J. Phys. Chem. 1992, 96, 4382-4388
4382
Oxidation Klnetics of Uitrasmali Colloidal Chalcopyrlte (CuFeS,) with One-Electron Oxidants Ewen J. Silvester,+Franz Crieser,**+Dan Meisel,*?*Thomas W. Healy: and James C. Sullivan$ School of Chemistry, University of Melbourne, Parkville, Victoria, 3052 Australia, and Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439 (Received: November 8, 1991)
The sto ped flow technique has been used to study the oxidation of ultrasmall colloidal chalcopyrite (CuFeS2) particles (diameter 50-90 by the one-electron oxidants hexacyanoferrate(II1) (Fe(CN),’-) and hexachloroiridate(1V) (IrC162-)in aqueous solution. The stoichiometry of the initial oxidation step corresponds to the reaction CuFeS2(,)+ 5H20 Cu(OH),(,, + Fe(OH),,,) + 2S0,,, + 5H+ + 5e-. Further oxidation of the elemental sulfur results in the formation of S2032- and S042in solution, as determined by ion chromatography. Time-resolved measurement of both the oxidant and chalcopyrite concentrations has provided a detailed understanding of the mechanism involved in the oxidation of colloidal chalcopyrite.
1)
-
Introduction The oxidation of chalcopyrite has received considerable attention due to the possible applications in the hydrometallurgy and flotation of this mineral.’-3 These studies have generally been performed on solid CuFeS2 electrodes or suspensions of large particle size in which very little of the material (), does overlap with the wavelength range studied, the extinction coefficients at the wavelengths of (1) Munoz, P. B.; Miller, J. D.; Wadsworth, M. E. Metall. Trans. B 1979, IOB, 149. (2) Dutrizac, J. E. Metall. Trans. B 1978, 9B, 431. (3) Linge, H. G. Hydrometallurgy 1976, 2, 51. (4) Kamat, P. V.; Dimitrijevic, N. M.; Fessenden, R. W. J . Phys. Chem. 1987, 91, 396. (5) Baral, S.; Fojtik, A.; Weller, H.; Henglein, A. J. Am. Chem. Soc. 1986, 108, 375. (6) Albery, W. J.; Brown, G . T.; Darwent, J. R.; Saievar-Iranizad, E. J . Chem. Soc., Faraday Trans. I 1985, 81, 1999. (7) Dimitrijevic, N . M.; Kamat, P. V. Radial. Phys. Chem. 1988,32, 53. (8) Hayes, R.; Freeman, P. A.; Mulvaney, P.; Grieser, F.; Healy, T. W.; Furlong, D. N. Ber. Bunsen-Ges. Phys. Chem. 1987, 91, 231. (9) Silvester, E. J.; Healy, T. W.; Grieser, F.; Sexton, B. A. Longmuir 1991, 7, 19. (10) Hart, E. J.; Anbar, M. The Hydrated Electron; John Wiley & Sons: New York, 1970; p 198. (1 1) Kiss, A. V.; Abraham, J.; Hegedus, I. Z . Anorg. Allg. Chem. 1940, 244, 98. (12) Balzani, V.; Carassiti, V . Photochemistry of Coordination Compounds; Academic Press: London, 1970; p 146, (13) Poulsen, I. A.; Garner, C. S. J . Am. Chem. SOC.1962, 84. 2032.
0 1992 American Chemical Society
The Journal of Physical Chemistry, Vol. 96, No. 11, 1992 4383
Oxidation Kinetics of CuFeS2 1.5 I
4 4
I
1.0
I
51
P)
0
W
m
0)
51
e 51
C
0.9
u c
f!
0
2
0.8
0.7
0.0' 300
"
"
"
"
"
400
500
600
700
800
0.6
0
Results Oxidation by Ferricyanide. The typical spectral changes which occur upon mixing chalcopyrite and ferricyanide are shown in Figure 1. The insert in Figure 1 shows the absorption spectra of transparent chalcopyrite (a) and ferricyanide (b). The kinetics of chalcopyrite oxidation were investigated over a range (lo4 to 2.5 X lo-' M) of initial ferricyanide concentrations. The sol absorbance decays at 480 nm corresponding to this range of oxidant concentrations are shown in Figure 2a. (The solid lines shown in this figure correspond to calculated decay curves based on the kinetic mechanism presented in the Discussion section.) Bleaching of the sol absorption is clearly faster at higher ferricyanide concentrations. At the higher initial ferricyanide concentrations, and at longer times than that shown in Figure 2a, absorbance bleaching at 480 nm was as great as 90%, indicating little contribution from oxidation products at this wavelength. At all but the highest initial Fe(CN)63- concentration, the reaction rate changes abruptly after bleaching of 3&40% of sol absorption (Le., at a normalized absorbance of -0.7-0.6). This is most likely the point at which surface oxidation product(s) start (14) Jagensen, C. K. Acto Chem. Scand. 1956, 10, 500.
50
40
Time ( 5 )
h (nm) Absorption spectra of a chalcopyrite-ferricyanide reaction mixture at 2-s intervals, from t = 0 to t = 16 s (initial concentrations: Fe(CN)63-at lo-' M and CuFeSl at 1.1 X lo4 M). Insert: absorption spectra of pure components, (a) chalcopyrite and (b) ferricyanide. Figure 1.
interest are very low compared to that of either IrCls2-or CuFeSz (tX,470C 30 M-l cm-').I4 Absorption by chalcopyrite oxidation products appeared to be negligible at wavelengths greater than 420 nm. The absorption that was observed below this wavelength was attributed to the oxides (or hydroxides) of copper and iron. In all systems, chalcopyrite sols were mixed in the stopped flow apparatus so as to achieve a concentration of (1.1-1.2) X M after mixing. The pH during reaction was controlled by the addition of disodium tetraborate (borax) buffer to the chalcopyrite sol at a concentration of 20 mM prior to mixing (i-e., 10 mM after mixing). The pH of reaction mixtures at this buffer concentration was in the range 9.1 f 0.1. All experiments were conducted at a temperature of 25 OC. Determination of Sulfur-Oxy Anions. Due to the small size of CuFeS,! particles, separation of the aqueous phase required the use of ultrafiltration techniques. Sulfur-oxy anions in solution were separated from the colloidal fraction by filtration of 5 mL of sol through an Amicon PM30 membrane in a 8010 stirred cell assembly. Filtrates were buffered with carbonate buffer (0.1 g L-I) and analyzed by ion chromatography (Dionex 20001/sp chromatography module with a Dionex "Ion Pac" anion-exchange column). Integrated peak areas were calibrated against the peak areas of standard sulfate and thiosulfate solutions. CIllculatiOa of ''Reaction Extent". Some of the results presented in this report are expressed in terms of 'reaction extent". This parameter refers to the percentage of chalcopyrite oxidized, as determined by the change in sol absorbance at either 470 or 650 nm.
30
20
10
0.05
b
0
10
20
50
40
30
Time (s) Figure 2. (a) Bleaching of CuFeS2 sol absorbance at 480 nm (initial absorbance normalized to 1.O) upon reaction with ferricyanideat initial concentrations of (m) lo4, (0)2.5 X lo4, (0) 5 X lo4, ( 0 ) and (m) 2.5 X lo-' M. Chalcopyrite sol initially at 1.1 X lo4 M. The solid
lines are calculated decays, based on the kinetic mechanism presented in the text. (b) Bleaching of ferricyanide absorbance at 420 nm, corresponding to CuFeSz sol decays shown in (a). Symbols and solid lines have the same meaning.
f h
r
2.0 2.0 1
l
O
.
10.0 O
IV)
I In
5 X
w
4 B
*.
5.0
1.0
0
c 0 P) c.
m
0.0
0
20
40
60
80
100
Time (s)
reaction (in M s-I) for (0) chalcopyrite and ( 0 ) ferricyanide for one reaction system studied (initial concentrations: Fe(CN),'-at 5 X lo4 M; CuFeS2at 1.1 X lo-' M). The data points are numerically calculated first derivatives of the corresponding sol and oxidant absorbance decay curves. Figure 3. Rates of
to influence the reaction kinetics. The corresponding ferricyanide absorbance decays, which are shown in Figure 2b, do not show an abrupt change in the rate of reaction at any point. In all systems, although it is more evident at intermediate initial Fe(CN)s3- concentrations, there is an induction period in the sol bleaching. A similar, but shorter, induction is observed in Fe(CN)63-decays. In Figure 3 is shown the rate of reaction, for both reacting species,for one of the systems studied. In this figure, the induction is clearly observed as a maximum in the rate of reaction at a short time after mixing. As way of a comparison, if there was no induction period, the maximum rate of reaction would occur at zero time. The point of maximum rate of decay of ferricyanide precedes that of the sol, reflecting the shorter
Silvester et al.
4384 The Journal of Physical Chemistry, Vol. 96, No. 11, 1992
TABLE I [CuFeSJ reacted, M
[Fe(CN)& reacted, M
1.26 X
1.91 X
1.26X lo4
4.59 X
1.26 X lo4
7.93 X
fibs480 nm (chalcopyrite)
loJ lo4 lo4
0.784 €480 nm(chn1copyrite)
4.6 x 10-4
101
57
90
87
[Sot-],M
(% total S)
[Fe(CN),)-] calcd, M
1.73 X
4.7 x 104 (2) 7.8 X 10” (3) 1.8 x 10-5
7.1 X
0.252 0.514
reaction extent, %
2.0 x 10-4
ferricyanide accounted for, % 103
[S202-], M (5% total S)
(14) 2.12 x 10-5 (17) = 7070 M-’cm-l
lo4
37
(7) (AbS(480)initial = 0.8g4)
1.5 v)
3 ID
s1
1.0
X
c 0
c
0
K c
0
a
0.5
# A*‘
v)
E
/* /*
0.0 0.000
I . ,
0 Q
0.001
iJ
0.002
L
oo
20
40
60
80
Reaction Extent (%) Figure 4. Rate of chalcopyrite oxidation (in M s-I) by ferricyanide at the 50% reaction point. Ferricyanide concentrations were determined from remaining oxidant absorbance at 420 nm.
Figure 5. Reaction stoichiometry of chalcopyrite oxidation by ferricyanide, as determined from reaction rates of both oxidant and sol, as described by eq 1 in the text: for initial ferricyanide concentrations of ( 0 ) 5 X lo4, and (0)lo-’ M. (0)2.5 X
induction time of this species. Also evident from this figure is the abrupt change in the rate of reaction of the sol compared to that of ferricyanide. At a reaction extent of 50%, it is conceivable that an accumulating product layer would influence the reaction kinetics due to impeded transport of reactants or products. Figure 4 shows the rate of sol oxidation for reactions which have proceeded to an extent of 508, showing clearly that the reaction rate is strongly (although not linearly) dependent on the remaining oxidant concentration. Clearly, electron transfer readily occurs across any product layer present, and while there must be accompanying chargebalancingprocesses occurring, the transport of these species does not appear to be rate determining. Decays in the absorbance of both sol and oxidant were analyzed by determination of the “time-resolved” rate of reaction for both species (as shown in Figure 3). This allowed the “time-resolved” stoichiometry of the oxidation reaction to be calculated, Le.
can be attributed to the further oxidation of So, which would be expected at a pH of 9.1,15Le. 2S0(,,+ 3H20 S2032-+ 6H+ + 4e(3)
This kinetically derived reaction stoichiometry is shown in Figure 5 for a range of initial ferricyanide concentrations. In all cases, the number of electrons removed per CuFeSz molecule is in the range 4-6 until approximately 30-35% of the reaction has occurred. The decrease in the reaction ratio to this point reflects the longer induction in the sol absorbance decays compared to that of the oxidant. A slight increase in the initial reaction stoichiometry is observed with increasing oxidant concentration, which could be attributed to more extensive oxidation. Reaction stoichiometries are not shown for the highest or lowest initial femcyanide concentrations due to considerable scatter in this data. Using an average number of five electrons, the following corrosion reaction can be postulated: CuFeS,,,,
-+
+ 5H20 Cu(OH),(,,
Fe(OH),(,,
+ 2S0,,, + 5H+ + 5e-
(2)
Beyond 30% reaction, a sharp rise in the electrons transferred per CuFeSz molecule occurs,to a value of approximately 9. This
and S2032-+ 5 H 2 0
-
2S042- + 10H+
+ 8e-
(4)
The formation of both thiosulfate and sulfate was confirmed by analysis of the filtrates of ferricyanide oxidized sols, oxidized in the absence of oxygen. On the basis of the measured thiosulfate and sulfate concentrations, and the bleaching of the sol absorbance at 480 nm, the expected consumption of oxidant was calculated, assuming that oxidation in this system proceeds via reactions 2-4, i.e.
(5) Table I shows a comparison of the calculated ferricyanide consumption with the known reacted concentration, for three systems, 30 min after mixing. The good correlation observed up to a reaction extent of nearly 60% indicates that the reactions given above do describe the oxidation of chalcopyrite by ferricyanide. In these experiments a longer reaction time was used to allow complete reaction of the oxidant, and sulfur-oxy anion desorption to occur, although it would appear that at the highat femcyanide concentration some sulfur-oxy anions remain particle bound. The formation of sulfur-oxy anions in the absence of oxygen is consistent with isotopic measurements on oxidized metal sulfide products, which show that up to 100%of sulfate oxygen can be derived from water16 even when 02(*) is available. Eff- of Added spit. The background ionic strength was raised from 0.02 M set by the borax buffer up to 0.2 M by the addition of sodium nitrate. The effect of sodium nitrate on the kinetics (15) Garrels, R. M.; Christ, C. L. Solutions, Minerals and Equilibria; Freeman, Cooper and Co.: San Francisco, 1965. (16) (a) Toran, L.; Harris, R. F. Geochim. Cosmochim. Acta 1989, 53, 2341. (b) Reedy, B. J.; Beattie, J. K.; Lowson, R.T. Geochim. Cosmochim. Acta 1991, 55, 1609.
The Journal of Physical Chemistry, Vol. 96, No. 11, 1992 4385
Oxidation Kinetics of CuFeSz
0
3 -
...
,
I
,
.
A
8
0
0.6 0
5
Time ( 5 ) Figure 6. Bleaching of CuFeS, sol absorbance at 480 nm (initial absorbance normalized to 1.0) upon reaction with ferricyanide at an initial concentration of 5 X lo4 M. Ionic strength at (0)0.02 M (lo-, M borax buffer), ( 0 )0.05 M (borax buffer plus sodium nitrate), ( 0 ) 0.1 M, and (B) 0.2 M. Sol bleaching by sodium nitrate alone at 0.2 M is also shown @).
0.5
0
10
20
Time (s) Figure 7. Bleaching of CuFeS, sol absorbance at 480 nm (initial absorbance normalized to 1.O) upon reaction with ferricyanide at an initial concentration of 5 X lo4 M; in (0) lo-, M borax buffer, ( 0 ) lom2M borax buffer 3 X M LiN03, ( 0 ) lo-, M borax buffer 3 X lo-, M NaN03, and (B) lo-, M borax buffer + 3 X M KN03.
+
+
of sol bleaching is shown in Figure 6. This shows a considerably enhanced rate of sol bleaching at higher salt concentrations. Also shown as a blank is the effect of sodium nitrate at the highest concentration used upon the sol absorbance. Clearly bleaching due to nitrate, whether by aggregation or oxidation, is minimal in this system over the time scales considered. Time-resolved reaction stoichiometries for these decays exhibited behavior identical to that shown in Figure 5 . Specific interaction effects of the cation upon the reaction rate were investigated by comparison of the effects of lithium, sodium, and potassium nitrates at equal added concentrations, as shown in Figure 7. The data show that the identity of the cation has a strong influence on the rate of reaction, indicating that enhancement of the reaction rate at higher ionic strengths is not solely due to electrostatic effects. The observed reaction rate follows the order K+ > Na+ > Li+, which follows the generally observed adsorbability of these cations onto negatively charged surfaces." Electrokinetic studies on natural chalcopyrite particles are consistent with a negatively charged particle at this pH (pH 9.1).'* This specific cation effect can be attributed to an uouter-spherebridging" phenomenon that is frequently observed in electron-transfer processes involving transition-metal complexes and in particular has been observed for electron exchange between ferri- and ferrocvanide in s o l ~ t i o n . ~In~ this model. the two
-
'
0
10
I 10
20
30
Time ( 5 ) Figure 8. Bleaching of CuFeS, sol absorbance at 480 nm (initial absorbance normalized to 1.0) upon reaction with ferricyanide at an initial concentration of 5 X lo4 M with ( 0 )no added sulfide and (0)sulfide M (25% excess in total sulfur). added at 5 X
reacting species and the cation (in this case Li+, Na+, or K+) are involved in a transition-state complex that facilitates electron transfer.z0 In the case of chalcopyrite oxidation, the transition-state complex would involve a Fe(CN)63- molecule and a negatively charged surface site, bridged by one or more monovalent cations. While there is almost certainly an ionic strength contribution to the enhanced rate of reaction due electrostatic screening effects, this is difficult to separate from the cation specific effect upon the reaction rate. Effect of Added Sulfkle. It has been shown previously that the sulfide ion has a protective action on cadmium sulfide, preventing photochemical o x i d a t i ~ n . ~This ~ - ~ protection ~ is attributed to the reaction of S2-with valence band holes generated by light absorption. The mechanism of chalcopyrite oxidation will be discussed later; however, it is appropriate at this stage to point out that if corrosion does not occur upon the removal of an electron from a CuFeS, particle, a similar protective action by an electron donor such as S2-could occur. Figure 8 shows the bleaching of a chalcopyrite sol with a 25% excess of sulfide compared to a sol without added sulfide. In the presence of excess sulfide a long induction is observed in sol bleaching which could be due to corrosion protection by this anion. Furthermore, it appears likely that a low level of excess sulfide may exist in all chalcopyrite sols, which would account for the observed shorter induction periods observed in all sol decays, even when there is no added sulfide. An excess of sulfide is not unreasonable considering both the expected negative surface charge and the potential determining role of S2- for sulfide mineralsz4 Although stopped flow investigation of the direct reaction of ferricyanide with aqueous sulfide revealed a very fast reaction (half-life C 0.5 s) between these solution species, the observed corrosion protection afforded by excess sulfide over much longer time scales would suggest that this reaction does not occur in colloidal systems. This would suggest that the bulk of the excess sulfide is surfaceadsorbed and does not react directly with solution ferricyanide at the same rate as does aqueous Sz-. Oxidation by Hexachlmiridate(W). Reactions of IrCbz- with chalcopyrite sols were much faster than for similar initial concentrations of ferricyanide. Time-resolved spectra of these reactions were therefore not obtained for this system. Absorbance measurements were recorded at 100-ms intervals, this being as fast as possible with the detection system used. Parts a and b of Figures 9 show the absorbance decays of chalcopyrite and hexto lW3 M) achloroiridate(IV),respectively, for a range ( 5 X of initial hexachloroiridate(1V)concentrations. (The solid lines (20) Cannon, R. D. Electron Transfer Reactions; Buttenvorths: London,
(17) Hunter, R. .I. Foundations of Colloid Science; Oxford University Press: New York, 1989; Vol. 1. (18) Ney, P.Zeta Potentiale und Flotierbarkeit von Mineralen; Springer-Verlag: Wien, 1973. (19) Shporer, M.;Ron, G.; Loewenstein, A.; Navon, G. Inorg. Chem. 1965, 4, 361.
1980.
(21) Henglein, A. Ber. Bunsen-Ges. Phys. Chem. 1%2, 86, 301. (22) Biihler, N.; Meier, K.;Rebcr, J. J . Phys. Chem. 1984, 88, 3261. (23) Ellis, A. B.; Kaiser, S. W.: Wrighton, M. S. J . Am. Chcm. Sot. 1976, 98, 1635. (24) Park, S. W.; Huang, C. P. J. ColloidInterfaceSci. 1W7,117, 431.
Silvester et al.
4386 The Journal of Physical Chemistry, Vol. 96, No. 11, 1992
TABLE I1 [CuFeS2] reacted, M 1.33 X lo-“
[IrC162-] reacted, M 1.75 X lo-”
1.33 X lo4
4.38
X
0.378
1.33 x 10-4
8.75 x 10-4
0.562
[S2032-l,M (% total S ) 1.7 X 10” (2) 8.8 X 10” (6) 1.8 x 10-5 (14)
AAbS650n? (chalcopyrite) 0.182
6650 nm(cha1copynk)
1.2
=
cm-’
[SO,2-1, M (% total S ) 3.6 X 10” (1) 5.9 x 10” (2) 1.9 X (7)
[ calcd, M 2.0 X lo-“
IrC1,2accounted for, % 114
reaction extent, % 27
4.29 x 10-4
98
56
7.0
80
83
X
lo4
(AbS(6SO)mit!al= 0.679)
a
1.o
