20 0.04 501E 200 0.000 0.0002 24.8 23.8 0.66 0.75 3.62 0.20 0.04 501F 250 0.000 0.0001 26.0 25.3 0.75 0.84 3.89 0.30 0.05

106C 150 0.005 0.005 20.5 19.4 0.56 0.64 4.55 0.10 0.01 106D 200 0.005 0.005 19.8 19.2 0.52 0.61 5.05 0.14 0.06

107B 100 0.005 0.004 17.8 16.5 0.61 0.67 4.22 0.12 0.01 107C 150 0.005 0.004 20.5 19.4 0.56 0.64 4.49 0.18 0.03 107D 200 0.005 0.004 22.3 21.5 0.58 0.68 4.95 0.18 0.04 107E 250 0.005 0.004 22.2 21.7 0.64 0.73 5.48 0.31 0.01 107G 250 0.005 0.004 19.4 19.0 0.56 0.64 5.54 0.21 0.02 1071 150 0.005 0.004 15.4 14.8 0.74 0.49 4.60 0.16 0.03 107J 100 0.005 0.004 12.0 11.4 0.42 0.46 4.38 0.10 0.02

105A 25 0.022 0.018 6.5 6.1 0.61 0.62 5.77 0.45 0.04 105B 100 0.022 0.018 16.4 15.3 0.56 0.63 4.89 0.59 0.08 105C 150 0.022 0.018 19.2 18.2 0.52 0.62 5.15 0.77 0.12 105D 200 0.022 0.017 21.1 20.2 0.55 0.65 5.60 1.23 0.09 105E 250 0.022 0.016 21.5 21.0 0.62 0.72 6.11 0.85 0.25 105G 250 0.022 0.016 21.7 21.2 0.63 0.72 6.11 0.77 0.18 105H 200 0.022 0.017 19.3 18.7 0.51 0.61 5.63 0.81 0.07 1051 150 0.022 0.018 16.6 15.9 0.46 0.54 5.21 0.44 0.09 105J 100 0.022 0.018 13.2 12.5 0.46 0.52 4.98 0.24 0.07

104C 150 0.042 0.032 17.7 16.8 0.48 0.59 5.44 2.53 0.32 104D 200 0.042 0.031 18.6 18.0 0.49 0.60 5.90 2.57 0.48

504A 25 1.00 0.53 15.3 13.1 1.30 2.02 6.89 51 9 504C 100 1.00 0.47 31.3 27.2 0.99 1.84 6.05 144 26 504D 150 1.00 0.42 39.7 35.3 1.00 1.86 6.24 229 38 504E 200 1.00 0.37 47.1 43.4 1.15 1.96 6.62 480 51 504F 250 1.00 0.32 54.3 51.0 1.44 2.13 7.04 669 50

108C 150 0.67 0.31 8.9 8.5 0.24 0.90 6.73 9.07 0.26 108E 200 0.67 0.28 5.7 5.7 0.15 0.81 7.37 10.2 1.4 108F 250 0.67 0.24 9.2 9.1 0.26 0.90 7.65 11.0 1.7 108H 200 0.67 0.27 8.2 8.1 0.22 0.87 7.22 10.2 2.3 108K 100 0.67 0.34 5.4 5.3 0.19 0.85 6.63 7.88 0.73

103A 25 0.15 0.11 1.4 1.4 0.14 0.28 7.18 0.45 0.12 103C 200 0.15 0.088 21.1 20.4 0.55 0.75 6.31 1.75 0.25 103D 250 0.15 0.079 24.6 23.9 0.71 0.89 6.75 1.53 0.36

502A 25 4.10 1.35 8.2 7.6 0.75 4.36 7.54 250 83 502C 100 4.10 1.16 16.5 15.4 0.54 4.35 6.71 630 78 502D 150 4.10 1.00 21.5 20.3 0.55 4.38 6.88 920 238 502E 200 4.10 0.88 25.7 24.6 0.64 4.43 7.25 1400 310 Ag2S solubility in sulfide solutions up to 250°C 305

Table A-2 (Continued)

Sample Temp. ENa aHS PH2S fH2S aH2S ES pH EAg Standard deviation No. (°C) (m) (atm) (atm) (m) (ppm) (ppm) 502F 250 4.10 0.80 29.1 28.2 0.80 4.48 7.70 2140 194 109A 25 0.91 0.49 1.5 1.5 0.15 1.03 7.80 8.2 2.6 109B 100 0.91 0.43 2.4 2.4 0.08 0.98 7.07 15.4 0.5 109C 150 0.91 0.39 2.9 2.9 0.08 0.98 7.29 15.0 0.7 109F 250 0.91 0.29 1.5 1.5 0.06 0.94 8.39 15.7 1.6 109H 200 0.91 0.34 2.7 2.7 0.07 0.97 7.77 12.1 3.2 1091 100 0.91 0.43 1.2 1.2 0.04 0.94 7.37 11.6 1.0

505A 25 1.01 0.53 1.0 1.0 0.10 1.09 8.01 15.3 5.4 505C 100 1.01 0.47 1.7 1.7 0.06 1.06 7.26 20.0 8.6 505D 150 1.01 0.42 2.0 2.0 0.06 1.06 7.48 43.0 4.9 505F 250 1.01 0.32 3.6 3.6 0.11 1.09 8.15 138 45

102C 200 2.85 0.72 0.8 0.8 0.06 2.86 8.17 49.4 8.2 102D 250 2.85 0.63 2.2 2.2 0.13 2.88 8.40 70.2 10.2 102E 200 2.85 0.73 3.4 3.4 0.11 2.90 7.92 109 24 102G 25 2.85 1.10 1.9 1.9 0.19 2.95 8.05 37.7 4.2

506A 25 1.02 0.53 0.0004 1.02 10.44 2.10 1.5 506C 100 1.02 0.47 0.0026 1.02 8.63 3.80 2.8 506D 150 1.02 0.42 0.0066 1.02 8.42 10.3 0.8 506F 250 1.02 0.31 0.029 1.02 8.72 29.2 0.3

503A 25 4.09 1.32 0.0007 4.05 10.58 24.0 7.8 503C 100 4.09 1.13 0.0058 4.05 8.67 35.0 7.2 503D 150 4.09 0.98 0.018 4.05 8.36 42.0 12.1 503E 200 4.09 0.86 0.044 4.05 8.40 142 20 503F 350 4.09 0.77 0.093 4.05 8.62 336 15