Table 111. Measurement Differences Between Methods
Anaheim Los Angeles San Bernardino San Diego
Blas, pprn SO2 a Mean and estd SD
Mean and estd SD, pprn SO2 Conductlrnetrlc Pararosanlllne
Data palrs
0.0057 f 0.007 1 0.0191 f 0.0081 0.0064 f 0.0067 0.0019 f 0.0030
160 101 130 25
0.0050 f 0.0034 0.0071 f 0.0050 0.0024 f 0.0017 0.0035 f 0.0034
0.0008 f 0.0065 0.0120 f 0.0072 0.0040 f 0.0072 -0.0016 f 0.0040
Max
0.024 0.040 0.029 0.013
ConductimetricSO2 minus pararosaniline Son.
Literature Cited
Table IV. Relative Bias at 0.04 ppm SOz by Each Method Conductlrnetrlc SO2 Blas at 0.04 pprn Carrel Cond SOz, a ppm coeff
Anaheim Los Angeles San Bernardino San Diego Mean
0.0282 0.0267 0.0391 0.0258 0.0300
0.65 0.79 0.54
0.55
(1) “Directory of Air Quality Monitoring Sites Active in 1973”,
Pararosanlllne SO2 Bias at 0.04 ppm Corre~ PRA SOz, a pprn coeff
0.0008b 0.0048 -0.0582’ -0.0301 -0.0207
* By linear regression. * Mean bias, not significant. meaningful negative bias.
0.04 ppm
0.08 0.15 0.40 0.67
is largest
Conclusions Estimated bias between the conductimetric and pararosaniline methods a t 0.04 ppm SO2 by the conductimetric method ranges to the order 0.04 ppm. Estimated random measurement error ranges (99.70/0 confidence) are f0.02 and f0.006 ppm for the conductimetric and pararosaniline methods, respectively. The conductimetric method is unsuitable for determining compliance with 0.04 or 0.05 ppm 24-h average ambient air quality standards for SOz. Acknowledgment
N. R. Crawford and F. W. Morgan’s statistical advice and assistance and E. G. Loffay and J. R. Ugolini’s computer data handling were invaluable.
EPA-450/2-75-006, Mar. 1975. (2) Cooper, R. E., “Statistics for Experimentalists”, Pergamon Press, 1969. (3) Blacker, J. H., Confer, R. G., Brief, R. S., J. Air Pollut. Control Assoc., 23, 525 (1973). (4) Hocheiser. S.. Santer. J.. Ludmann, W. F., ibid., 16, 226 (1966). (51 Terabe. M.. Oomichi. S..Benson. F. B.. Newill. V. A.. Thomoson. J. E., ibid., 17,673 (1967). (6) Staff Rep. 75-12-2, “Consideration of the California 24-Hour Ambient Air Quality Standard for Sulfur Dioxide”, California Air Resources Board, June 11,1975. (7) , , Booras. S. G.. Zimmer. C. E.. J. Air Pollut. Control Assoc... 18.612 , (1968). (8) Stevens, R. K., Hodgeson, J. A., Ballard, L. F., Decker, C. E., “Ratio of Sulfur Dioxide to Total Gaseous Sulfur Compounds and Ozone to Total Oxidants in the Los Angeles Atmosphere-An Instrument Evaluation Study”, in “Determination of Air Quality”, G. Mamantov and W. D. Shults, Eds., pp 83-108, Plenum, New York, N.Y., 1972. (9) “Air Quality Criteria for Sulfur Oxides”, National Air Pollution Control Admin., Jan. 1969. (10) Katz, M., “Inorganic Gaseous Pollutants”, in “Air Pollution”, A. C. Stern, Ed., 2nd ed., Chap. 17, Academic Press, New York, N.Y., 1968. (11) Neuscheler, R. C., “Selection of Continuous Sulfur Dioxide Monitors for Ambient and Source Concentration Levels”, “Instrumentation for Air Monitoring”, ASTM STP 555, pp 9-19, American Society for Testing Materials, 1973. (12) Purdue, L. J., “Performance Evaluation of SO2 Monitoring Instruments”, ibid., pp 3-8. (13) “The Environmental Protection Agency’s Research Program with Primary Emphasis on the Community Health and Environmental Surveillance System (CHESS):An Investigative Report”, for Committee on Science and Technology, U.S.House of Representatives, Nov. 1976. I
,
Received for review J u l y 22,1977. Accepted October 25,1977.
Limits of Coprecipitation of Cadmium and Ferrous Sulfides Paul E. Framson” and James 0. Leckie
.
Environmental Engineering and Science, Department of Civil Engineering, Stanford University, Stanford, Calif. 94305
Cadmium was precipitated from aqueous sulfide to ascertain the chemical nature of cadmium sulfide phases precipitated in aquatic ecosystems. Monitoring cell-dimension trends of sulfides precipitated from thioacetamide solutions of increasing ferrous content allowed estimation of an upper bound of 3.5 10-4 for the distribution coefficient describing Fez+ cocrysta~~ization with greenockite. This result suggests that cadmium precipitates primarily through surface exchange with ferrous monosulfide substrates and as essentially unsubstituted CdS in natural sulfidic systems.
Sediments possess the interesting ability to act as both a sink and a source for many trace contaminants. For example, significant spatial and temporal variations in zinc and cad0013-936X/78/0912-0465$01.00~0 0 1978 American
Chemical Society
mium in the Corpus Christi, Tex., estuarine system demonstrate that those sediments act alternately as a sink during the summer months when high concentrations of sulfides are present in the surficial sediments and as a source during the winter when more oxidizing conditions prevail (1).In anoxic benthic marine, estuarine, and freshwater environments, sulfate-reducing bacteria can generate total aqueous sulfide as high as millimolar concentration ( 2 , 3 ) .Under sufficiently reducing conditions, common sedimentary mineral phases such as hydrous oxides and hydroxides are replaced by sulfides. The pK,, of the heavy metal sulfides range from 25.2 for ZnS(s) to 53.2 for HgS(s) ( 2 ) .Hence, a large fraction of aqueous heavy metal contaminants is likely bound as sedimentary sulfides under anaerqbic conditions. Certain human activities such as dredging, which disrupts sediments, can effect oxidative dissolution of ferrous sulfides, Volume 12, Number 4, April 1978 465
prompting concern that associated heavy metals may also be released from their sulfide phases. To resolve this concern, we must first ascertain the chemical nature of the heavy metal sulfide phase precipitated under natural conditions. Cadmium was selected as the subject in the endeavor to answer this question due to the relatively extensive information concerning the aqueous and surface chemistry of cadmium and because Cd is known to be toxic. Because in aquatic sediments, iron sulfides are the dominant sulfides, cadmium will exist in intimate association with iron in the sulfide solid phase. The possibilities for cadmium precipitation as its sulfide are: cadmium surface exchange with or adsorption onto iron(I1) sulfide substrate; solid solution formation of ferrous and cadmium sulfide: cocrystallization as cadmium sulfide host phase substituted with Fez+, or cocrystallization as ferrous sulfide host phase substituted with Cd2+;and precipitation as essentially pure cadmium sulfide. This study suggests which of these phenomena are most probable.
Theory We first consider the general thermodynamic boundaries of equilibrium cocrystallization. The substitution of a foreign cation B+ into the host solid AX may be considered a two-step process of surface exchange followed by diffusion and incorporation into the crystal lattice ( 4 ) , A+(ads)
+ B+(aq)
+
B+(ads) AX(so1id)
--
B+(ads)
+ A+(aq)
(1)
BX(so1id)
+ A+(ads)
(2)
BX(so1id)
+ A+(aq)
(3)
the overall equilibrium being AX(so1id)
+ B+(aq)
The equilibrium distribution coefficient D for Equation 3 is defined as
D
= (E)solici/
(E)q
(4)
where x denotes mole fraction in the solid phase. Assuming = 1,D can be formulated from thermodynamic parameters as
XAX
where K,, is the equilibrium ion activity product, y2 the mean aqueous activity coefficient product, f the solid state activity coefficient, R the universal gas constant, and T the absolute equals the modification of the free energy temperature. of the AX lattice due to the substitution of B+, or through slight reformulation, the change in the chemical potential of the solid solute B X due to nonideality of solid solution. Equation 5 requires that some B X be incorporated into the AX lattice when [B+] is nonzero. Were A G i x zero, D would nearly equal the ratio of the ion activity products. Seldom zero however, AGix is minimized and D maximized when the mixed crystal least distorts the host lattice and offers the solute an environment most identical to that in its own Lure solid. We speculate shortly on the relative values of A G i x terms in the Fe-Cd-S system. This thermodynamic perspective on cocrystallization is properly qualified by the fact that cocrystallization does not often proceed in a near equilibrium fashion because surface exchange rates generally radically exceed solid state diffusion rates. Typically, there exists a direct relationship between amount of a specie precipitated from solution and concentration remaining in solution a t that instant as stated in the Doerner-Hoskins relation,
hcA,
466
Environmental Science & Technology
where dB and dA represent increments of B + and A+ precipitated, and [B+]and [A+] represent their aqueous concentrations. The heterogeneous distribution coefficient describes cocrystallization resulting in a heterogeneous distribution of foreign species within the host lattice when aqueous solution composition changes during the course of precipitation. X approaches D as the precipitation rate approaches zero. Surface reaction constitutes not only a critical step in cocrystallization, but in and of itself, a prominent course of coprecipitation. Most widely documented of sulfide surface reactions is the exchange reaction conforming to the metathetical equilibrium K
+
AzlmS + 2 / n B f nf) 2/mA+m B21,S; K = Kf:/K::
(7)
whereby aqueous cation B+“ displaces cation A + m from the solid sulfide when the solubility of B21nS is less than that of A21,S. Accordingly, ferrous sulfide, nearly the most soluble of the heavy metal sulfides, has been enlisted as an effective analytical adsorbent for other heavy metal cations including Cd2+( 5 ) ,whereas cadmium sulfide was found to remove from aqueous solution Fe2+ most feebly among seven cations studied (6). The metathetical reaction typically proceeds by rapid formation of a mono- to trilayer coating of the displacing cation, whereas subsequent formation of a new sulfide crystalline phase proceeds much more slowly reflecting low solid state diffusion rates. Phillips and Kraus ( 7 ) document formation of new crystalline phases by Ag+ and Cu2+conversion of zinc, cadmium, lead, arsenic, or cupric sulfides. Gaudin et al. (8) also find Ag+ to convert zinc sulfide to a new crystalline phase, notably rejecting mixed crystal formation as a mechanism for this process. James and Parks (9) find two models to describe moderately well the nonmetathetical reaction of aqueous Zn2+ with solid HgS. The first model postulates a simple adsorption reaction whose driving force arises from electrostatic attraction between surface and adsorbate, change in the solvation energy of adsorbate, and modification of the chemical free energy of the adsorbate; the other model postulates an exchange between adsorbed H+ and Zn2+ in the sulfide surface layer. To speculate on the degree to which coprecipitation in the Fe-Cd-S system involves not only surface phenomena but crystal lattice phenomena as well, a brief examination of the crystal chemistry of pertinent solids in that system is useful (Table I) (10-12). Cadmium sulfide precipitates from aqueous solution a t ambient temperatures and pressures as either of two polytypes, greenockite and hawleyite. Both are founded on close-packed arrays of sulfur, greenockite possessing hexagonal- and hawleyite cubic-close packing. Hawleyite is metastable to greenockite at 25-900 “C (13),but the two polytypes are so nearly isoenergetic that the coordination state of the aqueous Cd2+ specie kinetically influences which of the two precipitates. Halide ligands very strongly dictate greenockite crystallization whereas SO:- and NO,, which complex Cd+2 less strongly, favor hawleyite crystallization (14).Therefore, greenockite exclusively should precipitate from marine waters, whereas hawleyite should precipitate as well to a minor degree in freshwater environments. A host of different iron sulfides form in aqueous systems (3,15,16).Mackinawite, FeSI-,, is thought to control aqueous [Fe2+]in freshwater-reducing environments, whereas greigite, Fe3S4, is suspected of controlling [Fez+] in some marine-reducing environments (17).Both of these forms are metastable to yet other naturally occurring iron sulfides. However, be-
Table 1. Crystal Chemistry of Minerals in Fe-Cd-S System a Metal Ion
Cd2+
Fe2+
a
Coordination no., spin
Bond dislances
ionic radius
IV VI
IV High VI Low High
(A)
Mineral
0.84 0.95
Greenockite CdS Hawleyite CdS Mackinawite FeS1-,
0.63 0.61 0.77
(A)
Description
Sulfur in hcp array. Cd2+ occupies half the tetrahedral interstices. Highly ionic Same as greenockite except sulfur in ccp array Sulfur in nearly ccp array. Fe2+ tetrahedrally coordinated, high spin. Strong Fez+ back-donation to sulfur. Metallic bonding
Cd-S
= 2.53
Cd-S
= 2.52
Equll activity product
Fe-S = 2.23 Fe-Fe = 2.60
Modified from refs. 75-77.
cause mackinawite is that form essentially always precipitated initially, only mackinawite is considered in this discussion of coprecipitation with cadmium sulfide. Cadmium sulfide is among the most ionic of sulfides; therefore, Fe2+ substitution in the CdS lattice is controlled primarily by geometric restrictions. The cubic- and hexagonal-close packed sulfur frameworks of the cadmium sulfides are highly stable, quite independently of the identity of the cation that fills the interstices (18).Pappalardo and Dietz (19) found Fe2+to diffuse at 600 "C into the tetrahedral interstices of CdS. In fact, Skinner and Bethke (20) found Fez+ to substitute to at least 0.47 mol fraction in hydrothermally (800-950 "C) formed greenockite, despite the approximate 25% difference in fourfold coordinated Cd2+and Fe2+ionic radii (see Figure 1). (Mole fraction of cation 1 is here defined as nc,/(Zzo nc,),where nc, is the number of moles of ith cation Ci in a solid containing rn distinct cations.) It is anticipated that Fe2+ substitutes to a similar degree in hawleyite, the immediate CdSi- polyhedra being essentially identical in the two polytypes. Relatively then, the A?& term for Fe2+ cocrystallization with greenockite and hawleyite is expected to favor this process. Geometrical restrictions alone highly disfavor Cd2+ substitution into mackinawite, considering that the Cd-S bond distance exceeds the Fe-S bond distance by 0.3 A. Additionally, mackinawite possesses a highly covalent bond character due to strong Fe2+ d-orbital back-donation to sulfur. This covalency is enhanced by metallic d-orbital overlapping between neighboring Fez+ ions, which display a conspicuously low Fe-Fe distance of 2.60 h;. By populating the o* energy levels and overlapping unfavorably with both neighboring anions and cations, Cd2+would highly destabilize the covalent mackinawite structure. This prediction is supported by comparison with the FeZn-S system. Pyrrhotite, Fel-,S, is a highly covalent ferrous sulfide similar to mackinawite. The ionic radii of tetrahedrally coordinated high spin Fe2+ and Zn2+ differ by only 6%, and the Zn2+ orbital energy levels are more nearly isoenergetic with those of Fe2+than are those of Cd2+. Suggested then by better orbital overlap and geometrical compatibility, Zn2+ substitution in pyrrhotite is favored more highly than Cd2+ substitution in mackinawite. Yet, Barton and Toulmin (21) found that Zn2+ substitutes insignificantly at 850 OC-less than 0.006 mol fraction into the pyrrhotite structure. Hence, the A c ~ , sterm is predicted to highly disfavor Cd2+ cocrystallization with mackinawite.
Experimental Work and Results In anticipation that ferrous sulfide and cadmium sulfide behave as two solids of limited mutual miscibility, work began with an attempt to precipitate both solids from a solution
800-950°C
'.'.
6 50
-
6.403+----7
3900 01 0 2 0 3 04 0 5 06 07 08 09 I O Fe/(Fe+Cd),
Mole Fraction in Solid
Figure 1. Cell dimensions of hydrothermal Fe-bearing greenockites (modified from reference 20)
initially equimolar in ferrous iron and cadmium. Hopefully, these two solids could then be separated and analyzed to determine the extent of cocrystallization. Solutions, to 10-l molar NazS, CdS04, and FeS04, were prepared in essentially anoxic conditions, and slow conventional precipitations were effected by "titrating" one of the reactant species with the other with constant stirring, perfusion with N2, and in some cases, pH-statting. This approach confirmed the thermodynamic prediction that no ferrous sulfide forms until essentially all Cd2+has precipitated from solution when aqueous sulfide is not present in excess. Conditions of sulfide excess produce an intractable ferrous sulfide colloid in which no cadmium sulfide is detected visually. Furthermore, this approach is plagued by the following adversities: solids formed are nearly universally colloidal rendering separation and analytical techniques difficult; Fe(I1) ion oxidation by 0 2 is difficult to eliminate entirely; and the precipitation microenvironment cannot be controlled sensitively enough to prevent slight formation of iron hydroxides. After attempts to precipitate only CdS from solutions of varied [Fe2+] were frustrated by cogeneration of iron hydroxides, the entire conventional precipitation technique was discarded in favor of precipitation from homogeneous solution (PFHS),whereby one of the reactant precipitating species is Volume 12, Number 4, April 1978
467
the re11 dimension values and respective srandard denations appearing in Tables I1 and 111 and Figures 3 and 4. Simultaneous atwmpts to produce ferrous sulfide by T A A hydrolysis surreeded only under mildly alkaline cnnditions and werecomplirated by rhecogenerarim ofminor amounts of iron hydroxides.
Discussion and Conclusion The trends in CdS cell dimensions demonstrate that Fez+ does not cocrystallize significantly with CdS over the range of aqueous solution compositions studied at 25 "C. Due to similarity in crystal structure, hawleyite would be expected to exhibit a decrease in cell dimension with increasing iron content similar in magnitude to that of greenockite. Data from precipitate Ga allow calculation of an upper limit for the distribution coefficient,
Figure 2. Comparison of powder patterns of hawleyite precipitated by thioacetamide hydrolysis (left)and by conventional mixing (right)
formed slowly and uniformly throughout the precipitation medium. P F H S allows standardization of the precipitation microenvironment and reproduction of the slow rate of precipitation in natural sediments. This slow precipitation additionally a more highly . . crystalline precipitate . generates (Figure 2) and allows a closer approach to equilibrium . .. distribution of snecies between solid and a aueous solutions. Subsequent CdS precipitations were ef fected by slow acid hydrolysis of thioacetamide (TAA) undeir strict anoxic conditions. Senarate series utilizine C1- and SO?? 30:- environments ;e and greenockite. were run to enable study of both hawleyite ions of varied Fez+ These solids were precipitated from solutions concentrations, aged for 10 days to two weeks, and freeze dried. Dehve Scherrer powder diffractions were performed on the precipiratcc in an arfempt to defwt any trends in cell dimensions rhat would signal Fe'- cocrystilllizarion. T h e hswleyites were irradiared fur 20 h w i t h Mn-filtered Fe Kti radiation, and the greenockites for 12-15 h with Ni-filtered Cu K a radiation. The reflection Hragg angles were measured by traveling microscope, modified hy internal standard cim rections (except i n the cases of two hawlvyite samples!, and indwed. The resultnnt data were then subjected to a IWSIsquares cell dimeniion refinement program (221, generating ~
~~
Noting that greenockite cell dimension e exhibits the greatest sensitivity to FeS mole fraction io the solid, we apply Skinner and Bethke's (20) equation relating FeS mole fraction to cell dimension c to the observed difference in c dimension exhibited by greenockites GI and Gs. This calculation generates a maximum Z F value ~ for greenockite Ga of 0.074 f 0.019. Gs precipitated from a solution possessing an average value of [Fe2+]/[Cdz+]equal to 2.3 X lo2 during the course of precipitation (computing nonchlorocomplexed ion concentrations; stability constants are from refs. 3 and 23). The above distribution coefficient is then at maximum 3.5 f 1 X lo-*. Comparison with the work of Rittner and Schulman (13) yields insight into the chemical basis of this very low distribution hution coefficient. They found Cd2+to cocrystallize significantly with HgS under conditions similar to those of this
I
. -~.~.-.. .. .~~..._ ~ ~ . ~ . . ~- ._. .._ . ~~.~..~ - . ~.~ Table 11. Hawleyite Precipitations by TAA Hydrolysis H, H2
H3
Total [metal] m 0.1 M Total [SO:-] '5. 0.1 M [TAA] IU 0.025 M
5.833 f 0.002 5.83 f 0.01
0 0.050 0.10
Initial nH = 1.5 ~~
Ha H5
~~~
Final pH = 1.5 T = 22-25OC
5.834 f 0.001 5.83f 0.01
0.15 0.20
Table 111. Greenockite Precipitations by TAA Hydrolysis a "l111.1
PreclpIldbn
Gs Gio GI,
0.0800 0.0800 0.0800 0.0800 0.0800 0.0800 0.0400 . 0.00335 0.00336 0.00337 0.00335
T = 22-25
"c.
GI GZ
G G4 G5
G5
GI G8
468
Cdt (molar)
lnlllal
lnlllal
Fa,
WAI
(molar)
(molar)
0
0.02 0.02 0.02 0.02 0.02 0.02 0.01 0.003 0.003 0.003 0.03
0.0080 0.0200 0.0400 0.0600 0.0800 0.0600 0.0301 0.0634 0.33 2.5
Environmental Science a Technology
l"lIl*l PH
1.4 1.4
1.4 1.4 1.4 1.4 1.4 1.5 1.5 1.5 1.5
&Yi (molar)
0.200 0.216 0.240 0.280 0.320 0.360 0.240 0.100 0.167 0.70 5.0
Preclpllate cell dlmenslons Innlsl Fetl(Cdt +Fat)
0 0.091 0.200 0.33 0.428 0.50 0.60 0.90 0.95 0.99 0.995
and SO a
(A)
= (A)
4.130 + 0.001
6.722f 0.005
4.129 f 0.001
6.703 f 0.002
4.131 f 0.002 4.1305 f 0.0009 4.1295 f 0.0009
6.696 0.005 6.695 f 0.002 6.695 f 0.002
No CdS precipitate
No CdS precipitate
+
I
25-C
t
[TAA l 5]
5.90
i
5.70 5.60 I 0
1
Eiperimcntol Conalttons
0.025M
[SO:-],
0.1 M
[Mstol],
0.1 M
1 i
I
0.1
0.2
1
I
0.3
0.4
0.5
Initial Fe+/(Fe, +Cdt)
Figure 3. Cell dimension trend of hawleyites precipitated from solutions of increasing Fe mole fraction
6’80r--l-l
term, the AC!esterm must highly disfavor Cd2+cocrystallization with FeS. Hence, despite the favorable value of 10s for K,F,”S/K,C,dS, Cd2+ cocrystallization with ferrous sulfide a t ambient conditions is predicted to be as limited as Fe2+ cocrystallization with cadmium sulfide. Though formulation of a well-grounded conclusion demands further study, especially of Cd2+ cocrystallization with ferrous sulfide, this study suggests strongly that crystalline solids precipitated from the aqueous Fe-Cd-S system are highly unsubstituted with the foreign cations Fe2+ or Cd2+. The dominant forms of cadmium precipitation in a natural sulfidic system are, therefore, predicted to be Cd2+ surface exchange with ferrous sulfide substrate, and precipitation as essentially unsubstituted CdS. Acknowledgment
The authors thank Gordon E. Brown, Michael B. Nelson, Mark Taylor, and Keith Keefer for their technical advice and assistance. Literature Cited (1) Holmes, C. W., Slade, E. A., McLerran, C. J., Enuiron. Sci. Technol., 8,255 (1974). (2) Leckie, J. O., Nelson, M. B., “The Role of Natural Heterogeneous
Sulfide Systems in Controlling the Concentration and Distribution of Heavy Metals”, Dept. of Civil Engineering, Stanford University, Stanford, Calif., presented at the 2nd Int. Symp. on Environmental Biogeochemistry, Burlington, Ont., Canada, Apr. 1975. (3) Goldhaber, M. B., Kaplan, I. R., in “The Sea”, Vol 5, E. D. Goldberg, Ed., Wiley-Interscience, New York, N.Y., 1974. (4) Walton, A. G., “The Formation and Properties of Precipitates,” Interscience, New York, N.Y., 1967. (5) Caletka, R., Tymbyl, M., Kotas, P., J . Chromatogr., 111, 93
25.C.
pH I 5
3900 01 0 2 03 04 05 06 07 08 09 I O Initial Fet/(Fe++Cdt)
Figure 4. Cell dimensions of greenockites precipitated from solutions of increasing Fe mole fraction
study. Consider again the thermodynamic formulation of the equilibrium distribution coefficient:
For Cd2+ cocrystallization with HgS, K$s/K,c,ds = 10-52/10-26 = 10-26 whereas for Fe2+ cocrystallization with CdS,
K,c61s/Kf;s = 10-268/10-18 = 10-8 Since the ratios of solubilities favor Fe2+ cocrystallization with CdS far more than Cd2+cocrystallization with HgS, the A?&S term must, relative to the ACHgsterm, disfavor Fe2+ cocrystallization with CdS. Considering that the AG& term was predicted to favor cocrystallization far more than the A?&es
(1975). (6) . , Tvaeai.V. A.. Petrova. N. A..Treskunov.R. L..Electrokhim.. . 4.. (2 ); 119 (1968): (7) Phillips, H. O., Kraus, K. A., J . Chromatogr., 17,549 (1965). ( 8 ) Gaudin, A. M., Fuerstenau, D. W., Turkanis, M. M., Trans. A I M E , 268,65 (1957). (9) James. R. 0..Parks. G. A.. AIChE S v m o . Ser.. 150.157 (1975). (10) Wuensch, B. J., MineraL‘Soc. Am.,Shbrt Course Notes, I, W1 (1974). (11) Shannon, R. P., Prewitt, C. T., Acta Crystallogr., B,25 (51,925 (1969). (12) Uda, M., 2. Anorg. Allg. Chem., 361,94 (1968). (13) Rittner, E. S., Schulman, J. H., J. Phys. Chem., 47, 537 (1943). (14) Sato, R., Itoh, H., J p n . J. Appl. Phys., 3 (lo),625 (1964). (15) Berner, R. A., J . Geol., 72,293 (1964). (16) Rickard, D. T., Stock. Contrib. Geol., 20,49 (1969). (17) Doyle, R. W., Am. J. Sci., 266,980 (1968). (18) Evans. R. C.. “Crvstal Chemistrv”. “ . Cambridge Univ. Press. ‘ London, England, 1964. (19) Pauualardo, R., Dietz, R. E., Phys. Reu., 123 (4). 1188 (1961). (20) Skinner, B. J., Bethke, P. M., A h . Mineral., 46,1382 (1961). (21) Barton, P. B., Jr., Toulmin, P., Econ. Geol., 61,815 (1966). (22) Evans, H. T., Jr., Appleman, D. E., Handwerker, P. S., A m . Crystallogr. Assoc., A n n . Meeting Program, 42 (1963). (23) Sillen, L. G., Martell, A. E., “Stability Constants of Metal-Ion Complexes”, Suppl. # 1, Chemical Society, London, England, 1971. 1
Received for review January 6, 1977. Accepted October 31, 1977. Partial support from Air Force contract F29601-75-C-0028.
Volume 12, Number 4, April 1978
469
