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Solubility product constants of covellite and a poorly crystalline copper sulfide precipitate at 298 K
Geochimica et Cosmochirnk.a Acla Vol. 53. pp. 229-236 Copyright © 1989 Pergamon Press plc. Printed in U.S.A.
0016-7037/89/$3.00 + .00
Solubility product constants of covellite and a poorly crystalline copper sulfide precipitate at 298 K* DAM1AN SHEAt and GEORGE R. HELZ Department of Chemistry and Biochemistry, University of Maryland, College Park, MD 20742, U.S.A. ( Received Apri120, 1988; accepted in revised form August 8, 1988) Abstract--The equilibrium constant at 25"C for the following reaction has been measured in NaC1 media by an indirect method: CuS(COv) + H+(aq) ~ Cu2+(aq) + HS-(aq),
K~ = Mcu2.M,s-(10 +p")
where CuS(cov) designates synthetic covellite. Values ofpK~ are 21.39, 21.04 and 20.95 at NaCI = 0.2, 0.7 and 1.0 M, respectively; the uncertainty in these K~ values is _+0.15. The free energy of formation of covellite, for which published values are discordant, is calculated to be -11.83 _+0.4 kcal/mole at 298 K (-49.50 + 1.7 kJ/mole). This value is obtained by extrapolating the measured pK,p values to infinite dilution with corrections for CI- complexing. Applying similar CI- complexing corrections, based on recent measurements by Seward, to preyiously published solubility data for galena yields a revised pK°wfor galena of 12.78. A poorly crystalline precipitate, obtained by mixing Cu 2÷ and HS- solutions, yielded a reversible solubility product 3 orders of magnitude greater than that of covellite but about 3 orders of magnitude less than that of a truly amorphous phase, super-cooled liquid CuS. The poorly crystalline phase has not been studied previously. Its bulk composition was Cu ,.,aS, but microprobe analysis revealed that it was a partially exsolved mixture of roughly Cu L, ,S and Cu t.32S (similar to known blaubleibender covellites). It was kinetically unstable, and converted to coveUitewhen thermally annealed or when exposed to polysulfide solutions. Because of its instability, a material of this nature is unlikely to account for the amorphous copper sulfide alleged to occur in the Red Sea Brine deposits. However, it is possible that on short time scales dissolved Cu in sulfidic waters is controlled by metastahle, rather than stable phases, as is known to be the case with dissolved Fe. INTRODUCTION
system. We have measured the solubility product ofcovellite and of a poorly crystalline, metastable phase formed by rapid mixing of CuCI2 and NariS solutions. Combining these measurements with data from the literature, we have estimated the solubility of some additional phases. To determine solubility products, we employed the indirect method of UHLER and HELZ (1984). This involves using a chelating agent, Y, to enhance the solubility of a sulfide mineral sufficiently to enable convenient measurement. Equilibrium constants for the following three reactions are determined and then multiplied to get the solubility product constant, K,p.
M U C H C U R R E N T A T T E N T I O N i s being devoted to understanding the geochemical behavior of trace metals in low temperature, sultidic environments (BOULEGUEet al., 1982; JACOBS and EMERSON, 1982; EMERSON et al., 1983; KREMLING, 1983; DYRSSEN, 1985; CARIGNAN and NRIAGU, 1985; JACOBS et aL, 1985; GAmLARD et al., 1986). A m o n g the m a n y problems is the need to establish what phases control trace metal solubilities in these environments. Interfaces between deep-seated sulfidic waters and oxic surface waters are important sites of sulfide mineral deposition. Such interfaces occur in groundwaters around sulfide ore deposits, in pore waters of modern sediments, and in fjords or other types of anoxic marine basins. Mixing of waters across such interfaces can readily generate very high degrees of supersaturation with respect to trace metal sulfide minerals, permitting deposition ofmetastable phases. Indeed, metastable blaubleibender covellite phases, in which the C u / S ratio exceeds 1.0, are often found in environments of supergene enrichment. BROCKAMP el al. (1978) even report that an amorphous copper-bearing sulfide phase occurs in deposits from the Red Sea thermal brines. Thus it is conceivable that metastable Cu-S phases may control copper solubilities at temperatures occurring near the Earth's surface. In this regard, copper would be analogous to iron, which is known to be controlled by metastable phases (e.g. DAVISON, 1980). This paper concerns the solubility of phases in the Cu-S
CuS(s) + H Y 3- ~ CuY 2- + H S -
Kdiu
(1)
Kcuv*
(2)
H2Y 2 - ~:~ H + + H Y 3-
Ka3
(3)
CuS(s) + H + ~ Cu 2+ + H S -
K~
(4)
CuY 2- + 2H + ~ Cu 2+ + H2Y2-
In measurements on the poorly crystallized precipitate, we used the chelator, ethylenediaminetetraacetic acid ( E D T A ) . For covellite, which is much less soluble, we had to use the stronger chelator, trans 1-2 diaminoeyclohexyltetraacefic acid (DCTA). EXPERIMENTAL Materials
* This paper was part ofan invited presentation at the MERI conference. ' Present address: Department of Chemistry, Southeastern Massachusetts University, North Dartmouth, MA 02747, U.S.A.
Covellite was synthesized by mixing about 10 g of a 1:1 molar ratio of 99.99% pure Cu powder and 99.999% pure S powder in an evacuated pyrex vessel. This mixture was allowed to react for one day at room temperature and was then heated to 2500C for 30 days. 229
230
D. Shea and G. R. Helz (H4 DCTA) was twice recrystallized from water, dried and then standardized by potentiomctric titration against a standard Cu(II) solution.
Methods
30 °
40 °
50 °
60 °
2e FIG. 1. X-ray powder diffraction patterns for (a) covellite and (b) poorly crystalline CuS (Cu ka radiation ).
The product was washed and in some cases leached in a deoxygenated 0. l M DCTA solution in a N2-filled glove box. The X-ray diffraction pattern for this sample exhibited sharp covellite peaks (see Fig. 1). The peak positions and intensities were in excellent agreement with those of POTTER and EVANS (1976), indicating a well-crystallized material. Ten covellite samples of approximately 0. ! g (weighed to 0.1 mg) were dissolved in warm aqua regia and Cu analysis was performed using plasma emission spectrometry to determine the weight percent of Cu in these samples. The average atomic Cu/S ratio was found to be 1.008 __.0.013. The poorly crystalline CuS was made in an N2-filled glove box by mixing 500 mL of N2-purged 1.00 M CuCI2 (pH = 4) with 500 mL of N2-purged 1.00 M NaFIS (pH = 8-11). After 10 minutes of shaking, the solution was filtered and the dark blue residue was washed with distilled water and allowed to dry under vacuum in the antechamber of the glove box. The X-ray diffraction pattern of this product was obtained within a few hours and is also shown in Fig. 1. This pattern exhibited a very high background with broadened peaks of low intensity. In addition, the strong ( 108 ) and ( ! 16) covellite reflections, occurring at 55* and 61.5" 2# respectively, in Fig. 1 are absent in the poorly crystalline CuS. The absence of these two reflections is characteristic of blaubleibender covellite (PUTNIS et al., 1977). The single particle electron diffraction patterns of approximately 1 #m particles were also very poor, exhibiting diffuse rings about the electron beam. Copper analysis was performed on the poorly crystalline product after washing with 0.1 M Naris (pH = 12 ) to remove any elemental sulfur (S °) as polysulfides. An atomic Cu/S ratio of 1.189 + 0.015 was determined. Comparison of this value with the corresponding value determined before the Naris wash suggests that the poorly crystalline material contained about 3 tool% elemental sulfur. Single particle X-ray emission analysis was performed on several grains of poorly crystallized CuS. The results are plotted as a histogram in Fig. 2. There is a wide range in the composition of individual particles, but it is apparent that two CuS ratios are dominant (1.11 and 1.32). The ratios are in good agreement with those obtained by GOBLE and SMITH ( 1973 ) for blaubleibender covellite ( 1.12 ___0.01 and 1.32 ___0.04). GOBLE (1980) has subsequently reported blaubliebender coveUite to consist of two new minerals, yarrowite (Cu 1.125S)and spionkopite (Cu j.39S). He suggests that the CuS value of 1.39 for spionkopite is the ideal value and the value of 1.32 commonly reported may represent a mixture of these two new phases. All solutions were made with distilled demineralized water from a MiUipore Corp. Milli-Q system. All reagents were of analytical grade unless otherwise specified. Ionic strength was adjusted with NaCI or NaCIO4. Solutions of EDTA were preparedfrom the dried disodium salt. Solutions of DCTA were prepared in a manner similar to that of MAROERUM and BYDALEK(1963). The tetra acid form of DCTA
Bisulfide stock solutions were prepared by bubbling high purity H2S gas through a deoxygenated standard NaOH solution sealed in a septum vial. Polysulfide solutions were made by dissolving solid elemental sulfur (S °) into an HS- solution to obtain the desired concentration of dissolved zero-valent sulfur S(0). All pH measuremerits were made inside a N2-filled glove box with a Ross pH electrode (Orion Research, Inc.) and an Orion Research Model 701 Digital pH meter. The pH electrode was calibrated with either a National Bureau of Standards TRIS reference buffer (SRM 921,923) or with Scientific Products, pH Standard Reference Buffers. The third and fourth ionization constants of H4DCTA were determined by titrating 1.041 mM H2DCTA with 0.04473 M NaOH. Three replicate titrations were performed at ionic strengths of 0.20, 0.70 and 1.0 M (adjusted with NaCI). These were the same ionic strengths used in the solubility experiments below. The solution was purged with N2 gas to prevent CO2 contamination in the alkaline region of the titration curve. The Cu (II) -EDTA and Cu (I1)-DCTA stability constants were determined by the titration of a 0.500 mM CuCI2 solution with a 0.500 mM solution of the respective ligands. The progress of the titration was monitored with the Cu (II) specific electrode. The pH was maintained at 5.00 with an 0.02 M acetate buffer and the stoichiomctric ionic strength was adjusted to 0.20, 0.70 and 1.0 M with NaCI. Solubility experiments were performed by adding 0.10 g of the appropriate CuS to a 125 mL septum vial. These vials were transferred to a N2-filled glove box, where 100 mL of either an EDTA or DCTA solution ( N2 purged) was added. The pH was buffered with TRIS ( 1 raM). All vials were then sealed and placed in a shaker bath inside the glove box. The temperature of the bath was maintained at 25 _+0.2oc. At predetermined times, 20 mL samples were taken with a syringe and filtered through 0.45 #m Millipore filters to remove any suspended CuS. The pH was checked, and then sulfide analyses were performed using a solid-state AglAgaS electrode inside the glove box. A calibration curve was constructed using standard HS- solutions buffered by TRIS and similar in composition to the test solutions. In many cases, sulfide analyses were also performed amperomctrically by oxidizing all the sulfide with an excess of a standardized 12 solution. The residual 12was then back titrated with phenylarsine oxide. Plasma emission spectrometry at 324.75 nm was used to determine the concentration of total dissolved Cu. All electron diffraction and elemental microprobe analyses were performed on a JEOL JEM 200 CX electron microprobe operating in the STEM mode and equipped with an energy dispersive X-ray 1o
G) hi / ID
g-
k/Y
"7-
8-
6-
h 0
54-
¢r Iii m
Z o
,
.
o . g m'
.
.
.
.
.
.
.
.
.
1.05
Cu/S
j
.
.
.
.
.
1. ~ 5
(molar
,
.
.
.
.
1.25
.
.
.
.
.
.
.
1.3~
ratio)
FIG. 2. Microprobe analyses of poorly crystalline CuS. CuS values were obtained using covetlite as a standard (Cu:S = 1.00).
Solubility product of covellite detector. Electrondiffractionlmttemswere obtained at a beam voltage of 200 kV. Copper and sulfur determinations were made using the Cu Kot and S Ka emimionsat 15 kV. CoveUitewas used as a standard. Background corrections were made for atomic number, absorbance and fluorescence.
231
Table 1.
Proton and DCTA. a
Cu(II)
pE.e3
Medium
Stability
PKa4
Constants
for
PKCuDCTA'
RESULTS
0.20 thtCl
6.10+.01 10.40+_.02 22.14_-4-0.09
The method of BJERRUM (1941) Was used to determine the third and fourth ionization constants of D C r A by titration with NaOH. This method was used previously in determining the values ofpKa3 (6.18 _+0.03) and pK,4 (10.06 + 0.01) for EDTA (UHLER and HELZ, 1984). Plots of ti (the average protonation number of the ligand) vs. pH are shown in Figs. 3a and 3b, respectively. The pK, values obtained from these plots are listed in Table 1. Our values are in good agreement with those in the literature (ANDEREGG, 1967; M_ARTELLand SMITH, 1982). The Cu-EDTA and Cu-DCTA dissociation constants were determined from the following relationship
0.70 Nat1
5.88+.01 9.55+.02 20.71+--0.11
1.0 he1
5.84±.02
(Cu2+)(H2y -2) Kc,v, = (CuY_2)(H+) 2
(5)
where Kc~v, is the stability constant of the Cu-ligand species at the specified conditions ( I - 0.20, 0.70 and 1.0 M; T = 298 K; pH = 5.00 buffered with 0.02 M acetate). This was done by titrating a H2Y -2 solution with a Cu +2 titrant, while monitoring free copper (i.e. Cu not complexed by EDTA or DCTA) with a Cu ion selective electrode. The concentration of the free ligand, H2Y-2, can be calculated from the total
pH
,o.5
10,O
9"50. 0
0.2
0.¢
ole
0.8
9.35+.03 20.28_~.10
a Kz 3 [H+]IW)CT&-3I/IH2DCTA-2] Ka4 [e+][De~A-4I/[nDerA-3I KCuDCTA= [Cu+2][DCTA-4]/[CuDCTA -2] .
Errors represent I std. 3 replicate experiments.
deviation
from
the
mean of
free ligand concentration and the acidity constants given above. Twelve discrete values of log Kc~v, were calculated from three separate titrations for both Cu-EDTA and CuI)C~A. In the case of Cu-DCTA, titrations were performed at NaCI concentrations of 0.20, 0.70 and 1.0 M to give pKc,ocrA- values of 5.37 _+0.04, 5.04 +_0.06 and 4.87 _+0.05, respectively. These values apply to a medium containing the specified Nat1 concentration plus 0.02 M acetate at pH 5.0. In order to use them with the other constants, which were measured in acetate-free NaCI solutions, it is necessary to make a correction for CuAc + and CuAc ° complexing. Correction factors (i.e. 1 - ZacuAc,) were calculated from copper acetate association constants in MARTELLand SMITH(1982). The corrections increase pKc,v, by 0.27, 0.24 and 0.22 at I = 0.2, 0.7 and 1.0 respectively. The uncertainty in these corrections is about +0.05, based on the uncertainty in the association constants. The Cu-EDTA titrations were done at an ionic strength of 0.20 M only and an average value of pKCuEOTA"= 2.14 + 0.02 Was obtained. The error represents one standard deviation from the mean of 12 values. Coneeted to acetate-free medium, this becomes pKcuEDTA"= 2.41. Metal-ligand stability constants are usually expressed with respect to the deprotonated ligand, and for purposes of comparison with the values in the literature, we have computed such constants as follows
7.5-
pKcuv = pKc,v,, + PKt3 + p K g .
(6)
7.O -
6.S-
pH
8.o-
$.5 5.0 4,5 4,O l .o
,
,
,
,
i .=
1.4
i .s
i .s
2,O
FIG. 3. Determination of acidity constants for DCTA at 298 K. The average protonation number, ~ = 2-(Co.- - [OH-])/Yt, w h e r e C O I l - is the concentration of NaOH added by titration, OH is the free hydroxyl concentration and Yt is the total DCTA concentration. I = 0.20 M. a) pK.4 for HI~'-'TA-a is equal to the pH value at ~ = 0.5. b) pK.3 for H2DCTA-2 is equal to the pH value at ri = 1.5.
Values of pKcuv in acetate-free Nat1 solutions are given in Table 1 for Y = DCTA. A value ofpKc,e = 18.65 :t: 0.07 at I = 0.20 is obtained for Cu-EDTA. This agrees very well with MARTELL and SMITH'S(1982) value of 18.70 at 25"C and I --0.1. The solubility of eovellite was measured at various H S and HDCTA-3 concentrations at pH 8.1 and I = 0.20, 0.70 and 1.0 M. The solubility data are listed in Table 2 along with the calculated values of log Koiu and log K~ (defined by reactions 1 and 4, respectively). The reversibility of the solubilities was demonstrated by first allowing an HS- solution to equilibrate with covellite and then doubling the HS- concentration. The Cu concentration was observed first to rise to a plateau, and then after adding the HS-, to fall by a factor of 2 to a new steady state (SHEA, 1985).
232
D. Shea and G. R. Heiz T a b l e 2.
(I)
- l o g M CuT
Covelllta
- l o g M HS-
• 20
6.45+.04
5.50
Solubillty
D a ta .
- l o g M HDCTA- 3
2.30
l o g Kdlss
l o g Ksp
-9.65+.03
-21.39+0.13
.70
6.52~.02
5.58
2.30
-9.80~.02
-20.96+--0.14
.70
6.43~.04
5.54
2.00
-9.97~.04
-21.13+0.16
.70
5.67+.03
5.50
1.30
- 9 . 8 7 -+- . 0 4
- 2 1 . 0 3m +0.16
1.0
6.62+.03
5.64
2.30
--9.96+.04
--20.89+0.16
1.0
6.47+.02
5.52
2.00
--9.99+.03
--20.92+0.15
1.0
5.93+.03
5.45
1.30
- - 1 0 . 0 8 +. 0 3
--21.01+0o15
1.0
6.24+.01
5.08
1.30
--I0.02+.02
--20.95+0.14
1.0
6.81+.03
4.55
1.30
- - 1 0 . 0 6 +. 0 4
--20.99+0.16
4.05
t.30
--
1.0
<7.20
T - 298 K, pH = 8.1 (TRIS b u f f e r e d ) , i o n i c s t r e n g t h Errors represent
( I ) a d j u s t e d w i t h NaCI.
t h e s t a n d a r d d e v i a t i o n of 3 r e p l i c a t e
Kdis8 = [CUT] [HS-]/[HDCTA-3],
Ksp
--
experiments.
= Kdlss KCuDCTA" Ka3
[HS-] d e t e r m i n e d by a a p e r o m e t r i c t i t r a t i o n .
In one test, covellite was leached with DCTA to dissolve 25% of the original mass. Then K.p was measured on the leached solids. The average K,p in this test was within the uncertainty of the K,p determined on unleached covellite. Leaching should have eliminated small particles, strained crystals, sharp corners and other features which might lead to artificially high solubilities. Solubility measurements for poorly crystalline CuS are presented in Table 3. Despite the fact that this material is not a pure phase, it is possible to describe the data with the following equilibrium: x - 1 S(s) + 1 CuxS(s) + HEDTA3-(aq) ~X
X
CuEDTA2-(aq) + H S - ( a q )
Kd~.
3 that reproducible K,p values were obtained for poorly crystalline products prepared independently, at different times. The range of values is actually quite small considering the lack of crystallinity of these samples. The solubility of the poorly crystalline CuS was measured in an EDTA solution as a function of time to determine if reversible equilibrium was achieved. The data are plotted in Fig. 4. It appears that the poorly crystallized CuS reaches metastable equilibrium from both undersaturated and oversaturated solutions within about 6 hours. The poorly crystalline CuS was stable for at least 3 weeks in a 0.20 M NaC1 solution containing 10 mM HS-. However, it recrystallized within about 4 days when 1 mM S(0) was added to this solution in the form of polysulfide ions. The
(7) T a b le 3.
When this is combined with reactions (2) and ( 3 ) we get: H+ + x - 1 S(s) + 1 CuxS(s) ~ Cu2+(aq) + H S - ( a q ) X
S o l u b i l i t y Data f o r Poorly C r y s t a l l i n e CuS.
Sample
- l o g CUT
- l o g KS-
- l o g EDTA
l o g F,Sp
A-1
6.62+.04
5.08
2.30
-17.99
A-1
6.93+. 05
5.12
2.60
-18.04
X
(8) (Cu2+)(HS - ) K,o = (H+)(aso)( x _ l)/x"
(9)
Note that the equilibrium constant expression for (8) is identical in form to that for (4) except for the activity of sulfur term in the denominator. Since the poorly crystalline material contained elemental sulfur, we have assumed that the activity of sulfur is unity. However, this assumption may be a source of systematic error, as we have not proved that the sulfur which was present was in the stable, rhombic form. In any case, the calculated value ofK, p is insensitive to the assumed activity of sulfur because the exponent in Eqn. (9) is small. Values of log K,p are given in Table 3 for each solubility experiment. The mean value, -18.02, is 3 orders of magnitude greater than Ksp for covellite. It should be noted in Table
4.96
3.00
--
A-I
7.12+.09
4.50
2.30
-17.91
A-2
6.75+.06
5.04
2.30
-18.08
A-I
<7.40
A-3
6.58+.03
5.09
2.30
-17.96
A-4
6.82+.06
5.02
2.30
-18.12
A-5
6.70+.07
5.04
2.30
-18.03
T - 298 K, pH - 8 . 1 (TRIS b u f f e r e d ) ,
I = 0 . 2 0 ( a d j u s t e d wlch BaCl),
0 . 1 0 g CuS i n 100 aL s o l u t i o n . E r r o r s a r e s t a n d a r d d e v i a t i o n of 3 r e p l i c a t e S a a p l e s A - l , A-2, A-3, A-4, and A-5 w e r e to t e s t
reproducibility
of th e s y n t h e s i s .
experiments.
made at
different
times.
Solubility product of covellite 55 6.0
Mmm~m"m~
6.5
•
•
o ~ z5 !
8.0 8.5
233
conditions were probably appropriate i'or recrystallization of Rickard's initial precipitates to more stable forms during storage in solution. Above pH 8.5 Rickard observed blaubleibender covellite, rather than normal covellite. For given total polysulfide and bisulfide concentrations, the activity of S(0) will be lower at higher pH. This may account for the failure of sulfur-deficient blaubleibender covellite to recrystallize to covellite in Rickard's more alkaline runs. Free energy o f covellite
0
d5
IiO TIME
'15
2.0
(days)
FiG. 4. ]_08 concentration of Cu as a function of time for poorly crystalline CuS. T = 298 K, ! = 0.2 M. Squares represent the solubitity in a solution containing [ H S - ] = 0.01 raM, [ E D T A ] = 5.0 raM. Triangles represent measurements where the above solution was spiked with 0.09 m M HS- after 8 hours, thus demonstrating reversibility.
It is possible to derive a value for the free energy of formarion ofcovellite from our solubility measurements by using published values for the free energy of formation of Cu 2+ and H S - (from Roam et al., 1978). However, it is first necessary to correct our K~ values to zero ionic strength. Assuming no complexing between Cu 2+ and NaCI, the relationship between K,p in an ionic medium and its corresponding value at infinite dilution (i,e., K°~) would be:
~p "Yc.,÷~.s-K,~
(10)
K,p = Mc,:+MHs-(IO+pH).
(1 i)
=
X-ray diffraction pattern of the solid was taken after storage in the S(0) solution for 14 days. This pattern was similar to that of the covellite, indicating that the poorly crystalline CuS inverts to less soluble covellite in the presence of S(0). The role of S(0) is probably to produce polysulfide ligands by reacting with HS-. Polysulfides f o r m strong complexes with Cu (SHEA and HELZ, 1988), thus increasing solubility and facilitating recrystallization. The thermal stability of poorly crystalline CuS was investigated by placing a bulk sample in an evacuated pyrex vessel and heating to 250°C for 30 days. The X-ray diffraction pattern of the annealed product indicated that it was crystalline coveUite. Although no extraneous peaks were observed, an additional, Cu-rich phase must also have been present to preserve mass balance. The 3 mol% S Opresent in the poorly crystalline product is insufficient to convert it all to covellite.
where
Using our value of Kw at I = 0.2 M and activity coefficients calculated from the Davies equation, we obtain pK~,p= 22.06 and a free energy of formation of covellite o f - 1 1 . 5 kcal/ mole ( - 4 8 . 3 kJ/mole). In Fig. 5, the curve corresponding to Eqn. (10) is shown as a dashed line. This curve fails to pass through the data at I = 0.7 and 1.0. One way of improving the fit to the data is to hypothesize that copper forms complexes with e l - in the NaC1 medium. Assuming the existence of only one such complex, CuCI+, the following modified form of Eqn. (10) can be derived:
DISCUSSION RICKARD (1972) also haS studied precipitates formed by mixing CuSO4 and NariS solutions. In contrast to our results, Rickard failed to observe the poorly crystalline phase, obtaining instead either covellite or blaubleibender covellite, depending upon pH. Rickard's procedure differed from ours in that he used a 30 to 1 molar excess of NariS, whereas we used a stoichiometric, 1 to 1 ratio. Rickard also stored his precipitates in solution for 7 days before X-ray analysis, whereas we vacuum dried our precipitates within an hour of formation. The difference between Rickard's results and ours can be accounted for as follows. The precipitation of CuxS ( x > 1) from a solution containing cupric and bisulfide ions generates S(0) and acidity by the reverse of reaction 8. In our experiments, the final pH would be near one, whereas in Rickard's it would be near its original, neutral to alkaline value, owing to his excess HS-. Under these conditions, S(0) precipitated as solid sulfur in our experiments and was observed in the product. However, it would have remained in solution as polysulfide ions in Rickard's experiments. Since polysulfide ions solubilize copper minerals by strongly complexing Cu,
20
.--21 i 0
MATHIEU I~ I e~ R'CKERTT (1972)/ff " ~ ,
0 J
-D~VTE~- . . . .
et oL(1978)
23 • WAGMAN el" al. (1982) O POTTER ' (1977)
• PANKRATZ, et ol. (1987)
24
6
o12
' o14
' o16
' ale
'
FIG. 5. Extrapolation of K,o to infinite dilution and comparison with published data. Dashed line indicates extrapolation according to equation 10 and solid line indicates extrapolation according to equation 12.
234
D. Shea and G. R. Helz
K°o =
Table 4.
7c.2"7Hs-K~p I +/~c.cl+Mo-Vc.2+
'
(12)
Where K°cucl÷ = acuct-/acu2+act-. We have treated K°,o+ as a variable parameter and chosen its value so as to make a plot of Eqn. (12) exactly fit the data at I = 0.2 and I = 1.0, as shown by the solid line in Fig. 5. Activity coefficients were again evaluated from the Davies equation. The best value of /~cucl. was l 1.1. Independent measurements of this constant tend to be lower than this value (OHLSONand VANNERBURG, 1974), but of similar magnitude. Based on the solid line in Fig. 5, the preferred value of pK~p is 22.27 _-+ 0.3 and the corresponding free energy of formation of covellite is -11.83 -+ 0.4 kcal/mole (-49.50 _+ 1.7 kJ/mole). The uncertainties in the pK~o measurements in NaC1 media are approximately -+0.15, but the uncertainty in pK~o is greater owing to the uncertainty in the extrapolation to I = 0. The uncertainty of +0.3, assigned above, represents our estimate of the likely uncertainty inherent in this extrapolation. The free energy of formation of covellite, in kcal/ mole, is reported to be -11.7 _+ 1.0 by ROBIE et al. (1978), -11.58 by MATHIEU and RICKERT (1972), but -12.89 + 0.05 by POTTER (1977), --12.8 by WAGMANet aL (1982) and -13.44 by PANKRATZ et al. (1987). Our best estimate of AG ° is in excellent agreement with the first two values but clearly can not be reconciled with the latter values (see Fig. 5). Especially troubling is our disagreement with Potter's value, which is a high precision result based on electrochemical measurements. Possibly Potter failed to determine the potential of the covellite-rhombic sulfur assemblage, as he believed, and instead obtained the potential o f a covellite-metastable sulfur assemblage. Amorphous forms of sulfur are well known. The discrepancy between our AG ° value and Potter's is of the right magnitude for a crystalline to amorphous transition. Potter used a sulfide-free, CuSO4 aqueous electrolyte in his cell. This electrolyte is a very poor solvent for elemental sulfur. If amorphous sulfur formed in such a cell, it would not be likely to recrystallize to a stable form. (See also the discussion by NOWAK et aL, 1984, concerning additional potential problems associated with use of CuSO4 electrolyte.) In a paper on the solubility product of galena (UHLERand HELZ, 1984), no correction for Pb-chloride complexing was made in the extrapolation to infinite dilution. When this correction is made with the complex stability constants of SEWARD(1984), a pK°p of 12.78 is obtained. This is lower than the previously recommended value of 12.25, but it is still substantially larger than that obtained by extrapolation of high temperature free energy data, i.e. 14.68 according to ROBIE et aL (1978).
Comparison with FeS Table 4 compares pK~,pvalues ofcovellite and poorly crystalline CuS with analogous phases in the Fe-S system. Also shown are estimated pK~tp values for supercooled CuS and FeS liquids. These latter values, which have been estimated by extrapolation of high temperature thermodynamic data, probably indicate the solubilities of the least stable condensed phases that could possibly precipitate from aqueous solution.
Values of pKsp for CuS and PeS Phases at 2980K. o
pKsp
Phase
Data Source
CuS
Covellite
(CuS)
Poorly Cryscallne Supercooled
CuI.18S
Liquid CuS
22.27
~his study
18.90
This study
15.63
Larraln, et al. 1979
ZeS Troillte
(FeS)
Poorly Crystalline Supercooled
Fel.13 S
Liquid ~eS
6.06
Robie, et el. 1978
2.94
Berner 1964, 1967
0.51
Robie, et al. 1978
Even though the iron phases are uniformly much more soluble than the corresponding copper phases, there are some striking analogies between the two sets of daia. In both cases, the stable crystalline phase is about 6 orders of magnitude less soluble than the amorphous, supercooled liquid, and the poorly crystalline phase is roughly midway between the supercooled liquid and the crystalline phase in solubility. Furthermore, in both cases, the poorly crystalline phase possesses a metal to Sulfur ratio in excess of unity. Surprisingly, these analogies exist despite differences in the crystal structures of covellite and troilitc and despite differences in the common valencies of the two metals.
Phase diagram Using the free energy of formation of covellite together with free energies for other stable Cu-S phases permits construction of a phase diagram, as shown in Fig. 6. The free energy values used were - 17.71 kcal/ mole for anilite, - 19.08 kcal/mole for djurleite and -19.39 kcal/mole for low chalcocite. These values were obtained by correcting POTTER'S (1977) data to be consistent with our free energy for coveUite. These values agree reasonably with those of MATHIEU and RICKERT ( 1972 ). Figure 6 illustrates several important points about the thermodynamics of copper in sulfidic waters. First, it is clear that the solubility of copper in natural sulfidic waters can not be defined without knowledge of the activity of S o. Unfortunately, in very few field studies has an effort been made to quantify this variable. SHEAand HELZ (1988) discuss methods for doing this. Second, two intermediate phases, anilite and djurleite, are stable at temperatures found near the Earth's surface. In most previous efforts to model the behavior of copper in sulfidic waters, saturation only with respect to chalcocite and covellite has been considered. Finally, covellite is the stable phase at high activities of S °, where polysulfide aqueous species have maximum stability. In an effort to account for the solubility of copper in sulfidic waters, some workers have explored equilibria between chaicocite and dissolved copper polysulfide complexes (e.g., BOULEGUE et al., 1982; JACOBS et al., 1985). Such equilibria are metastable. Chalcocite will be converted to a more sulfur-rich solid phase by taking up S(0) from polysulfides until either the chalcocite disappears or the polysullides are depleted to immeasurably low concentrations.
Solubility product of covellite
pooR~ x'rA_.L-~
18
19
-
LOW CHALCOCITE AI
LITE
(
.-.~S
tJ COVELLITE
Cu,,S
•, 2 0 Z
CuS
0.7 M NelCl
// /
22
/
/
' 23
24
t-O
,J JNSATURATED SOLUTION
J i - log
o
as*
FIG. 6. Phase diagram for CuS system at 25°C. Contours indicate solubility limits at various concentrations of NaCI. The activity of sulfur, on the abscissa,is definedwith respect to stablerhombic sulfur. Also shown in Fig. 6 is the solubility of our poorly crystalline material. This has been plotted as a dotted line with a slope consistent with the average composition of Cu,.,sS and an intersection at log as0 = 0 determined by the measured K,p. This is not a rigorous procedure, as the poorly crystalline material is not a pure phase and its composition can be expected to change with as 0. However, the position of the dotted line serves to place this material in the context of the stable mineral phases in the Cu-S system. Possibility o f a poorly crystalline C u S phase in nature In view of the probable role of poorly crystalline FeS in controlling iron solubility in natural waters (DAVISON, 1980), it is of interest to consider whether poorly crystalline CuS may play a similar role. BROCKAMP el al. (1978) investigated the chemical form of Cu and Zn in Red Sea Brine deposits. Some of these deposits contained no X-ray detectable chalcopyrite or sphalerite and yet contained up to several percent Cu and Zn. Using a selective extraction procedure that they claim isolates the sulfide-bound fraction of these metals in the deposits, BROCKAMPet al. (1978) showed that an appreciable fraction of both metals appears to be sulfide-bound. Lacking X-ray diffraction evidence for the existence of crystalline sulfides, they concluded that the Cu and Zn are contained in amorphous sulfide phases. If the poorly crystalline copper sulfide that we have studied is to be found anywhere in nature, one would expect to find it where solutions of markedly different composition and temperature mix, as at the top of the Red Sea brine pools. Aside from rapid mixing of disparate solutions, it is difficult to envision physical processes in nature that could produce the high degree of supersaturation required to precipitate the poorly crystalline material. Nevertheless, the thermal instability of this material and its rapid conversion to covellite when exposed to solutions containing S(0) in the form of polysulfides, suggests that this material is unlikely to be a persistent phase in nature. To
235
find it would at the very least require particular precautions during sampling to prevent production of S (0) by oxidation. No such precautions were reported by BROCKAMP et al. (1978). On the basis of these considerations, we doubt that the poorly crystalline material we studied is the X-ray amorphous material inferred to exist in the Red Sea Brine deposits. Because of its instability, poorly crystalline CuS would be expected to have only an ephemeral existence in nature and would not be expected to accumulate in sediments. However, the possibility remains that the poorly crystalline phase could control copper solubilities for short periods in highly dynamic environments. If indeed a non-crystalline form of sulfide-bound copper exists in the Red Sea deposits, then we suspect it is a Cu-Fe sulfide phase. Phases in the Cu-Fe-S system are more sluggish in equilibrating than phases in the Cu-S system (BARTON and SKINNER, 1979). Thus the chances are better that an amorphous Cu-Fe sulfide would be persistent enough to be observed. Acknowledgements--Acknowledgement is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society for financial support. The National Bureau of Standards provided access to the electron microscope, and this work was assisted by Patrick Sheridan. Editorial handling: S. A. Wood REFERENCES ANDEREGGG. (1967) Komplexone, L.X. Die Protonierungskonstanten einiger Komplexone in verschiedenen wasserigen Salzmedien. Heir. Chim. Acta 50, 2333-2340. BARTONP. B. JR. and SKJNNERB. J. (1979) Sulfidemineralstabilities. In Geochemistry of Hydrothermal Ore Deposits (ed. H. L. BARNES), 2nd edn., pp. 278-403. L Wiley & Sons. BERNERR. A. (1964) Iron sulfidesformed from aqueous solution at low temperatures and atmospheric pressure. J. Geol. 72, 293-306. BERNERR. A. (1967) Thermodynamic stability of sedimentary iron sulfides. Amer. J. Sci. 265, 773-785. BJERRUMJ. ( 194l) Metal Amine Formation in Aqueous Solution. P. Haase & Sons, Copenhagen, 208p. BOULEGUEJ., LORDC. J. and CHURCHT. M. (1982) Sulfurspeciation and associated trace metals (Fe, Cu) in the pore waters of Great Marsh, Delaware. Geochim. Cosmochim. Acta 46, 453--464. BROCKAMP O., GOULART E., HARDER H. and HEYDEMANN A.
(1978) Amorphous copper and zinc sulfides in the metalliferous sediments of the Red Sea. Contrib. Mineral Petrol. 68, 85-88. CARtGNANR. and NRIAGUJ. O. (1985 ) Trace metal deposition and mobility in the sediments of two lakes near Sudbory, Ontario. Geochim. Cosmochim. Acta 49, 1753-1764. DAVlSONW. (1980) A criticalcomparisonof the measuredsolubilities of ferrous sulfide in natural waters. Geochim. Cosmochim. Acta 44, 803-808. DYRSSEND. (1985) Metal complex formation in sulfidic seawater. Mar. Chem. 15, 285-293. EMERSONS., JACOBSL. and TREBOB. (1983) The behavior of trace metals in marine anoxicwaters: Solubilitiesat the oxygen-hydrogen sulfide interface. In Trace Metals in Sea Water (eds. C. WONGet al. ), pp. 579-608. Plenum Publishing. GAILLARDJ.-F., JEANDEL C., MICHARD G., NICHOLAS E. and RENARD D. (1986) Interstitial water chemistry of Villefranche Bay sediments: Trace metal diagenesis.Mar. Chem. 18, 233-247. GOBLER. J. (1980) Copper sulfides from Alberta: Yarrowite CugSs and Spionkopite Cu39S2s. Can. Mineral. lg, 511-518. GOBLER. J. and SMITHD. E. W. (1973) Electron microprobe investigation of copper sulfides in the Precambrian Lewis Series of S.W. Alberta, Canada. Can. Mineral. 12, 95-103.
236
D. Shea and G. R. Helz
JACOBSL. and EMERSONS. (1982) Trace metal solubility in an anoxic fiord. Earth Planet. Sci. Lett. 60, 237-252. JACOBSL., EMERSONS. and SKEIJ. (1985) Partitioning and transport of metals across the O2/H2S interface in a permanently anoxic basin: Framvaren Fjord, Norway. Geochim. Cosmochim. Acta 49, 1433-1444. KltEMLING K. (1983) The behavior ofZn, Cd, Cu, Ni, Co, Fe and Mn in anoxic Baltic waters. Mar. Chem. 13, 87-108. LARRAINJ. M., LEE S. L. and KELLOGGH. H. (1979) Thermodynamic properties of copper-sulfur melts. Can. Metallurg. Quart. 8, 395-400. MARGERUMD. W. and BYDALEKT. J. (1963) Multidentate ligand kinetics. V. Copper (1I) and cyciohexylenediamine tetraacetonickelate (II). Inorg. Chem. 2, 683-688. MARTELLA. E. and SMITHR. M. (1982 ) Critical Stability Constants. Vol. 5, First supplement. Plenum Press, New York. MATHIEU H. J. and RICKERT H. (1972) Eiectrochemischthermodynamische Untersuchungen am System Cu-S. Z. Phys. Chem. Neue Folge 79, 315-330. NOWAK P., BARZYK W. and POMIANOWSKIA. (1984) The applicability of EMF measurements to evaluation of thermodynamic properties of the Cu-S system. J. Electroanal. Chem. Interfacial Electrochem. 171, 355-358. OHLSON M. and VANNERaERGN.-G. (1974) Stability constants for chloride complexes of copper (lI) in sulfuric acid solution, Acta Chem. Scand. A28, 1031-1035. PANKRATZL. B., MAH A. D. and WATSONS. W. (1987) Thermodynamic properties of sulfides. U.S. Bur. Mines. Bull. 689, 427p.
POTTER R. W. (1977) An electrochemical investigation of the system copper-sulfur. Econ. Geol. 72, 1524-1542. POTTER R. W. and EVANS H. T. (1976) Definitive X-ray powder data for covellite, anilite, djudeite and chalcocite. J. Res. U.S. Geol. Surv. 4, 205-212. PUTNIS A., GRACEJ. and CAMERONW. E. (1977) Blaubleibender covellite and its relationship to normal covellite. Contrib. Mineral. Petrol. 60, 207-217. RICKARD D. Z. (1972) Covellite formation in low temperature aqueous solutions. Mineral. Deposita 7, 180-188. ROaIE R. A., HEM]NGWAVR. S. and FISHERJ. R. (1978) Thermodynamic properties of minerals and related substances at 298.15 K and l bar pressure and at higher temperatures. U.S. Geol. Surv. Bull. 1452, 456p. SHEAD. ( 1985 ) Solubility and precipitation of copper sulfide. Ph.D. dissertation, Univ. of Maryland, 250p. SHEA D. and HELZ G. R. (1988) The solubility of copper in sulfidic waters: Sulfide and polysulfide complexes in equilibrium with covellite. Geochim. Cosmochim. Acta 52, 1815-1825. SEWARDT. M. ( 1984 ) The formation of lead (II) chloride complexes to 300oC: A spectrophotometric study. Geochim. Cosmochim. Acta 48, 121-134. UHLER A. D. and HELZ G. R. (1984) Solubility product of galena at 298 K. Geochim. Cosmochim. Acta 48, 1155-1160. WAGMAN D. D., EVANS W. H., PARKER V. B., SCHUMM R. H., HALOW 1., BAILEYS. M., CHURNEY K. L. and NUTTALLR. L. ( 1982 ) The NBS tables of chemical thermodynamic properties. J. Phys. Chem. Ref. Data. I I: Supplement 2, 392p.
0016-7037/89/$3.00 + .00
Solubility product constants of covellite and a poorly crystalline copper sulfide precipitate at 298 K* DAM1AN SHEAt and GEORGE R. HELZ Department of Chemistry and Biochemistry, University of Maryland, College Park, MD 20742, U.S.A. ( Received Apri120, 1988; accepted in revised form August 8, 1988) Abstract--The equilibrium constant at 25"C for the following reaction has been measured in NaC1 media by an indirect method: CuS(COv) + H+(aq) ~ Cu2+(aq) + HS-(aq),
K~ = Mcu2.M,s-(10 +p")
where CuS(cov) designates synthetic covellite. Values ofpK~ are 21.39, 21.04 and 20.95 at NaCI = 0.2, 0.7 and 1.0 M, respectively; the uncertainty in these K~ values is _+0.15. The free energy of formation of covellite, for which published values are discordant, is calculated to be -11.83 _+0.4 kcal/mole at 298 K (-49.50 + 1.7 kJ/mole). This value is obtained by extrapolating the measured pK,p values to infinite dilution with corrections for CI- complexing. Applying similar CI- complexing corrections, based on recent measurements by Seward, to preyiously published solubility data for galena yields a revised pK°wfor galena of 12.78. A poorly crystalline precipitate, obtained by mixing Cu 2÷ and HS- solutions, yielded a reversible solubility product 3 orders of magnitude greater than that of covellite but about 3 orders of magnitude less than that of a truly amorphous phase, super-cooled liquid CuS. The poorly crystalline phase has not been studied previously. Its bulk composition was Cu ,.,aS, but microprobe analysis revealed that it was a partially exsolved mixture of roughly Cu L, ,S and Cu t.32S (similar to known blaubleibender covellites). It was kinetically unstable, and converted to coveUitewhen thermally annealed or when exposed to polysulfide solutions. Because of its instability, a material of this nature is unlikely to account for the amorphous copper sulfide alleged to occur in the Red Sea Brine deposits. However, it is possible that on short time scales dissolved Cu in sulfidic waters is controlled by metastahle, rather than stable phases, as is known to be the case with dissolved Fe. INTRODUCTION
system. We have measured the solubility product ofcovellite and of a poorly crystalline, metastable phase formed by rapid mixing of CuCI2 and NariS solutions. Combining these measurements with data from the literature, we have estimated the solubility of some additional phases. To determine solubility products, we employed the indirect method of UHLER and HELZ (1984). This involves using a chelating agent, Y, to enhance the solubility of a sulfide mineral sufficiently to enable convenient measurement. Equilibrium constants for the following three reactions are determined and then multiplied to get the solubility product constant, K,p.
M U C H C U R R E N T A T T E N T I O N i s being devoted to understanding the geochemical behavior of trace metals in low temperature, sultidic environments (BOULEGUEet al., 1982; JACOBS and EMERSON, 1982; EMERSON et al., 1983; KREMLING, 1983; DYRSSEN, 1985; CARIGNAN and NRIAGU, 1985; JACOBS et aL, 1985; GAmLARD et al., 1986). A m o n g the m a n y problems is the need to establish what phases control trace metal solubilities in these environments. Interfaces between deep-seated sulfidic waters and oxic surface waters are important sites of sulfide mineral deposition. Such interfaces occur in groundwaters around sulfide ore deposits, in pore waters of modern sediments, and in fjords or other types of anoxic marine basins. Mixing of waters across such interfaces can readily generate very high degrees of supersaturation with respect to trace metal sulfide minerals, permitting deposition ofmetastable phases. Indeed, metastable blaubleibender covellite phases, in which the C u / S ratio exceeds 1.0, are often found in environments of supergene enrichment. BROCKAMP el al. (1978) even report that an amorphous copper-bearing sulfide phase occurs in deposits from the Red Sea thermal brines. Thus it is conceivable that metastable Cu-S phases may control copper solubilities at temperatures occurring near the Earth's surface. In this regard, copper would be analogous to iron, which is known to be controlled by metastable phases (e.g. DAVISON, 1980). This paper concerns the solubility of phases in the Cu-S
CuS(s) + H Y 3- ~ CuY 2- + H S -
Kdiu
(1)
Kcuv*
(2)
H2Y 2 - ~:~ H + + H Y 3-
Ka3
(3)
CuS(s) + H + ~ Cu 2+ + H S -
K~
(4)
CuY 2- + 2H + ~ Cu 2+ + H2Y2-
In measurements on the poorly crystallized precipitate, we used the chelator, ethylenediaminetetraacetic acid ( E D T A ) . For covellite, which is much less soluble, we had to use the stronger chelator, trans 1-2 diaminoeyclohexyltetraacefic acid (DCTA). EXPERIMENTAL Materials
* This paper was part ofan invited presentation at the MERI conference. ' Present address: Department of Chemistry, Southeastern Massachusetts University, North Dartmouth, MA 02747, U.S.A.
Covellite was synthesized by mixing about 10 g of a 1:1 molar ratio of 99.99% pure Cu powder and 99.999% pure S powder in an evacuated pyrex vessel. This mixture was allowed to react for one day at room temperature and was then heated to 2500C for 30 days. 229
230
D. Shea and G. R. Helz (H4 DCTA) was twice recrystallized from water, dried and then standardized by potentiomctric titration against a standard Cu(II) solution.
Methods
30 °
40 °
50 °
60 °
2e FIG. 1. X-ray powder diffraction patterns for (a) covellite and (b) poorly crystalline CuS (Cu ka radiation ).
The product was washed and in some cases leached in a deoxygenated 0. l M DCTA solution in a N2-filled glove box. The X-ray diffraction pattern for this sample exhibited sharp covellite peaks (see Fig. 1). The peak positions and intensities were in excellent agreement with those of POTTER and EVANS (1976), indicating a well-crystallized material. Ten covellite samples of approximately 0. ! g (weighed to 0.1 mg) were dissolved in warm aqua regia and Cu analysis was performed using plasma emission spectrometry to determine the weight percent of Cu in these samples. The average atomic Cu/S ratio was found to be 1.008 __.0.013. The poorly crystalline CuS was made in an N2-filled glove box by mixing 500 mL of N2-purged 1.00 M CuCI2 (pH = 4) with 500 mL of N2-purged 1.00 M NaFIS (pH = 8-11). After 10 minutes of shaking, the solution was filtered and the dark blue residue was washed with distilled water and allowed to dry under vacuum in the antechamber of the glove box. The X-ray diffraction pattern of this product was obtained within a few hours and is also shown in Fig. 1. This pattern exhibited a very high background with broadened peaks of low intensity. In addition, the strong ( 108 ) and ( ! 16) covellite reflections, occurring at 55* and 61.5" 2# respectively, in Fig. 1 are absent in the poorly crystalline CuS. The absence of these two reflections is characteristic of blaubleibender covellite (PUTNIS et al., 1977). The single particle electron diffraction patterns of approximately 1 #m particles were also very poor, exhibiting diffuse rings about the electron beam. Copper analysis was performed on the poorly crystalline product after washing with 0.1 M Naris (pH = 12 ) to remove any elemental sulfur (S °) as polysulfides. An atomic Cu/S ratio of 1.189 + 0.015 was determined. Comparison of this value with the corresponding value determined before the Naris wash suggests that the poorly crystalline material contained about 3 tool% elemental sulfur. Single particle X-ray emission analysis was performed on several grains of poorly crystallized CuS. The results are plotted as a histogram in Fig. 2. There is a wide range in the composition of individual particles, but it is apparent that two CuS ratios are dominant (1.11 and 1.32). The ratios are in good agreement with those obtained by GOBLE and SMITH ( 1973 ) for blaubleibender covellite ( 1.12 ___0.01 and 1.32 ___0.04). GOBLE (1980) has subsequently reported blaubliebender coveUite to consist of two new minerals, yarrowite (Cu 1.125S)and spionkopite (Cu j.39S). He suggests that the CuS value of 1.39 for spionkopite is the ideal value and the value of 1.32 commonly reported may represent a mixture of these two new phases. All solutions were made with distilled demineralized water from a MiUipore Corp. Milli-Q system. All reagents were of analytical grade unless otherwise specified. Ionic strength was adjusted with NaCI or NaCIO4. Solutions of EDTA were preparedfrom the dried disodium salt. Solutions of DCTA were prepared in a manner similar to that of MAROERUM and BYDALEK(1963). The tetra acid form of DCTA
Bisulfide stock solutions were prepared by bubbling high purity H2S gas through a deoxygenated standard NaOH solution sealed in a septum vial. Polysulfide solutions were made by dissolving solid elemental sulfur (S °) into an HS- solution to obtain the desired concentration of dissolved zero-valent sulfur S(0). All pH measuremerits were made inside a N2-filled glove box with a Ross pH electrode (Orion Research, Inc.) and an Orion Research Model 701 Digital pH meter. The pH electrode was calibrated with either a National Bureau of Standards TRIS reference buffer (SRM 921,923) or with Scientific Products, pH Standard Reference Buffers. The third and fourth ionization constants of H4DCTA were determined by titrating 1.041 mM H2DCTA with 0.04473 M NaOH. Three replicate titrations were performed at ionic strengths of 0.20, 0.70 and 1.0 M (adjusted with NaCI). These were the same ionic strengths used in the solubility experiments below. The solution was purged with N2 gas to prevent CO2 contamination in the alkaline region of the titration curve. The Cu (II) -EDTA and Cu (I1)-DCTA stability constants were determined by the titration of a 0.500 mM CuCI2 solution with a 0.500 mM solution of the respective ligands. The progress of the titration was monitored with the Cu (II) specific electrode. The pH was maintained at 5.00 with an 0.02 M acetate buffer and the stoichiomctric ionic strength was adjusted to 0.20, 0.70 and 1.0 M with NaCI. Solubility experiments were performed by adding 0.10 g of the appropriate CuS to a 125 mL septum vial. These vials were transferred to a N2-filled glove box, where 100 mL of either an EDTA or DCTA solution ( N2 purged) was added. The pH was buffered with TRIS ( 1 raM). All vials were then sealed and placed in a shaker bath inside the glove box. The temperature of the bath was maintained at 25 _+0.2oc. At predetermined times, 20 mL samples were taken with a syringe and filtered through 0.45 #m Millipore filters to remove any suspended CuS. The pH was checked, and then sulfide analyses were performed using a solid-state AglAgaS electrode inside the glove box. A calibration curve was constructed using standard HS- solutions buffered by TRIS and similar in composition to the test solutions. In many cases, sulfide analyses were also performed amperomctrically by oxidizing all the sulfide with an excess of a standardized 12 solution. The residual 12was then back titrated with phenylarsine oxide. Plasma emission spectrometry at 324.75 nm was used to determine the concentration of total dissolved Cu. All electron diffraction and elemental microprobe analyses were performed on a JEOL JEM 200 CX electron microprobe operating in the STEM mode and equipped with an energy dispersive X-ray 1o
G) hi / ID
g-
k/Y
"7-
8-
6-
h 0
54-
¢r Iii m
Z o
,
.
o . g m'
.
.
.
.
.
.
.
.
.
1.05
Cu/S
j
.
.
.
.
.
1. ~ 5
(molar
,
.
.
.
.
1.25
.
.
.
.
.
.
.
1.3~
ratio)
FIG. 2. Microprobe analyses of poorly crystalline CuS. CuS values were obtained using covetlite as a standard (Cu:S = 1.00).
Solubility product of covellite detector. Electrondiffractionlmttemswere obtained at a beam voltage of 200 kV. Copper and sulfur determinations were made using the Cu Kot and S Ka emimionsat 15 kV. CoveUitewas used as a standard. Background corrections were made for atomic number, absorbance and fluorescence.
231
Table 1.
Proton and DCTA. a
Cu(II)
pE.e3
Medium
Stability
PKa4
Constants
for
PKCuDCTA'
RESULTS
0.20 thtCl
6.10+.01 10.40+_.02 22.14_-4-0.09
The method of BJERRUM (1941) Was used to determine the third and fourth ionization constants of D C r A by titration with NaOH. This method was used previously in determining the values ofpKa3 (6.18 _+0.03) and pK,4 (10.06 + 0.01) for EDTA (UHLER and HELZ, 1984). Plots of ti (the average protonation number of the ligand) vs. pH are shown in Figs. 3a and 3b, respectively. The pK, values obtained from these plots are listed in Table 1. Our values are in good agreement with those in the literature (ANDEREGG, 1967; M_ARTELLand SMITH, 1982). The Cu-EDTA and Cu-DCTA dissociation constants were determined from the following relationship
0.70 Nat1
5.88+.01 9.55+.02 20.71+--0.11
1.0 he1
5.84±.02
(Cu2+)(H2y -2) Kc,v, = (CuY_2)(H+) 2
(5)
where Kc~v, is the stability constant of the Cu-ligand species at the specified conditions ( I - 0.20, 0.70 and 1.0 M; T = 298 K; pH = 5.00 buffered with 0.02 M acetate). This was done by titrating a H2Y -2 solution with a Cu +2 titrant, while monitoring free copper (i.e. Cu not complexed by EDTA or DCTA) with a Cu ion selective electrode. The concentration of the free ligand, H2Y-2, can be calculated from the total
pH
,o.5
10,O
9"50. 0
0.2
0.¢
ole
0.8
9.35+.03 20.28_~.10
a Kz 3 [H+]IW)CT&-3I/IH2DCTA-2] Ka4 [e+][De~A-4I/[nDerA-3I KCuDCTA= [Cu+2][DCTA-4]/[CuDCTA -2] .
Errors represent I std. 3 replicate experiments.
deviation
from
the
mean of
free ligand concentration and the acidity constants given above. Twelve discrete values of log Kc~v, were calculated from three separate titrations for both Cu-EDTA and CuI)C~A. In the case of Cu-DCTA, titrations were performed at NaCI concentrations of 0.20, 0.70 and 1.0 M to give pKc,ocrA- values of 5.37 _+0.04, 5.04 +_0.06 and 4.87 _+0.05, respectively. These values apply to a medium containing the specified Nat1 concentration plus 0.02 M acetate at pH 5.0. In order to use them with the other constants, which were measured in acetate-free NaCI solutions, it is necessary to make a correction for CuAc + and CuAc ° complexing. Correction factors (i.e. 1 - ZacuAc,) were calculated from copper acetate association constants in MARTELLand SMITH(1982). The corrections increase pKc,v, by 0.27, 0.24 and 0.22 at I = 0.2, 0.7 and 1.0 respectively. The uncertainty in these corrections is about +0.05, based on the uncertainty in the association constants. The Cu-EDTA titrations were done at an ionic strength of 0.20 M only and an average value of pKCuEOTA"= 2.14 + 0.02 Was obtained. The error represents one standard deviation from the mean of 12 values. Coneeted to acetate-free medium, this becomes pKcuEDTA"= 2.41. Metal-ligand stability constants are usually expressed with respect to the deprotonated ligand, and for purposes of comparison with the values in the literature, we have computed such constants as follows
7.5-
pKcuv = pKc,v,, + PKt3 + p K g .
(6)
7.O -
6.S-
pH
8.o-
$.5 5.0 4,5 4,O l .o
,
,
,
,
i .=
1.4
i .s
i .s
2,O
FIG. 3. Determination of acidity constants for DCTA at 298 K. The average protonation number, ~ = 2-(Co.- - [OH-])/Yt, w h e r e C O I l - is the concentration of NaOH added by titration, OH is the free hydroxyl concentration and Yt is the total DCTA concentration. I = 0.20 M. a) pK.4 for HI~'-'TA-a is equal to the pH value at ~ = 0.5. b) pK.3 for H2DCTA-2 is equal to the pH value at ri = 1.5.
Values of pKcuv in acetate-free Nat1 solutions are given in Table 1 for Y = DCTA. A value ofpKc,e = 18.65 :t: 0.07 at I = 0.20 is obtained for Cu-EDTA. This agrees very well with MARTELL and SMITH'S(1982) value of 18.70 at 25"C and I --0.1. The solubility of eovellite was measured at various H S and HDCTA-3 concentrations at pH 8.1 and I = 0.20, 0.70 and 1.0 M. The solubility data are listed in Table 2 along with the calculated values of log Koiu and log K~ (defined by reactions 1 and 4, respectively). The reversibility of the solubilities was demonstrated by first allowing an HS- solution to equilibrate with covellite and then doubling the HS- concentration. The Cu concentration was observed first to rise to a plateau, and then after adding the HS-, to fall by a factor of 2 to a new steady state (SHEA, 1985).
232
D. Shea and G. R. Heiz T a b l e 2.
(I)
- l o g M CuT
Covelllta
- l o g M HS-
• 20
6.45+.04
5.50
Solubillty
D a ta .
- l o g M HDCTA- 3
2.30
l o g Kdlss
l o g Ksp
-9.65+.03
-21.39+0.13
.70
6.52~.02
5.58
2.30
-9.80~.02
-20.96+--0.14
.70
6.43~.04
5.54
2.00
-9.97~.04
-21.13+0.16
.70
5.67+.03
5.50
1.30
- 9 . 8 7 -+- . 0 4
- 2 1 . 0 3m +0.16
1.0
6.62+.03
5.64
2.30
--9.96+.04
--20.89+0.16
1.0
6.47+.02
5.52
2.00
--9.99+.03
--20.92+0.15
1.0
5.93+.03
5.45
1.30
- - 1 0 . 0 8 +. 0 3
--21.01+0o15
1.0
6.24+.01
5.08
1.30
--I0.02+.02
--20.95+0.14
1.0
6.81+.03
4.55
1.30
- - 1 0 . 0 6 +. 0 4
--20.99+0.16
4.05
t.30
--
1.0
<7.20
T - 298 K, pH = 8.1 (TRIS b u f f e r e d ) , i o n i c s t r e n g t h Errors represent
( I ) a d j u s t e d w i t h NaCI.
t h e s t a n d a r d d e v i a t i o n of 3 r e p l i c a t e
Kdis8 = [CUT] [HS-]/[HDCTA-3],
Ksp
--
experiments.
= Kdlss KCuDCTA" Ka3
[HS-] d e t e r m i n e d by a a p e r o m e t r i c t i t r a t i o n .
In one test, covellite was leached with DCTA to dissolve 25% of the original mass. Then K.p was measured on the leached solids. The average K,p in this test was within the uncertainty of the K,p determined on unleached covellite. Leaching should have eliminated small particles, strained crystals, sharp corners and other features which might lead to artificially high solubilities. Solubility measurements for poorly crystalline CuS are presented in Table 3. Despite the fact that this material is not a pure phase, it is possible to describe the data with the following equilibrium: x - 1 S(s) + 1 CuxS(s) + HEDTA3-(aq) ~X
X
CuEDTA2-(aq) + H S - ( a q )
Kd~.
3 that reproducible K,p values were obtained for poorly crystalline products prepared independently, at different times. The range of values is actually quite small considering the lack of crystallinity of these samples. The solubility of the poorly crystalline CuS was measured in an EDTA solution as a function of time to determine if reversible equilibrium was achieved. The data are plotted in Fig. 4. It appears that the poorly crystallized CuS reaches metastable equilibrium from both undersaturated and oversaturated solutions within about 6 hours. The poorly crystalline CuS was stable for at least 3 weeks in a 0.20 M NaC1 solution containing 10 mM HS-. However, it recrystallized within about 4 days when 1 mM S(0) was added to this solution in the form of polysulfide ions. The
(7) T a b le 3.
When this is combined with reactions (2) and ( 3 ) we get: H+ + x - 1 S(s) + 1 CuxS(s) ~ Cu2+(aq) + H S - ( a q ) X
S o l u b i l i t y Data f o r Poorly C r y s t a l l i n e CuS.
Sample
- l o g CUT
- l o g KS-
- l o g EDTA
l o g F,Sp
A-1
6.62+.04
5.08
2.30
-17.99
A-1
6.93+. 05
5.12
2.60
-18.04
X
(8) (Cu2+)(HS - ) K,o = (H+)(aso)( x _ l)/x"
(9)
Note that the equilibrium constant expression for (8) is identical in form to that for (4) except for the activity of sulfur term in the denominator. Since the poorly crystalline material contained elemental sulfur, we have assumed that the activity of sulfur is unity. However, this assumption may be a source of systematic error, as we have not proved that the sulfur which was present was in the stable, rhombic form. In any case, the calculated value ofK, p is insensitive to the assumed activity of sulfur because the exponent in Eqn. (9) is small. Values of log K,p are given in Table 3 for each solubility experiment. The mean value, -18.02, is 3 orders of magnitude greater than Ksp for covellite. It should be noted in Table
4.96
3.00
--
A-I
7.12+.09
4.50
2.30
-17.91
A-2
6.75+.06
5.04
2.30
-18.08
A-I
<7.40
A-3
6.58+.03
5.09
2.30
-17.96
A-4
6.82+.06
5.02
2.30
-18.12
A-5
6.70+.07
5.04
2.30
-18.03
T - 298 K, pH - 8 . 1 (TRIS b u f f e r e d ) ,
I = 0 . 2 0 ( a d j u s t e d wlch BaCl),
0 . 1 0 g CuS i n 100 aL s o l u t i o n . E r r o r s a r e s t a n d a r d d e v i a t i o n of 3 r e p l i c a t e S a a p l e s A - l , A-2, A-3, A-4, and A-5 w e r e to t e s t
reproducibility
of th e s y n t h e s i s .
experiments.
made at
different
times.
Solubility product of covellite 55 6.0
Mmm~m"m~
6.5
•
•
o ~ z5 !
8.0 8.5
233
conditions were probably appropriate i'or recrystallization of Rickard's initial precipitates to more stable forms during storage in solution. Above pH 8.5 Rickard observed blaubleibender covellite, rather than normal covellite. For given total polysulfide and bisulfide concentrations, the activity of S(0) will be lower at higher pH. This may account for the failure of sulfur-deficient blaubleibender covellite to recrystallize to covellite in Rickard's more alkaline runs. Free energy o f covellite
0
d5
IiO TIME
'15
2.0
(days)
FiG. 4. ]_08 concentration of Cu as a function of time for poorly crystalline CuS. T = 298 K, ! = 0.2 M. Squares represent the solubitity in a solution containing [ H S - ] = 0.01 raM, [ E D T A ] = 5.0 raM. Triangles represent measurements where the above solution was spiked with 0.09 m M HS- after 8 hours, thus demonstrating reversibility.
It is possible to derive a value for the free energy of formarion ofcovellite from our solubility measurements by using published values for the free energy of formation of Cu 2+ and H S - (from Roam et al., 1978). However, it is first necessary to correct our K~ values to zero ionic strength. Assuming no complexing between Cu 2+ and NaCI, the relationship between K,p in an ionic medium and its corresponding value at infinite dilution (i,e., K°~) would be:
~p "Yc.,÷~.s-K,~
(10)
K,p = Mc,:+MHs-(IO+pH).
(1 i)
=
X-ray diffraction pattern of the solid was taken after storage in the S(0) solution for 14 days. This pattern was similar to that of the covellite, indicating that the poorly crystalline CuS inverts to less soluble covellite in the presence of S(0). The role of S(0) is probably to produce polysulfide ligands by reacting with HS-. Polysulfides f o r m strong complexes with Cu (SHEA and HELZ, 1988), thus increasing solubility and facilitating recrystallization. The thermal stability of poorly crystalline CuS was investigated by placing a bulk sample in an evacuated pyrex vessel and heating to 250°C for 30 days. The X-ray diffraction pattern of the annealed product indicated that it was crystalline coveUite. Although no extraneous peaks were observed, an additional, Cu-rich phase must also have been present to preserve mass balance. The 3 mol% S Opresent in the poorly crystalline product is insufficient to convert it all to covellite.
where
Using our value of Kw at I = 0.2 M and activity coefficients calculated from the Davies equation, we obtain pK~,p= 22.06 and a free energy of formation of covellite o f - 1 1 . 5 kcal/ mole ( - 4 8 . 3 kJ/mole). In Fig. 5, the curve corresponding to Eqn. (10) is shown as a dashed line. This curve fails to pass through the data at I = 0.7 and 1.0. One way of improving the fit to the data is to hypothesize that copper forms complexes with e l - in the NaC1 medium. Assuming the existence of only one such complex, CuCI+, the following modified form of Eqn. (10) can be derived:
DISCUSSION RICKARD (1972) also haS studied precipitates formed by mixing CuSO4 and NariS solutions. In contrast to our results, Rickard failed to observe the poorly crystalline phase, obtaining instead either covellite or blaubleibender covellite, depending upon pH. Rickard's procedure differed from ours in that he used a 30 to 1 molar excess of NariS, whereas we used a stoichiometric, 1 to 1 ratio. Rickard also stored his precipitates in solution for 7 days before X-ray analysis, whereas we vacuum dried our precipitates within an hour of formation. The difference between Rickard's results and ours can be accounted for as follows. The precipitation of CuxS ( x > 1) from a solution containing cupric and bisulfide ions generates S(0) and acidity by the reverse of reaction 8. In our experiments, the final pH would be near one, whereas in Rickard's it would be near its original, neutral to alkaline value, owing to his excess HS-. Under these conditions, S(0) precipitated as solid sulfur in our experiments and was observed in the product. However, it would have remained in solution as polysulfide ions in Rickard's experiments. Since polysulfide ions solubilize copper minerals by strongly complexing Cu,
20
.--21 i 0
MATHIEU I~ I e~ R'CKERTT (1972)/ff " ~ ,
0 J
-D~VTE~- . . . .
et oL(1978)
23 • WAGMAN el" al. (1982) O POTTER ' (1977)
• PANKRATZ, et ol. (1987)
24
6
o12
' o14
' o16
' ale
'
FIG. 5. Extrapolation of K,o to infinite dilution and comparison with published data. Dashed line indicates extrapolation according to equation 10 and solid line indicates extrapolation according to equation 12.
234
D. Shea and G. R. Helz
K°o =
Table 4.
7c.2"7Hs-K~p I +/~c.cl+Mo-Vc.2+
'
(12)
Where K°cucl÷ = acuct-/acu2+act-. We have treated K°,o+ as a variable parameter and chosen its value so as to make a plot of Eqn. (12) exactly fit the data at I = 0.2 and I = 1.0, as shown by the solid line in Fig. 5. Activity coefficients were again evaluated from the Davies equation. The best value of /~cucl. was l 1.1. Independent measurements of this constant tend to be lower than this value (OHLSONand VANNERBURG, 1974), but of similar magnitude. Based on the solid line in Fig. 5, the preferred value of pK~p is 22.27 _-+ 0.3 and the corresponding free energy of formation of covellite is -11.83 -+ 0.4 kcal/mole (-49.50 _+ 1.7 kJ/mole). The uncertainties in the pK~o measurements in NaC1 media are approximately -+0.15, but the uncertainty in pK~o is greater owing to the uncertainty in the extrapolation to I = 0. The uncertainty of +0.3, assigned above, represents our estimate of the likely uncertainty inherent in this extrapolation. The free energy of formation of covellite, in kcal/ mole, is reported to be -11.7 _+ 1.0 by ROBIE et al. (1978), -11.58 by MATHIEU and RICKERT (1972), but -12.89 + 0.05 by POTTER (1977), --12.8 by WAGMANet aL (1982) and -13.44 by PANKRATZ et al. (1987). Our best estimate of AG ° is in excellent agreement with the first two values but clearly can not be reconciled with the latter values (see Fig. 5). Especially troubling is our disagreement with Potter's value, which is a high precision result based on electrochemical measurements. Possibly Potter failed to determine the potential of the covellite-rhombic sulfur assemblage, as he believed, and instead obtained the potential o f a covellite-metastable sulfur assemblage. Amorphous forms of sulfur are well known. The discrepancy between our AG ° value and Potter's is of the right magnitude for a crystalline to amorphous transition. Potter used a sulfide-free, CuSO4 aqueous electrolyte in his cell. This electrolyte is a very poor solvent for elemental sulfur. If amorphous sulfur formed in such a cell, it would not be likely to recrystallize to a stable form. (See also the discussion by NOWAK et aL, 1984, concerning additional potential problems associated with use of CuSO4 electrolyte.) In a paper on the solubility product of galena (UHLERand HELZ, 1984), no correction for Pb-chloride complexing was made in the extrapolation to infinite dilution. When this correction is made with the complex stability constants of SEWARD(1984), a pK°p of 12.78 is obtained. This is lower than the previously recommended value of 12.25, but it is still substantially larger than that obtained by extrapolation of high temperature free energy data, i.e. 14.68 according to ROBIE et aL (1978).
Comparison with FeS Table 4 compares pK~,pvalues ofcovellite and poorly crystalline CuS with analogous phases in the Fe-S system. Also shown are estimated pK~tp values for supercooled CuS and FeS liquids. These latter values, which have been estimated by extrapolation of high temperature thermodynamic data, probably indicate the solubilities of the least stable condensed phases that could possibly precipitate from aqueous solution.
Values of pKsp for CuS and PeS Phases at 2980K. o
pKsp
Phase
Data Source
CuS
Covellite
(CuS)
Poorly Cryscallne Supercooled
CuI.18S
Liquid CuS
22.27
~his study
18.90
This study
15.63
Larraln, et al. 1979
ZeS Troillte
(FeS)
Poorly Crystalline Supercooled
Fel.13 S
Liquid ~eS
6.06
Robie, et el. 1978
2.94
Berner 1964, 1967
0.51
Robie, et al. 1978
Even though the iron phases are uniformly much more soluble than the corresponding copper phases, there are some striking analogies between the two sets of daia. In both cases, the stable crystalline phase is about 6 orders of magnitude less soluble than the amorphous, supercooled liquid, and the poorly crystalline phase is roughly midway between the supercooled liquid and the crystalline phase in solubility. Furthermore, in both cases, the poorly crystalline phase possesses a metal to Sulfur ratio in excess of unity. Surprisingly, these analogies exist despite differences in the crystal structures of covellite and troilitc and despite differences in the common valencies of the two metals.
Phase diagram Using the free energy of formation of covellite together with free energies for other stable Cu-S phases permits construction of a phase diagram, as shown in Fig. 6. The free energy values used were - 17.71 kcal/ mole for anilite, - 19.08 kcal/mole for djurleite and -19.39 kcal/mole for low chalcocite. These values were obtained by correcting POTTER'S (1977) data to be consistent with our free energy for coveUite. These values agree reasonably with those of MATHIEU and RICKERT ( 1972 ). Figure 6 illustrates several important points about the thermodynamics of copper in sulfidic waters. First, it is clear that the solubility of copper in natural sulfidic waters can not be defined without knowledge of the activity of S o. Unfortunately, in very few field studies has an effort been made to quantify this variable. SHEAand HELZ (1988) discuss methods for doing this. Second, two intermediate phases, anilite and djurleite, are stable at temperatures found near the Earth's surface. In most previous efforts to model the behavior of copper in sulfidic waters, saturation only with respect to chalcocite and covellite has been considered. Finally, covellite is the stable phase at high activities of S °, where polysulfide aqueous species have maximum stability. In an effort to account for the solubility of copper in sulfidic waters, some workers have explored equilibria between chaicocite and dissolved copper polysulfide complexes (e.g., BOULEGUE et al., 1982; JACOBS et al., 1985). Such equilibria are metastable. Chalcocite will be converted to a more sulfur-rich solid phase by taking up S(0) from polysulfides until either the chalcocite disappears or the polysullides are depleted to immeasurably low concentrations.
Solubility product of covellite
pooR~ x'rA_.L-~
18
19
-
LOW CHALCOCITE AI
LITE
(
.-.~S
tJ COVELLITE
Cu,,S
•, 2 0 Z
CuS
0.7 M NelCl
// /
22
/
/
' 23
24
t-O
,J JNSATURATED SOLUTION
J i - log
o
as*
FIG. 6. Phase diagram for CuS system at 25°C. Contours indicate solubility limits at various concentrations of NaCI. The activity of sulfur, on the abscissa,is definedwith respect to stablerhombic sulfur. Also shown in Fig. 6 is the solubility of our poorly crystalline material. This has been plotted as a dotted line with a slope consistent with the average composition of Cu,.,sS and an intersection at log as0 = 0 determined by the measured K,p. This is not a rigorous procedure, as the poorly crystalline material is not a pure phase and its composition can be expected to change with as 0. However, the position of the dotted line serves to place this material in the context of the stable mineral phases in the Cu-S system. Possibility o f a poorly crystalline C u S phase in nature In view of the probable role of poorly crystalline FeS in controlling iron solubility in natural waters (DAVISON, 1980), it is of interest to consider whether poorly crystalline CuS may play a similar role. BROCKAMP el al. (1978) investigated the chemical form of Cu and Zn in Red Sea Brine deposits. Some of these deposits contained no X-ray detectable chalcopyrite or sphalerite and yet contained up to several percent Cu and Zn. Using a selective extraction procedure that they claim isolates the sulfide-bound fraction of these metals in the deposits, BROCKAMPet al. (1978) showed that an appreciable fraction of both metals appears to be sulfide-bound. Lacking X-ray diffraction evidence for the existence of crystalline sulfides, they concluded that the Cu and Zn are contained in amorphous sulfide phases. If the poorly crystalline copper sulfide that we have studied is to be found anywhere in nature, one would expect to find it where solutions of markedly different composition and temperature mix, as at the top of the Red Sea brine pools. Aside from rapid mixing of disparate solutions, it is difficult to envision physical processes in nature that could produce the high degree of supersaturation required to precipitate the poorly crystalline material. Nevertheless, the thermal instability of this material and its rapid conversion to covellite when exposed to solutions containing S(0) in the form of polysulfides, suggests that this material is unlikely to be a persistent phase in nature. To
235
find it would at the very least require particular precautions during sampling to prevent production of S (0) by oxidation. No such precautions were reported by BROCKAMP et al. (1978). On the basis of these considerations, we doubt that the poorly crystalline material we studied is the X-ray amorphous material inferred to exist in the Red Sea Brine deposits. Because of its instability, poorly crystalline CuS would be expected to have only an ephemeral existence in nature and would not be expected to accumulate in sediments. However, the possibility remains that the poorly crystalline phase could control copper solubilities for short periods in highly dynamic environments. If indeed a non-crystalline form of sulfide-bound copper exists in the Red Sea deposits, then we suspect it is a Cu-Fe sulfide phase. Phases in the Cu-Fe-S system are more sluggish in equilibrating than phases in the Cu-S system (BARTON and SKINNER, 1979). Thus the chances are better that an amorphous Cu-Fe sulfide would be persistent enough to be observed. Acknowledgements--Acknowledgement is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society for financial support. The National Bureau of Standards provided access to the electron microscope, and this work was assisted by Patrick Sheridan. Editorial handling: S. A. Wood REFERENCES ANDEREGGG. (1967) Komplexone, L.X. Die Protonierungskonstanten einiger Komplexone in verschiedenen wasserigen Salzmedien. Heir. Chim. Acta 50, 2333-2340. BARTONP. B. JR. and SKJNNERB. J. (1979) Sulfidemineralstabilities. In Geochemistry of Hydrothermal Ore Deposits (ed. H. L. BARNES), 2nd edn., pp. 278-403. L Wiley & Sons. BERNERR. A. (1964) Iron sulfidesformed from aqueous solution at low temperatures and atmospheric pressure. J. Geol. 72, 293-306. BERNERR. A. (1967) Thermodynamic stability of sedimentary iron sulfides. Amer. J. Sci. 265, 773-785. BJERRUMJ. ( 194l) Metal Amine Formation in Aqueous Solution. P. Haase & Sons, Copenhagen, 208p. BOULEGUEJ., LORDC. J. and CHURCHT. M. (1982) Sulfurspeciation and associated trace metals (Fe, Cu) in the pore waters of Great Marsh, Delaware. Geochim. Cosmochim. Acta 46, 453--464. BROCKAMP O., GOULART E., HARDER H. and HEYDEMANN A.
(1978) Amorphous copper and zinc sulfides in the metalliferous sediments of the Red Sea. Contrib. Mineral Petrol. 68, 85-88. CARtGNANR. and NRIAGUJ. O. (1985 ) Trace metal deposition and mobility in the sediments of two lakes near Sudbory, Ontario. Geochim. Cosmochim. Acta 49, 1753-1764. DAVlSONW. (1980) A criticalcomparisonof the measuredsolubilities of ferrous sulfide in natural waters. Geochim. Cosmochim. Acta 44, 803-808. DYRSSEND. (1985) Metal complex formation in sulfidic seawater. Mar. Chem. 15, 285-293. EMERSONS., JACOBSL. and TREBOB. (1983) The behavior of trace metals in marine anoxicwaters: Solubilitiesat the oxygen-hydrogen sulfide interface. In Trace Metals in Sea Water (eds. C. WONGet al. ), pp. 579-608. Plenum Publishing. GAILLARDJ.-F., JEANDEL C., MICHARD G., NICHOLAS E. and RENARD D. (1986) Interstitial water chemistry of Villefranche Bay sediments: Trace metal diagenesis.Mar. Chem. 18, 233-247. GOBLER. J. (1980) Copper sulfides from Alberta: Yarrowite CugSs and Spionkopite Cu39S2s. Can. Mineral. lg, 511-518. GOBLER. J. and SMITHD. E. W. (1973) Electron microprobe investigation of copper sulfides in the Precambrian Lewis Series of S.W. Alberta, Canada. Can. Mineral. 12, 95-103.
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D. Shea and G. R. Helz
JACOBSL. and EMERSONS. (1982) Trace metal solubility in an anoxic fiord. Earth Planet. Sci. Lett. 60, 237-252. JACOBSL., EMERSONS. and SKEIJ. (1985) Partitioning and transport of metals across the O2/H2S interface in a permanently anoxic basin: Framvaren Fjord, Norway. Geochim. Cosmochim. Acta 49, 1433-1444. KltEMLING K. (1983) The behavior ofZn, Cd, Cu, Ni, Co, Fe and Mn in anoxic Baltic waters. Mar. Chem. 13, 87-108. LARRAINJ. M., LEE S. L. and KELLOGGH. H. (1979) Thermodynamic properties of copper-sulfur melts. Can. Metallurg. Quart. 8, 395-400. MARGERUMD. W. and BYDALEKT. J. (1963) Multidentate ligand kinetics. V. Copper (1I) and cyciohexylenediamine tetraacetonickelate (II). Inorg. Chem. 2, 683-688. MARTELLA. E. and SMITHR. M. (1982 ) Critical Stability Constants. Vol. 5, First supplement. Plenum Press, New York. MATHIEU H. J. and RICKERT H. (1972) Eiectrochemischthermodynamische Untersuchungen am System Cu-S. Z. Phys. Chem. Neue Folge 79, 315-330. NOWAK P., BARZYK W. and POMIANOWSKIA. (1984) The applicability of EMF measurements to evaluation of thermodynamic properties of the Cu-S system. J. Electroanal. Chem. Interfacial Electrochem. 171, 355-358. OHLSON M. and VANNERaERGN.-G. (1974) Stability constants for chloride complexes of copper (lI) in sulfuric acid solution, Acta Chem. Scand. A28, 1031-1035. PANKRATZL. B., MAH A. D. and WATSONS. W. (1987) Thermodynamic properties of sulfides. U.S. Bur. Mines. Bull. 689, 427p.
POTTER R. W. (1977) An electrochemical investigation of the system copper-sulfur. Econ. Geol. 72, 1524-1542. POTTER R. W. and EVANS H. T. (1976) Definitive X-ray powder data for covellite, anilite, djudeite and chalcocite. J. Res. U.S. Geol. Surv. 4, 205-212. PUTNIS A., GRACEJ. and CAMERONW. E. (1977) Blaubleibender covellite and its relationship to normal covellite. Contrib. Mineral. Petrol. 60, 207-217. RICKARD D. Z. (1972) Covellite formation in low temperature aqueous solutions. Mineral. Deposita 7, 180-188. ROaIE R. A., HEM]NGWAVR. S. and FISHERJ. R. (1978) Thermodynamic properties of minerals and related substances at 298.15 K and l bar pressure and at higher temperatures. U.S. Geol. Surv. Bull. 1452, 456p. SHEAD. ( 1985 ) Solubility and precipitation of copper sulfide. Ph.D. dissertation, Univ. of Maryland, 250p. SHEA D. and HELZ G. R. (1988) The solubility of copper in sulfidic waters: Sulfide and polysulfide complexes in equilibrium with covellite. Geochim. Cosmochim. Acta 52, 1815-1825. SEWARDT. M. ( 1984 ) The formation of lead (II) chloride complexes to 300oC: A spectrophotometric study. Geochim. Cosmochim. Acta 48, 121-134. UHLER A. D. and HELZ G. R. (1984) Solubility product of galena at 298 K. Geochim. Cosmochim. Acta 48, 1155-1160. WAGMAN D. D., EVANS W. H., PARKER V. B., SCHUMM R. H., HALOW 1., BAILEYS. M., CHURNEY K. L. and NUTTALLR. L. ( 1982 ) The NBS tables of chemical thermodynamic properties. J. Phys. Chem. Ref. Data. I I: Supplement 2, 392p.