Thermodynamic calculations on sulfide flotation systems, II. Comparisons with electrochemical experiments on the galena-ethyl xanthate system

International Journal of Mineral Processing, 20 (1987) 267-290

267

Elsevier Science Publishers B.V., Amsterdam - - Printed in The Netherlands

T h e r m o d y n a m i c Calculations on Sulfide Flotation Systems, II. Comparisons with Electrochemical Experiments on the G a l e n a - E t h y l Xanthate System M.D. PRITZKER "~ and R.H. YOON

Department o[ Mining and Minerals Engineering, Virginia Polytechnic Institute and State University, Blacksburg, VA 24061 U.S.A.) (Received November 18, 1985; accepted after revision June 30, 1986)

ABSTRACT Pritzker, M.D. and Yoon, R.H., 1987. Thermodynamic calculations on sulfide flotation systems, II. Comparisons with electrochemical experiments on the galena-ethyl xanthate system. Int. J. Miner. Process,, 20" 267-290. Extensive thermodynamic calculations have been carried out to study the oxidation of galena and its flotation by potassium ethyl xanthate (KEX). In order to include some kinetic effects, three cases have been considered by assuming that sulfur oxidation can proceed as far as the formation of sulfate, thiosulfate and elemental sulfur. The results of these calculations have been compared with those of linear sweep voltammetry and intermittent galvanostatic polarization experiments conducted on galena at pH 9.2. Analysis of the experimental data indicates that in xanthate-free solutions, galena is oxidized primarily to PbO and So and, to a lesser extent, $2032-, rather than to the thermodynamically most favored PbOH + and S042-. Tests carried out in the presence of collector confirm the previous finding (Woods, 1971 ) that the interaction of xanthate with the mineral begins with chemisorption by a one-electron reaction. This cannot, of course, be predicted by calculations based on bulk thermodynamics.

INTRODUCTION

In certain cases, thermodynamics has been used successfully to interpret observed flotation behavior. For example, Walker et al. (1984) found that when a packed-bed pyrite electrode was contacted with 10-3M potassium ethyl xanthate (KEX) solution at pH 9.2, xanthate adsorption and flotation began to "1Currently at the Bureau of Geology and Mineral Technology, University of Arizona, Tucson, AZ 85721, U.S.A.

0301-7516/87/$03.50

© 1987 Elsevier Science Publishers B.V.

26~

occur only when the potential was raised above the equilibrium potential f'o~' the reaction: 2X

-,X2 +2e

lt~

However, in other cases, kinetics appears to be more important in controlling flotation. For example, Janetski et al. (1977) were able to explain the depression of pyrite by alkali in terms of the effect of pH on the reaction rates for xanthate oxidation and pyrite oxidation. One of the main objectives of the present work has been to closely examine how well thermodynamics can explain the flotation chemistry of the galena-ethyl xanthate system. Extensive computer calculations have been carried out with particular emphasis on determining the Eh-pH regions in which lead ethyl xanthate ( PbX2 and ethyl dixanthogen (X2) are stable. Since these are the species considered to induce hydrophobicity on galena, the results obtained can serve as a guide in predicting the conditions at which flotation is possible. Previously (Pritzker and Yoon, 1984), no restrictions were placed on the extent to which sulfur-bearing species could be oxidized so that the formation of sulfate ( the highest oxidation state of sulfur) was permitted. However, under the conditions typically encountered in flotation practice, sulfide oxidation is sluggish and, therefore, the metastable species such as elemental sulfur and thiosulfate can exist (Eadington and Prosser, 1969; Gardner and Woods, 1979; Guy and Trahar, 1984; Lamache et al., 1984). Consequently, similar calculations have been carried out for the two additional cases where the oxidation proceeds only as far as the formation of thiosulfate and elemental sulfur. Another objective of this paper is to present several useful ways in which the results of the calculations can be displayed. Of particular interest is the data concerning the effect of E h and pH on the mineral solubility and xanthate uptake. This will be shown in the form of two-dimensional diagrams, contour plots and three-dimensional block diagrams. Voltammetry and intermittent galvanostatic polarization (IGP)experiments have also been carried out on galena samples at pH 9.2. The results will be compared with each other and with some of the data obtained from the thermodynamic calculations. In the IGP method, a current is repeatedly applied to an electrode for a short duration and then turned off for usually an equal period of time. It was first applied by Nagel, Lange and their collaborators (1957) in their studies of equilibrium potentials for electrochemical processes occurring in metal-water binary systems. Later, Horvath and Hackl (1965) used it successfully to verify calculated E h - p H diagrams for metal-S-H20 ternary systems and to interpret the corrosion of metals in aqueous H2S environments. More recently, Thornbet (1983) used this technique and cyclic voltammetry to study the electro-

269 chemical reactions occurring at the surfaces of two nickel-iron sulfide minerals, i.e., pentlandite and violarite. EXPERIMENTAL

Materials The galena sample used in this work was a research grade ( > 99% pure) specimen obtained from Ward's Natural Science Establishment. It originated from the Brushy Creek Mine, Missouri. The collector reagent was prepared by dissolving commercial grade potassium ethyl xanthate in acetone and recrystallizing it in petroleum ether. After repeating this three times, the purified crystals were kept immersed in the ether. Prior to each experiment, a small amount of the reagent was removed and residual ether was driven off in a vacuum desiccator. Tests were conducted at pH 9.2 in 0.05 M Na2B407 solutions. All the chemicals used were of A.C.S. reagent grade and all the solutions were prepared with double-distilled water. Before each series of experiments, the 350-ml buffer solution was deoxygenated by purging it with low-oxygen nitrogen ( < 0.5 ppm 02 ) for at least one hour. Once the tests were begun, the gas sparger was raised above the solution and nitrogen was continually introduced to maintain a positive pressure.

Preparation of electrodes A galena specimen about I cm in length, with a rectangular cross-sectional area of 0.65 cm 2 was cut from a larger piece with a diamond saw. A copper lead was then fit into a small hole drilled into the back of the specimen and permanently attached with Electrodag 199, a carbon-based conducting cement. This assembly was then sealed in a glass holder with Buehler epoxy resin. The resistance of the electrode was measured at various times during its use and was always found to be between 5 and 10 ohms. Prior to each test run, the electrode surface was polished with 400- and 600-grit silicon carbide paper and then rinsed with distilled water.

Electrochemical measurements During the voltammetry experiments, a conventional three-electrode system was used. The electrode potential was controlled with a PAR Model 371 potentiostat/galvanostat and a PAR Model 175 programmer; voltammograms were recorded on a Hewlett Packard 7004B X-Y recorder. A saturated calomel electrode was used as the reference electrode, although all potentials reported here are expressed on the standard hydrogen scale.

270

For the IGP experiments, the PAR 371 unit was operated in the galvano~ static mode and only a two-electrode system was needed. By placing a resistor across the controlled current binding posts on the potentiostat, voltage pulses from the programmer were converted to current pulses. A standard 470- or 1500-ohm resistor was used, depending on the desired currents. Each curren~ pulse was applied for 1 second and then turned off for another second before being applied again. In order to monitor the change in electrode potential with time, the working electrode and reference electrode were also connected across a Keithley Model 642 electrometer. The output was recorded on a Pedersen 27MR strip chart recorder. RESULTS AND DISCUSSION

Parameters of calculations As mentioned earlier, sets of calculations were done for the following cases: Case I, oxidation of sulfide proceeds to sulfate; Case II, oxidation of sulfide proceeds only as far as the formation of thiosulfate; and Case III, oxidation of sulfide proceeds only as far as the formation of elemental sulfur. The assumptions and method for carrying out the computations were previously described in detail (Pritzker and Yoon, 1984; Pritzker, 1985). They were done for an Eh range from -- 2000 to + 2000 mV, for pH values from 0 to 14, and for xanthate additions of 0, 10 7, 10 ~, 10 -5 and 10-4M. Most of the thermodynamic data for the species considered can be found in a previous publication ( Pritzker and Yoon, 1984 ). Free energies for the species in Cases II and III that were not listed there are tabulated in Table I. W h e n sulfide (S 2 ) ions are allowed to oxidize to sulfate or thiosulfate {Case I or Case II), concentrations of polysulfide species, i.e. $22~, S:~2-, $42-~ and $52-, are negligibly small under any conditions. However, when oxidation of sulfide can proceed only to the zero oxidation state ( Case III), the presence of polysulfides can be significant. For this reason, polysulfides have been included TABLE I Free energy d a t a for a d d i t i o n a l species c o n s i d e r e d in this study AG ° (kcal/mole) "~

z~G° ( kcal/mole ) "~

$20:/~

- 127.2

S~ 2

19.8

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18.0 16.6

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- 134.0

$5 ~

15.7

"1All t h e data, except in t h e case of PbS2 O:,, were o b t a i n e d from Pourbaix, (1966). T h e free energy for PbS20:, was t a k e n from L a t i m e r (1952).

271

in the calculations for Case III. The free energy data for these species are also listed in Table I.

In the absence of xanthate

Eh-PH diagrams, contour plots and 3-D block diagrams. The Eh-pH diagrams for Cases I, II and III are shown in Figs. la, b and c, respectively. In every domain on these diagrams, the predominant lead-bearing and sulfur-bearing species (whether they are soluble or insoluble ) are specified. No xanthate was considered in these calculations. The main difference between them is the nature of the oxidation products below pH 10.1: Pb e+ and SO4e- (leading to PbSO4) are the oxidation products in Case I; Pb 2+ and SeO32- in Case II; and Pb 2+ and So in Case III. Above pH 10.1, PbO forms in all three instances. Obviously, under reducing conditions where HeS or H S - or S e- dominate, there is no difference between them. In order to more clearly see the combined effect of Eh and pH on lead solu"'-. ' ' ' ' '(a ) ",,

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bility, contour plots and three-dimensional block diagrams have been constructed using the Surface II Graphics System developed by the Kansas Geological Survey (Sampson, 1978; Fig. 2). The results for Case II are shown in Fig. 2. Since the Pb 2+ ion concentration is directly affected by which insoluble species is stable, the pattern of the contour lines (Fig. 2a) bears a resemblance to the stability domains of the Eh-pH diagram in Fig. lb. At high pH and Eh, PbO2 is stable and, consequently, the lines are determined by the Nernst equation for the Pb2+/PbO2 couple. Under reducing conditions, the lines correspond to the Pb2+/Pb ° couple and naturally are independent ofpH. The lead concentration becomes more strongly dependent upon pH and less so with respect to potential in the region between these two. A very distinct transition in the contour lines occurs as the upper boundary for the PbS domain is crossed. Although the lead concentration is independent of Eh both above and below this boundary, the solubility is considerably lower at a given Eh once PbS 1.4

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274

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becomes stable. Consequently, as the potential is lowered to the boundary line, each contour line abruptly shifts to a lower pH, where PbS becomes more soluble. Although the 3-D block diagram cannot provide quantitative information as easily as the contour plot, it does give a clearer pictorial display of the effect of Eh and pH on log [Pb 2÷ ]. Another general advantage of the block diagram over the contour plot is that considerably less computation time is required to generate it. Whereas the computer program must perform an interpolation to construct a contour line, it only has to scale the vertical dimension proportionally to the value of the dependent variable in the case of the block diagram.

Comparison with linear sweep voltammetry. Linear sweep voltammograms were obtained on a polished galena electrode in stirred and unstirred borate solutions (0.05 M, pH 9.2) containing no xanthate. For the experimental results shown in Fig. 3, the initial potential, the anodic limit and the scan speed were - 350 mV, 550 mV and 20 mV/sec, respectively. The current rise is gradual up to a potential of about 300 mV and then begins to rise more steeply thereafter. This has also been observed by Gardner and Woods (1979) and Hamilton and Woods (1984). The latter investigators have proposed that the initial oxidation of sulfides cannot be likened to any bulk reaction. Instead, it is a surface reaction involving the breakage of metal-sulfur

275

bonds and the strengthening of sulfur-sulfur and metal-oxygen bonds. Eventually, bulk oxide and sulfur nucleate, at which point the overall reaction can be expressed as: PbS+H20~PbO+S

° +2H + +2e

(2)

The shape of the current rise during the initial oxidation (i.e., below about 300 mV) was found to vary from test to test. In some instances, a small peak appeared; in others, only a plateau was observed. The randomness in this behavior is likely related to how the oxidation products form on the electrode surface. If they form uniformly on the surface, then the initial layer might inhibit further oxidation, which results in peak formation. If they grow as patches, then no distinction can be made between the completion of the first monolayer and the beginning of the next; consequently, a smoother current rise will result. Several of the characteristics noted by Gardner and Woods (1979) with regard to their study have been confirmed by our work. Firstly, stirring had no noticeable effect on the current during the anodic scan and on the charge assoI

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276

ciated with the more positive cathodic peak. This supports the conclusion or Gardner and Woods that PbO and S" are the oxidation products due to reaction 2 and that the first peak (C~ i on the cathodic scan is the reverse of this reaction, Secondly, a second cathodic peak (C=,), which coincides quite closely with the reversible potential for the reaction: P b O + 2 H + + 2 e +Pb ° + H,,()

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appears whenever galena is oxidized sufficiently during the anodic sweep. Gardner and Woods proposed the following explanation for this behavior. Galena can be oxidized beyond PbO and S ° to the formation of more PbO and oxysulfur anions, such as thiosulfate and sulfate. These reactions are irreversible so that when the scan is reversed, the excess PbO is reduced according to reaction 3. Based on the calculations for Case [II, PbO should not become stable until pH 10 is reached. Two possible reasons for the discrepancy between this and what the voltammograms indicate are that: (1) bulk free energy data have been used for these surface reactions, and ( 2 ) there may be a considerable difference between the surface concentration and the bulk concentrations of the reactants and/or products. One of the assumptions of these calculations is that mass transfer is so rapid that surface concentrations can be equated to bulk concentrations. Comparison of the voltammograms in Fig. 3 indicates that stirring diminishes the size of the C~ peak significantly, although it has very little effect on the anodic current, A~. This suggests that some of the PbO generated during the anodic scan is being dissolved when the solution is flowing. Apparently, the system is evolving toward the thermodynamically most favored state, i.e., soluble lead hydroxy complexes, hut mass transfer limitations are retarding the kinetics of this process. Analysis of the charges associated with the peaks in these voltammograms can be used to determine whether oxy-sulfur species are being produced and, if so, which one is predominant. Since reaction 2 is essentially reversible, the charge passed during the anodic sweep ( QA, ) should be matched by the charge associated with the first cathodic peak ( Qc, ), provided that PbO (or PbOH * ) and S o are the only oxidation products. However, of oxidation proceeds beyond the formation of S °, then QA, will exceed Qc,. The most likely oxy-sulfur species to be generated a r e S 2 O a 2 and 8042 according to the reactions: 2PbS+5H,~O~2PbO+S~O~

+ I O H ~ +8e

(4)

and P b S + 5 H 2 0 - ~ P b O + S O 4 ~ + I O H ~ +8e

(5)

If only thiosulfate is being formed, it can be shown that the charges associated

277

with the anodic sweep and the two cathodic peaks are related by the expression Q A , - Qc, = 2Qc~, where Qc2 is the charge corresponding to reaction 3. The relation becomes QA1 - Qcl = 4Qc2 when oxidation proceeds by reaction 5. In order to determine what species are being formed and the potential at which they begin to appear, scans have been performed at a constant sweep rate (20 mV/sec), but with varying anodic limits. These are shown in Fig. 4. The oxidation of galena below 200 mV was not extensive enough for any definite conclusions to be made. The values of Qc,/QA, and (QA, - Qcl)/Qc2 for potentials beyond this are given in Table II. The PbO/Pb ° peak does not become distinct enough for a reliable value of Qc2 to be obtained until the anodic limit is extended to 445 mV. Nevertheless, the data indicates that thiosulfate and not sulfate is being generated in addition to elemental sulfur and that it is present at potentials as low as 195 mV. Based on this, the molar ratios of $2032to S Ohave been calculated and are also included in Table II. Up to about 445 mV, elemental sulfur predominates and the relative amount of the two remains essentially constant. However, when the potential is raised above this, the amount of thiosulfate being produced increases rapidly. An argument can be made that not all the S Othat is produced is removed during

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278 T A B L E II D e t e r m i n a t i o n o f o x i d a t i o n p r o d u c t s on g a l e n a as a f u n c t i o n of a n o d i c limit A n o d i c limit ( m V )

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the cathodic sweep, particularly if the oxidation has been considerable. A value of Qcl/QA~ less than unity will, therefore, not necessarily mean that an oxysulfur anion is being formed. If this were so, one would expect the ratio to decrease as the anodic limit is extended since this tends to make the oxidation product more difficult to remove. However, the results in Table II indicate that until 445 mV is reached, the ratio does not change. Furthermore, one might also expect that varying the scan rate would affect the ratio if the removal of S Ois a problem, since this would alter the time allowed for reduction. Analysis of our data shows, however, that all the Qc,/QA, values are between 0.58 and 0.68 for a series of experiments in which the scan rate varies from 5 to 60 mV/sec and the anodic limit is held at 345 mV. These results support the conclusion that thiosulfate is indeed being formed. In addition, it is difficult to come to another conclusion considering that the ratios (QA,-Qc~)/Qc~ in Table II are so close to the ideal value of 2. Gardner and Woods (1979) and Lamache et al. (1984) have also reported detecting $20:32- along with S o as oxidation products.

Comparison with IGP. W h e n a current pulse is applied, the electrode potential is affected in two ways. The first is a rapid, oscillatory change as an immediate response to the pulses, whereas the second is a continual shift in the direction of the polarization. The result is that the upper and lower limits of the oscillations will continually move in the direction of the polarization. W h e n a faradaic process occurs and is able to keep up with the applied current, the potential will still oscillate, but now within fixed limits. This will appear as a plateau on a potential-time diagram. Typical I G P traces are shown in Figs. 5 and 6. The potential corresponding to maximum polarization for each point in time along the trace corresponds to the closed-circuit potential, while the minimum is the open-circuit potential. Current pulses of 0.05 m A / c m 2 were applied to a polished galena electrode in an unstirred borate solution for varying lengths of time and then were

279

reversed in order to reduce the oxidation products. The results are shown in Fig. 5. When galena is reduced without first oxidizing (Fig. 5a), the opencircuit potential quickly drops until it reaches a plateau at -720 mV. This is attributed to the reduction of galena according to the reaction: PbS+H + +2e~Pb ° +HS-

(6)

which has a Nernst potential of - 6 6 0 mv for [ H S - ] =10 -6 M that is very close to the observed open-circuit value. When the freshly polished galena electrode is oxidized, a well-defined plateau appears at an open-circuit potential just below 250 mV (Fig. 5b). The open-circuit potential along the plateau increases slowly with time, reaching approximately 300 mV after 6 minutes of polarization (Fig. 5d). This range

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Fig. 5. IGP diagrams for a polished galena electrode in an unstirred solution of 0.05 M Na2B407 solution in the absence of xanthate. Initial anodic polarization was carried out for 0 (a), 2 (b), 3 (c) and 6 (d) minutes. Current density=0.05 m A / c m 2. A, C2

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280 coincides well with the Nernst potential/br reaction 2 at this pH. It should be noted that at the applied current density of 0.05 mA/cm ~, we may be dealing with bulk oxidation of PbS to PbO and S Oby the time the plateau appears. During the first few seconds of' polarization, however, oxidation may involve the break-up of' the galena structure just at the mineral surface. This may be correspond to the current rise during the anodic sweep to 245 mV that can be seen in the voltammogram in Fig. 9. When the current is reversed after about 110 seconds of anodic polarization (Fig. 5b), an arrest (C1) forms at approximately - 1 0 0 mV, which can presumably be assigned to the reverse of reaction 2. Extending the duration of anodic polarization to 3 and 6 min (Figs. 5c and d) causes the PbO,S°/PbS arrest to broaden. In addition, another plateau ( C2 ) appears at - 4 0 0 mV and becomes more distinct as the extent of oxidation increases. This is consistent with what was observed with respect to the more negative cathodic peak in the voltammograms. Thus, it is likely that this plateau is due to reaction 3. The appearance of the PbO/Pb ° arrest indicates that thiosulfate is also being produced during oxidation. Comparison of the IGP diagrams in Fig. 5 shows that as the oxidation of galena proceeds, less time (and hence charge) is required to remove the oxidation products than it takes to produce them. This is also what is expected if an oxy-sulfur species, such as thiosulfate, is being formed. Although these results do suggest that reaction 4, as well as reaction 2, is occurring, no distinct arrest corresponding to this appears in any of the potential-time diagrams. This may be explained as follows. When anodic current is flowing, oxidation by both reaction 2 and reaction 4 can take place; however, when the circuit is opened, the potential can only be controlled by the reaction which can proceed in both directions. Since the oxidation of galena to PbO and $20:~ 2 is irreversible, the open-circuit potential will correspond to the Nernst potential of reaction 2 only. When the current pulse is increased to 0.5 mA/cm 2, the resulting traces change dramatically (Fig. 6). During oxidation, there is a very large fluctuation in potential fbr the first 2.5 minutes as the circuit is opened and closed. Presumably, the PbO layer is growing thick enough to passivate the electrode surface. Eventually, the current polarizes the electrode to a high enough potential that a new oxidation reaction (A4) begins to occur at an open-circuit potential of about 800 inV. With the onset of the new reaction, the difference between the closed- and open-circuit potentials during each pulse cycle diminishes. This change in behavior suggests that an oxidation product which is a good conductor is now being formed. The most probable species would be lead dioxide. PbO2 is an ntype semi-conductor with a band gap of 1.5 eV and a high conductivity, equivalent to that of the metal bismuth (Hampson, 1979). When the current is reversed, a small cathodic plateau appears at a potential

281 I

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I 12 (rain)

Fig. 6. IGP diagram for a polished galena electrode in an unstirred solution of 0.05 M Na2B407in the absence of xanthate. Current density = 0.5 mA/cm2. A~ = PbS/PbO,S°+PbS/PbO,S20~ 2A4 = PbS/Pb02,S2032- or PbO/Pb02 C3 = PbS/Pb°,HS-

C1 = PbO,S°/PbSorS°/HSC2 = PbO/Pb° C4 = PbO2/PbO

of about 650 mV. Carr and H a m p s o n (1971) conducted voltammetry studies on electrodeposited P b 0 2 in alkaline solutions and observed peaks for the reduction of Pb02 to PbO at a potential of about 500 mV. It is possible that the small cathodic plateau is due to the same reaction. With further cathodic polarization, the potential drops to where the reverse of reactions 2 and 3 can occur. The reduction of P b O to Pb ° is now predominant, indicating that oxidation is, for the most part, proceeding beyond the formation of elemental sulfur to that of oxy-sulfur species. It is therefore apparent from the electrochemical experiments that PbO, S O and $2032- are the primary oxidation products formed on galena at p H 9.2. Although $2032- is predicted to be the principal sulfur-bearing oxidation product by the calculations for Case II and S Oby those of Case III, it is impossible for the calculations to show both to be present simultaneously on the basis of equilibrium free energy data. It is clear, however, that under moderately oxidizing conditions, S Ois the predominant oxidation product.

In the presence of xanthate Eh-pH diagrams. With the K E X addition as low as 10 - 7 M, PbX2 forms only in Case II. The reaction in which PbXe is formed may be represented as follows:

282

2PbS+4X

+3H20--~2PbX2+S20:~ ~ +6H ~ +8e

tT

Although not shown in this paper, the region in which PbX2 is stable extends from pH 1.2 to pH 7.2 and from 60 to 350 mV. Beyond this upper potential limit, PbX2 decomposes by the reaction: PbX2--~ Pb 2" + X2 ( aq ) ± 2e

(8 )

Liquid dixanthogen, X2 (1), which may also render galena floatable if it adsorbs, cannot form at this low collector concentration since its solubility limit is about 1.3" 10 -s M. When the KEX concentration is increased to 10 -6 M, PbX2 can exist in all three cases considered, as shown in Fig. 7. In Case I (Fig. 7a), PbX2 is stable up to pH 8.4. The solid phase that can co-exist with PbX2 is primarily PbSO4, but PbS can also co-exist with it over a narrow potential range. The

2.0 1

II

I

I

I

I

l

I

I

1.6~

!

10 - 6 M

HS04

"

'

(a)

KEX

4

S04

PbS04+ Pb304

0.8 PbS04

U3 l-_J 0

0.4 PbX z + PbS04. PbXz+ PbS

t-

uJ

-0.4

~s +"Pb(o) ~ " ~

PbS~" Pb I

-0.8

I I

plb HzS

-1.2

I I

I

HS-

I I

-I.6

l I

1

-2.0 0

I 2

I 4

I 6

i

I

I 8

I I0

I 12

14

pH Fig. 7. Eh-pH diagrams of the PbS-H20-ethyl xanthate system at a collector addition of 10 6 M at T=298°K and P = I atm. for Case I (a), Case II (b) and Case III (c). The domains of the soluble sulfur-bearing species are delineated by ; separates the domains for the soluble lead-bearing species.

0

0

I

I

II

P b 2~"

2

I

I

I 4

HzS

PbX2+ PbS+ S

~"'1

-

II

Fig. 7 ( c o n t i n u e d ) .

-2.0

-I.6

-I.2 f

-0.8 -

u.l -0.4

o

'1

if)

04

0.8

1.2

1.6

2.0

S2032"

PbOz

I

i 8

i

6

pH

Pb

IPbS + Pb

PbS I

I

I0

I

HS-

I0 -6 M

I

(b)

12

I

KEX

!

14

tLd

o

O~ I-_.1

I

-2'00

-I.6

-I.2

-0.8

- 0.4

o

0.4

0.8

1.2

1.6,

20 I

I

2

/

4

I

Hzs

StPbS + PbX;

S + Pb X 2

S Pb Z÷

I

6

pH

I

I I I I

I I I

S+PbS I

8

I

PbS 4- Pb

I , s

I

_

i

I0

I

.s-

I0 -6 M

I

,

I I

I

PbO z + S 'l

,

I

s, < s;

i2

SsZ"

KEX

(c)

i

14

1

OO Co

284 upper Eh limit is 270 mV, and the lower Eh limit varies with p H / b l l o w i n g the P b S / P b S O 4 phase boundary. In Case II (Fig. 7b), PbX~ is fbund to be stable in a higher Eh region than in Case I ( up to 400 mV ), and the upper pH limit is 8.8. At higher Eh, PbXe is the only stable solid phase, but: at lower Eh, it can co-exist with PbS. It: (:an coexist with S Oonly at very low pH. When oxidation of sulfide is limited to the fbrmation of S ° ( Case III ), PbX2 is fbund to be stable over a narrower Eh region, as shown in Fig. 7c. The solid phases that can co-exist with PbX2 are PbS and S °. The reaction that forms PbX,~ in Case II remains the same as at 10 7 M K E X (eq. 2), while those for Cases I and III are: PbS+2X

+4H~O--~PbX:~-~SO4 ~ + 8 H ÷ +Be

(9)

and: PbS + 2X-- ~

PbX2 + S O+ 2e

0.6

(10)

I

(a) %

0.4t~

PbX 2

10-5M KEX C a s e 1[

\

0.2 I-. ..J o >

1

% o

i -0.2 -0.4

-0"60

I

2

I

4

I 10

I

6

pH

I 12

14

285

(b) Case

?

o-

Fig. 8. Effect of Eh and pH on % PbX2 at a 10 ~M KEX addition for Case II in the form of a contour plot (a) and a 3-D block diagram (b). The number next to each contour line corresponds to the value of % PbX2. Although not included in this paper, an increase in the xanthate level to 10 -5 and 10 -4 M enlarges all of the domains in which PbX2 is found. In each of the three cases, the upper pH limit for PbX2 stability rises to just above 10 at 10- ~ M K E X and 11.1 at 10 -4 M KEX. A xanthate concentration of 10 -4 M is also high enough t h a t X2(1) becomes stable as a result of the oxidation of lead xanthate and free xanthate ions.

Effect of Eh and pH on PbX2 formation. In our calculations, insoluble phases are also included in the mass balances and, thus, the number of moles of the stable solids can be directly determined. T h i s can be used to extract other information t h a t is useful for comparison with experimental adsorption or flotation data, e.g., the effect of Eh and pH on the fraction of added xanthate t h a t ends up as PbX2 (hereupon referred to as % PbX2). The combined effect of Eh and pH on % PbX2 for Case II at 10 -5 M K E X is shown in the contour plot and block diagram in Fig. 8. There appear to be two distinct regions for the formation of PbX2. The first one lies between pH 0.5 and 2.8 at potentials from - 300 to 200 mV, and the second occurs under more

286

oxidizing conditions (as high as about 400 mV ) and up to a p H of just beyond 10. Whereas % PbX2 reaches a maximum of only slightly above 40(~';, in the first region, it reaches virtually" 100% over a large portion of the second one. The upper limit at about pH 10 is determined by the competition between xanthate and hydroxyl anions for lead and is independent of the potential. Hydrolysis does not readily occur at low and neutral pH and, consequently, PbX2 can form. However, when the p H is increased above a value of about 10, it becomes dominant and prevents PbX2 from being stable.

Comparison with linear sweep voltammetry. Voltammograms at a sweep rate of 20 mV/sec were obtained for the galena electrode in unstirred 10 -3 and 10-4M K E X solutions. These are superimposed in Fig. 9, with one obtained in the absence of collector. The voltammogram for 10 -3 M K E X is similar to that reported by Woods (1971) for a collector addition of 9.5-10 ~M. During oxidation, a very distinct pre-wave or peak is always observed prior to the steep rise regardless of the scan rate. This is not always found to be so for 10-4 M KEX. At the higher K E X concentration, the more positive cathodic peak shifts toward lower potential.

zo

I

~

-----

NO

•

E

i

!

!

KEX 10-4M KEX

..............I0 . -3 M

=k

--, >-

,

~ /

/!~ .ff

KEX

pH9.2

I0

,

s..,~I

boo Z LU tm 0 Z LU ,Sf ~

n~ 0

-I0

'"

1 I'"

,.";

-20

(~ -6 0

I

I

I

-400 POTENTIAL

I

I

I

-200

I

0 (mV

vs.

I 200

A

SHE)

Fig. 9. Voltmnmograms of a polished galena electrode in unstirred solutions of 0.05 M Na2B407 containing 0, 10 4and 10-aM KEX. Sweep rate = 20 mV/see.

287

When chemisorption involves electron transfer, the amount of charge passed for the completion of monolayer coverage should be independent of scan rate. To determine whether xanthate chemisorption is responsible for the anodic pre-wave at 10 -3 M KEX, a series of scans at sweep rates varying from 10 to 50 mV/sec were obtained (Fig. 10). The charge associated with the pre-wave was measured by integrating the voltammograms up to the point of inflection between the pre-wave and the steep current rise. The results in Table III show that the charge passed is independent of scan rate and averages about 110 llC/cm 2. If a xanthate ion is considered to have a cross-sectional area of 21 A2, then adsorption by a one-electron process would require 74 pC/cm 2 for monolayer coverage. Thus, the charge density based on the apparent surface area is 1.5 times this value. Considering the relatively coarse polishing of the electrode before each test, a roughness factor of 1.5 is quite reasonable. Accepting this, then the data is consistent with the chemisorption reaction: X-

(11)

'Xads q-e

occurring until a potential of about 200 mV is reached. Woods (1971) came to the same conclusion on the basis of his study. Beyond this, the current rises much more steeply, presumably as multilayers of xanthate-bearing species form I

I

I

I

I

pH 9.2

4O

'

10-3M K E X

r

. J

no stirring

E

<~ 20 ::L >.p-

- -

I0 rnv/sec

......

is mv/sec

.......... .........

20 mv/sec 25 mv/sec

.J::"

.~'~ ~ ."--' ~~--/3/:/-~

:.~:~ ~, :~'~

Z

w c~

0

I-z uJ

tY

,.!i/

-2C

:: .....

// ./i .," //' . "

-4(

I.

-600

I

-400

I

-200

POTENTIAL

I

0

I

200

(my, SHE)

Fig. I0. Effect of scan rate on the single sweep voltammograms for a galena electrode immersed in a 0.05 M Na2B4Ov (pH 9.2 ) + 10 -'~M K E X solution. N o stirring.

28S 'FABLE III Effect of scan rate on the charge density of' the anodic pre-wave Scan rate ( mV/sec )

Charge density of anodic pre-wave (/zC/cm ~

10 15 20 25 50

112 114 109 107 107

on the surface due to reactions 1, 7 and/or10. The cathodic peaks in Fig. 10 that appear at potentials of about - 300 mV are associated with the removal of the xanthate-bearing species. The effect of scan rate on the anodic sweep for 10 4 M K E X has also been investigated. However, in this case, the charge passed decreases as the scan rate is increased. From these data alone, it is difficult to identify the reaction mechanism, but certainly the reaction is slower than at 10--'3 M KEX. Since xanthate chemisorption is a surface reaction, it is not accounted for in the thermodynamic calculations of this study. However, this type of reaction can, in principle, be included if the relation between the free energy of adsorption and surface coverage is known. It would also require that the surface area of the mineral be specified. S U M M A R Y AND C O N C L U S I O N S

The important results of this study can be summarized as follows. (1) Thermodynamic calculations on the P b S - K E X system have been carried out for three separate cases where the oxidation of sulfide is considered to proceed as far as the formation of sulfate, thiosulfate and elemental sulfur, in order to obtain information on both stable and metastable species. The calculations show that as the extent of sulfide oxidation progresses from elemental sulfur to thiosulfate and to sulfate, the lower potential edge for PbX2 stability becomes more cathodic. ( 2 ) The results of voltammetry and IGP experiments carried out at pH 9.2 in the absence of collector indicate that PbO and S Oare the primary oxidation products at lower applied potentials and currents. This is in partial agreement with the results of the calculations for Case III, in which soluble lead hydroxyl complexes and S o are found to be the predominant oxidation products under similar conditions. Stirring of the solution tends to dissolve away some of the PbO, thereby moving the system toward the more thermodynamically favored state. Although small amounts of thiosulfate are also detected at lower potentials, it is only above about 500 mV that it begins to appear in significant amounts.

289

( 3 ) The voltammograms obtained at 10 - 3M KEX addition show a pre-wave beginning at approximately - 100 mV, and suggest that xanthate chemisorbs on PbS via a one-electron transfer process. Once a monolayer is complete, oxidation may lead to the formation of X2 or PbX2. ACKNOWLEDGEMENTS

The authors express their gratitude to Dr. Ron Woods for his valuable discussions and to Beth Dillinger for reading and typing the manuscript. They also acknowledge the financial support of the National Science Foundation ( Project No. CPE-8303860) and the Virginia Mining and Minerals Resources Research Institute.

REFERENCES Cart, J.P. and Hampson, N.A., 1971. Differential capacitance and linear sweep voltammetry studies on polycrystalline lead and electrodeposited lead dioxide. J. Electrochem. Soc., 118: 1262-1268. Eadington, P. and Prosser, A.P., 1969. Oxidation of lead sulphide in aqueous suspensions. Trans. IMM, 78 C74. Gardner, J.R. and Woods, R., 1979. A study of the surface oxidation of galena using cyclic voltammetry. J. Electroanal. Chem., 100: 447. Guy, P.J. and Trahar, W.J., 1984. The influence of grinding and flotation environments on the laboratory batch flotation of galena. Int. J. Miner. Process., 12: 15-38. Hamilton, I.C. and Woods, R., 1984. A voltammetric study of the surface oxidation of sulfide minerals. In: P.E. Richardson, S. Srinivassan and R. Woods (Editors), Proceedings International Symposium on Electrochemistry in Mineral and Metal Processing. The Electrochemical Society, Pennington, N.J., 84-10: 259-285. Hampson, N.A., 1979. The aqueous system Pb 4+/Pb2+: Electrochemical aspects. In: A.T. Kuhn (Editor), The Electrochemistry of Lead. Academic Press, New York, N.Y., pp. 29-63. Horvath, J. and Hackl, L., 1965. Check of the potential/pH equilibrium diagrams of different metal-sulphur-water ternary systems by intermittent galvanostatic polarization method. Corros. Sci., 5: 525. Janetski, M.D., Woodburn, S.I. and Woods, R., 1977. An electrochemical investigation of pyrite flotation and depression. Int. J. Miner. Process., 4: 227. Lamache, M., Lam, O. and Bauer, D., 1984. Electrochemical oxidation of galena and ethyl xanthate. In: P.E. Richardson, S. Srinivasan and R. Woods (Editors), Proceedings, International Symposium on Electrochemistry in Mineral and Metal Processing. The Electrochemical Society, Pennington, N.J., 84-10: 54-65. Latimer, W., 1952. Oxidation Potentials. Prentice-Hall, Englewood Cliffs, N.J. Nagel, K., Ohse, R. and Lange, E., 1957. Z. Elektrochem., 61: 795. Pourbaix, M., 1966. Atlas of Electrochemical Equilibria. Gauthiers-Villas, Paris. Pritzker, M.D., 1985. Thermodynamic and Kinetic Studies of Galena in the Presence and Absence of Potassium Ethyl Xanthate. Ph.D. dissertation, Virginia Polytechnic Institute and State University, Blacksburg, VA. Pritzker, M.D. and Yoon, R.H., 1984. Thermodynamic calculations on sulfide flotation systems,

290 I. Galena-ethyl xanthate system in the absence of metastable species. Int. J. Miner. Process.. 12: 95-125. Sampson, R., 1978. Surface II Graphics System. Lawrence, Kans. Thornber. M.R., 1983. Mineralogical and electrochemical stability of the nickel-iron sulphides-pentlandite and violarite. J. Appl. Electrochem., 13: 253. Walker, G.W., Walters, C.P. and Richardson, P.E., 1984. Correlation of the electrosorption or' sulfur and thiol collectors with contact angle and flotation. Presented at the 165th Meeting of The Electrochemical Society, Cincinnati, Ohio. Woods, R., 1971. The oxidation of ethyl xanthate on platinum, gold, copper and galena electrodes. Relation to the mechanism of mineral flotation. J. Phys. Chem., 75: 354.