Ind. Eng. Chem. Re's.1990,29, 1178-1183
1178
Goldwasser, J.; Engelhardt, J.; Hall, W. K. The Isomerization and Metathesis of n-Butenes. I. Unreduced Molybdena-Alumina Catalysts. J. Catal. 1981a, 70,275-286. Goldwasser, J.; Engelhardt, J.; Hall, W. K. The Isomerization and Metathesis of n-Butenes. 11. Acid Catalysis on the Fully Oxidized Catalyst. J . Catal. 1981b,71,381-388. Haag, W. 0.;Pines, H. The Kinetics of Carbanion-Catalized Isomerization of Butenes and 1-Pentane. J. A m . Chem. SOC.1960,82, 387-391. Hisatsune, I. C Isomerization of Butenes by Atomic Nickel and Magnesium. J . Catal. 1982,75,425-427. Hoser, H.; Krzyzanowski, S. Studies on Isomerization of n-Butenes and 3,3-Dimethyl-l-Butene over Cobalt-Exchanged Zeolite X. J . &tal. 1975,38,366-374. Itoh, H.; Tada, A.; Hattori, H. Isomerization of Butenes over Aluminum Phosphate Catalyst. J . Catal. 1982,76,235-237. Kladnig, W. F. Use of Hammett Indicators for Acidity Measurements in Zeolites. J. Phys. Chem. 1979,83,765-766. Langner, B. E. Reactions of Olefins on Zeolites: The Change of the Product Distribution with Time on Stream in the Reaction of Butene-1 on Calcined NaNH,-Y. J . Catal. 1980,65,416-427. Lee, H. C.; Butt, J. B. Kinetics of the Desulfuration of Thiophene: Reactions of Thiophene and Butene. J. Catal. 1977,49,320-331. Manton, M. R. S.; Davidtz, J. C. Controlled Pore Sizes and Acive Site Spacings Determining Selectivity in Amorphous Silica-Alumina Catalysts. J . Catal. 1979,60,156-166. Martin, G. E.; Hill, L. W. Carbon Black-Catalyzed Olefin Isomerization. 11. Kinetics of a Three-Membered, First-Order, Reversible System. J . Catal. 1976,42,344-349. Mieville, R. L.; Meyers, B. L. Measuring Acidity by TemperatureProgrammed Desorption. J. Catal. 1982,74,196-198. Mintsa-Eya, V.;Hilaire, L.; Toroude, R.; Gault, F. G.; Moraweck, B.; Renouprez, A. Exchange and Isomerization of But-1-ene on S u p ported Pd, Ni, and Pd-Ni Alloys. J . Catal. 1982, 76, 169-181. Nagy, J. B.; Guelton, M.; Derouane, E. G. 13C NMR Investigation of the Isomerization of 1-Butene on a Synthetic Near-Faujasite Germanium Zeolite. J. Catal. 1978,55,43-51. Nakano, Y.; Iizuka, T.; Hattori, H.; Tanabe, K. Surface Properties of Zirconium Oxide and its Catalytic Activity for Isomerization of 1-Butene. J. Catal. 1979,57,1-10. Ragaini, V.The Isomerization of n-Butenes over Platinum Black in the Absence of Molecular Hydrogen. J . Catal. 1974, 34, 1-6.
Rodenas, E.; Yamaguchi, T.; Hattori, H.; Tanabe, K. Surface and Catalytic Properties of Ti02-A1203. J . Catal. 1981,69,434-444. Romero, A.; Bilbao, J.; Aguayo, A. T. Calcinacibn y Pastillado de Catalizadores de Silice-Alumina. An. Quim. 1983,79,393-398. Rosynek, M. P.; Fox, J. S. Characterization of Catalytic Lanthanum Oxide for Double Bond Isomerization of n-Butenes. J. Catal. 1977,49,285-293. Rosynek, M. P.; Fox, J. S.; Jensen, J. L. Kinetics and Mechanism of n-Butene Isomerization over Activated Lanthanum Oxide. J . Catal. 1981,71, 64-77. Sakai, Y.; Hattori, H. The Nature of Active Sites on Aluminum Phosphate for the Isomerization of Butenes. J . Catal. 1976,42, 37-47. Schwarz, J. A,; Russell, B. G.; Harnsberger, H. F. A Study of Pyridine Adsorbed on Silica-Alumina Catalysts by Combined Infrared Spectroscopy and Temperature-Programmed Desorption. J. Catal. 1978,54,303-317. Suzuki, I. On the Order of Catalytic Isomerization. J . Catal. 1981, 68,220-221. Takahashi, M.; Iwasawa, Y.; Ogasawara, S. The Nature of Adsorbed Sites on Catalysts. 11. Behavior of Basic Compounds on SilicaAlumina Catalyst at Elevated Temperatures. J. Catal. 1976,45, 15-24. Tanabe, K. Solid Acids and Bases; Academic Press: New York, 1970. Uematsu, T.; Tsukada, K.; Hashimoto, H. The Isomerization of 1Butene Over Cation-Exchanged Acidic Resin. J. Catal. 1974,32, 369-375. Uematsu, T.; Inamura, K.; Hirai, K.; Hashimoto, H. The Isomerization of Butenes over Doped Zinc Oxides. J . Catal. 1976,45, 68-76. Van Roosmalen, A. J.; Mol, J. C. Active Centers for the Metathesis and Isomerization of Alkenes on Tungsten4xide/Silica Catalysts. J . Catal. 1982, 78,17-23. Wei, J.; Prater, C. D. The Structure and Analysis of Complex Reaction Systems. In Advances i n Catalysis; Eley, D. D., Selwood, P. W.; Weisz, P. B., Eds.; Academic Press: New York, 1962;Vol. 13,p 203. Receiued for review May 15, 1989 Revised manuscript received January 18, 1990 Accepted February 21, 1990
Rate Law and Mechanism for the Oxidation of Copper(1) by Iron(II1) in Hydrochloric Acid Solutions R.J. Ortht and K. C. Liddell* D e p a r t m e n t of Chemical Engineering, W a s h i n g t o n S t a t e Uniuersity, P u l l m a n , W a s h i n g t o n 99164
The kinetics of the oxidation of Cu(1) by Fe(II1) in HCl solutions were studied electrochemically. A rotating ring disk electrode was used, and rigorous corrections were made to the ring current to account for overlap of the Fe(II)/Fe(III) and Cu(I)/Cu(II) couples. The rate was first order in both [Cu(I)] and (Fe(III)] and also depended on [Cl-1, and inversely on [H+]. A mechanism with two parallel reaction paths is proposed to account for these observations. The activation energy of the reaction is 86.4 kJ/mol. Understanding the rates and mechanisms of the reactions of copper species in aqueous chloride solutions is of growing practical importance as hydrometallurgicalroutes for copper production receive increased attention. In particular, several leaching processes have been developed that produce copper(1) from copper sulfide ores (e.g., Peters et al., 1981; Schweitzer and Livingston, 1982); most such processes have not yet been successfully operated
* To whom correspondence
should be addressed. Present address: Unocal Science and Technology, Unocal Corp., 376 S. Valencia Ave, P.O. Box 76,Brea, CA 92621. 0888-5885/90/2629-1178$02.50/0
for extended periods, however. One of the reactions needing study is the redox reaction between copper(1) and iron(II1); iron is virtually ubiquitous in sulfide ores and often occurs in the trivalent state. For example, in CuFeS2, the most abundant copper ore mineral, several lines of evidence indicate that copper is univalent and iron trivalent (Nakai et al., 1978; Abdullin et al., 1987). A detailed study of the copper(I)/iron(III) reaction has been reported by Parker and Espenson (1969),but a perchlorate medium was used. This provides a valuable benchmark for our electrochemical study, although the perchlorate system is chemically simpler since the Clod0 1990 American Chemical Society
Ind. Eng. Chem. Res., Vol. 29, No. 7, 1990 1179 anion, unlike C1-, does not form strong complexes with either Fe or Cu. Thus, our newly developed approach for the study of redox reactions (Orth, 1988,Orth et al., 1989) was tested by comparing our conclusions to those of the earlier work and was also used to investigate the role of complexation in Cu(I)/Fe(III) systems, thereby extending the earlier investigation. The analytical technique involves using the rotating ring disk electrode for determining the kinetics for the reaction between two components having overlapping currents. The technique was applied to the reaction between Cu(1) and Fe(I11) in 1 mol dms HC1. The experiments were carried out by applying constant disk currents that were less than the limiting current and conducting the experiments at different rotation speeds. Cu(I1) and Fe(II1) were initially present in solution; the disk was used to reduce the Cu(I1) to Cu(I), and then the Cu(1) reacted with the Fe(II1) as it moved radially across the insulated gap of the rotating ring disk electrode to the ring electrode. The extent of reaction was monitored by back-calculating the Fe(III)/Fe(II) current contribution at the ring electrode and measuring the ring current due to Cu(1) at 0.4 V. The kinetics of the reaction were determined by combining the experimental results with the solutions to the equations describing mass transport with chemical reaction for the rotating ring disk electrode. These equations were solved numerically by orthogonal collocation. By use of this method, numerous points on a ring current versus disk current working curve were used for a given rotation speed to elucidate the kinetics for the reaction between Cu(1) and Fe(II1) in 1 mol dm-3 HCl at 23 OC. A value of (8.5 & 2.1) X lo6 dm3/(mol-s)was determined for the second-order rate constant for the reaction of these two components under these conditions.
Methods Electrochemical measurements were carried out by using a Pt-Pt rotating ring disk electrode (RRDE) in a fivenecked flask. Following careful purging to remove oxygen from the solution, solid FeC13*6H20and CuC12.2H20were added in appropriate amounts and allowed to dissolve. Copper(1) was produced in situ by reduction of Cu(I1) at the disk electrode. Once formed, Cu(1) was simultaneously oxidized by Fe(1II) and transported radially by convection. As discussed in detail in previous publications, for all potentials of interest, the measured ring current results from the overlap of the Cu(II)/Cu(I) and Fe(III)/Fe(II) currents (Orth, 1988; Orth et al., 1989). A procedure involving the preparation of calibration curves for the individual redox couples, and orthogonal collocation simulations of the net ring current, was developed to overcome this problem. A detailed error analysis (Orth, 1988) indicates that the worst case error is about 20% ; significant error was avoided by taking the inequality in the diffusivities of Cu(1) and Fe(II1) (10.1 X lo4 and 4.6 X lo* cm2/s, respectively, in 1.0 mol/dm3 HCl) into account in the simulations. The range of concentrations investigated was 0.05-1.0 mol/dm3 H+, 0.08-1.03 mol/dm3 total C1-, and 0.22-0.44 mmol/dm3 Fe(II1). The bulk Cu(I1) concentration used was 16.4 mmol/dm3, and the disk current was selected to reduce a desired fraction of Cu(I1) to Cu(1). Electrode rotation speeds between 400 and 4000 rpm were used. As discussed in previous works (Orth, 1988; Orth et al., 1989), these concentrations and rotation speeds ensured that the Cu(1) and Fe(II1) concentrations were comparable in the reaction zone between the disk and ring and gave the best sensitivity and accuracy. Due to the thermal properties of Teflon, which was used as the insulating gap material
Table I. Hydrogen Ion Effect on k medium HC1, KCl, [H+], [Clz],: ionic mol/dm3 mol/dm3 mol/dm3 strength 0.05 0.28 0.30 0.05, 0.20 0.075 0.30 0.075, 0.175 0.28 0.1, 0.15 0.1 0.30 0.28 0.30 0.25 0.28 0.25, 0 0.05, 0.45 0.05 0.53 0.55 0.55 0.25 0.53 0.25, 0.25 0.55 0.5 0.53 0.5, 0 a
k, dms/(mol.s)
*
(1.2 0.2) X (7.2 i 1.0) X (6.1 i 0.8) X (3.0 f 0.4) X (4.0 i 1.0) X (9.2 1.5) X (5.1 i 0.7) X
lo0
IO6 lo6 lo6 lo6 lo6
IO6
Sum of contributions from HCl, CuC12-2H,0,and FeC13.6H20.
in the RRDE, the temperature range that could be studied was limited to 23-40 "C.
Results and Discussion Reaction Order. The Cu(I)/Fe(III) reaction may be written as Cu(1)
+ Fe(II1) & Cu(I1) + Fe(I1) k,
where reversibility has been assumed. The equilibrium lies far to the right (Orth, 1988; Orth et al., 1989; Kimura et al., 1984),and the kinetics of the forward reaction were studied under conditions such that the effects of the back reaction were negligible. Redox reactions between transition metals are ordinarily assumed to be first order in each metal, and this assumption was checked before proceeding with other experiments. Confirmation that the reaction is first order in Fe(II1) was obtained by comparing experimental and calculated Cu(1) ring currents obtained for Fe(II1) concentrations between 0.22 and 0.44 mmol/dm3; the calculations assumed a first-order dependence. Within experimental error, F-tests indicated that the agreement was independent of Fe(II1) concentration. For Cu(I), the situation may be more complicated. The Cu(1) concentration is controlled by the disk current. Therefore, calculated and experimental Cu(1) ring currents were compared for a range of disk currents (0-600 PA). Slight but consistent deviations at higher rotation speeds and lower disk currents were observed. A t these conditions, both the Cu(1) concentration and the time for the reaction are low. The resulting total ring current is also low, but the Fe(II)/ Fe(1II) contribution is proportionately large, causing a large absolute error. Although the possibility of systematic experimental error cannot be ruled out, the discrepancies may also result from a reaction order slightly different from 1.0. The reaction mechanism involves parallel pathways that would be affected differently by changes in species distributions, suggesting that this possibility should be considered further. The data discussed in this paper, however, can be satisfactorily explained at the 95% confidence level with a Cu(1) reaction order of 1.0. Effect of H+Concentration. Because iron(II1) hydrolyzes and precipitates readily, it was not possible to directly measure the kinetics in acid-free solutions. To isolate the effects of chloride ion and proton concentration, mixed HCl/KCl solutions were used. The acid concentration was varied from 0.05 to 0.25 mol/dm3 with total chloride at 0.25 mol/dm3 and from 0.05 to 0.50 mol/dm3 in 0.53 mol/dm3 chloride. The ionic strengths were 0.30 and 0.55 mol/dm3, respectively, and the temperature was 23 "C. Data are shown in Table I, and a linear relationship between the rate constant and the inverse of the proton concentration was found. The intercept of a plot of k versus l / [ H f ] was non-zero at the 95% confidence level.
1180 Ind. Eng. Chem. Res., Vol. 29, No. 7, 1990 Table 11. Chloride Ion Effect on k medium HCl, HC104, [Cl-lT? [Ht], ionic mol/dm3 mol/dm3 mol/dm3 strength 1.0 1.05 0.08 0.05, 0.95 1.0 1.05 0.28 0.25, 0.75 1.0 1.05 0.53 0.5, 0.5 1.0 1.05 0.7, 0.3 0.73 1.0 1.05 0.8, 0.2 0.83 1.0 1.05 1.03 1.0, 0 0.08 0.5 0.55 0.05, 0.45 0.5 0.55 0.28 0.25, 0.25 0.5 0.55 0.4, 0,l 0.43 0.5 0.55 0.5, 0 0.53 a
k, dm3/(mol.s) (8.7 f 1.5) X lo4 (1.8 f 0.3) X lo5 (2.6 f 0.5) X lo5 (4.6 f 0.7) X lo5 (7.2 f 1.2) X lo5 (8.5 f 2.1) X lo5 (7.2 f 1.5) X lo4 (2.1 f 0.4) X lo5 (3.6 f 0.5) X lo5 (5.1 f 0.7) X lo5
Table 111. Linear hlationshiD between k and ICl-l*/THtl slope, intercept, exDtl variable dm6/(mol*.s) dm3/(mol.s) R 7.0 x 105 10.5 x 104 HCl concn 0.996 7.6 x 105 8.5 x 104 0.971 Cl- concnu 0.999 7.0 x 105 8.3 x 104 H+ concnb 7.0 x 105 9.7 x 104 0.975 all expts
Data from Table 11. *Data from Table I.
Sum of contributions from HC1, CuC1,.2Hz0, and FeC13.6H20. 0 2 5 mol/dm' HCI 0 30 Ionic Strength
00
I 3 19
324
329
IITEMPERATURE ( IK
334
0.4
08
IO
I 6
Figure 2. k versus [C1-I2/[H+]for all sets of experiments. Intercept = (9.7 f 1.7) X lo4 dm /(mobs); slope = (7.0 f 1.8) X 106 dm6/ (mo12.s).
336
x
Figure 1. Arrhenius plot. Activation energy = 86.4 f 13.8 kJ/mol.
Mechanistic implications of the non-zero value are discussed below. Effect of Total Chloride Concentration. With mixed HC1/HC1O4 solutions, the effect of variations in the chloride concentration was investigated at H+ concentrations of 0.5 and 1.0 mol/dm3 and a t a temperature of 23 "C. The total chloride concentration was the sum of contributions from HC1, CuC12.2H20,and FeCl3.GH2O. At the two H+ concentrations, the ionic strengths were 0.55 and 1.05 mol/dm3, respectively. Data from these runs are given in Table 11. At the 95% confidence level, plots of the rate constant versus [c1-]T2 for each H+ concentration are linear and have comparable non-zero intercept values. Temperature Effect. Reaction rates were determined at 23, 30,35, and 40 OC. The acid concentration was 0.25 mol/dm3 HC1, and the total chloride was 0.28 mol/dm3. An Arrhenius plot of the data obtained is shown in Figure 1. The activation energy was 86.4 kJ/mol, which is typical of many redox reactions involving metals (Armstrong and Halpern, 1957; Basolo and Pearson, 1967; Bennett and Sheppard, 1962; Daugherty and Newton, 1963; Dulz and Sutin, 1963; Harkness and Halpern, 1959; Newton and Baker, 1963), and the R 2 value for the plot was 0.9882. Rate Law. The experiments with varying HCl, H+, and C1- concentrations indicated that the rate constant is linearly related to [C1-IT2/[H+];R 2 values obtained by fitting k versus [C1-]T2/[H+]are given in Table I11 and indicate a good correlation for each set of experiments. Ionic strength is not an important variable; it ranged from 0.1 to 1.05 mol/dm3 in the HC1 experiments, but this had no discernible effect on the fitting procedure. Bearing in mind
the non-zero intercepts of the k versus l/[H+] and [c1-]T2 plots, the 23 "C rate data were fitted to an equation of the form rate = k[Cu(I)][Fe(III)]
(la)
with
The values of c1 and c2 were found to be 9.7 X lo4 dm3/ (mo1.s) and 7.0 X lo5dm6/(mo12.s),respectively, with an R value of 0.975 when all of the data were fitted at once. (See Figure 2.) The activation energy and preexponential factor obtained for 0.25 mol/dm3 HC1 at varying temperatures were combined with the expression for concentration dependence to give the following constant for the oxidation of Cu(1) by Fe(II1) in chloride solutions:
k = 1.8 x
1015e-86.4x103/R
The units of k are dm3/(mol-s);the activation energy is in kJ/mol. Over the concentrations and temperatures investigated, the equation is accurate to within f25%. If it is assumed that the reverse reaction is first order in both Cu(I1) and Fe(I1) and that the system is sufficiently close to equilibrium, the reverse rate constant may be written as
Ind. Eng. Chem. Res., Vol. 29, No. 7, 1990 1181 A-
230c
'.
m
Table IV. Cu(1) and Fe(II1) Species and Their Respective Cumulative Formation or Hydrolysis Constants' species cum formation or hydrolysis const
cu+
CUCP cuc1,cuc1,2cuzc1:Fe3+ FeC12+ FeC12+ FeC1,O FeClL FeOHZ+ Fe(OH)*+ Fez(OH)z4+
/
l
/
iot
;t P E
U
c
-I
l
/
-.II 0.0
4.17 X 5.37 x 2.14 X 2.57 x 2.00 x 2.08 X 2.09 x
104 105 105 IOio
IOo
100 1O-I 10-3
10-3 lo4 (20 "C) 10-3
n V a l u e ~obtained from Kimura et al. (1984). Ionic strength = 1 at 25 OC unless otherwise noted.
/d
-6
1.40 x 2.74 x 1.80 x 4.65 X
"
0.2
'
I
0.4
'
"
0.6
I
0.8
'
'
I.o
HCI CONCENTRATION (mol/dm3 1 Figure 3. k, as a function of hydrochloric acid concentration. Obtained by combining the best line fitted to the data in Figure 2 with Kq values from references by Orth (1988) and Orth et al. (1989).
where K , is the equilibrium constant and may be expressed by
ing that both Cu(1) and Fe(II1) obey equilibrium species distributions and that the rate has a fit-order dependence on Cu(1) and Fe(III), the constants in Table IV may be used to factor out the concentrations of free Cu+ and Fe3+. For the [CU,C~,~-] [FeOH2+] pair, introduction of the equilibrium constants gives
r
1
[Cu+][C1-]4 Because of the pronounced effect of C1- in stabilizing univalent copper, Keqdepends quite strongly on solution concentrations and decreases with increasing chloride concentration. K , values at 23 "C have been reported previously (Orth, 1988; Orth et al., 1989). Combining these with the k data reported above results in the predicted k, values shown in Figure 3. As a result of the strong chloride ion dependence of K , and its lack of dependence on H+, the reverse rate constant is a strong nonlinear function of hydrochloride acid concentration. Similar trends were found for rate data obtained at constant total Cl- or constant H+ conditions. Reaction Mechanism. The observed reaction rate reflects the net result of all the possible interactions between copper(1) species (Cu+, CuClO, CuC12-, etc.) and iron(II1) species (Fe3+,FeCl+, FeOH+, etc.). Considering all possible pairwise interactions involving one univalent copper species and one trivalent iron species, the measured rate can be represented as rate = CCkijCiCj (4) i
j
where Ci and Cj are the concentrations of the individual Fe(II1) species i and Cu(1) species j and kij is the rate constant for the reaction between them. A plausible mechanism should reconcile this concept with the experimentally observed linear dependence of the rate constant On [ c l - ] ~ ~ / [ H + ] . The Cu(1) and Fe(II1) species for which equilibrium data have been reported at an ionic strength of 1.0 and temperatures near ambient are tabulated in Table IV, with their cumulative formation or hydrolysis constants. The very strong complexing of Cu+ by C1- may be noted, along with the weaker chloride complexing and much weaker hydrolysis of Fe3+. If each Cu(1) species is paired with each Fe(II1) species, eq 4 may be written out term by term. A total of 40 such terms can be written. A representative term of eq 4, for reaction of Cu2C12- with FeOH2+,would be proportional to the product k ( C u 2 C 4 ~ , F ~ H[CuzC142-] z+) [FeOH2+].Assum-
[H+l
[Cu+][Fe3+]
From experiment, however, a linear dependence of k on [C1-IT2/[H+]is observed (eq lb), as is a proportionality to [Cu+][Fe3+].Since the total metal concentrations are small relative to the total chloride level, [Cl-1, z [Cl-] i.e., it is not necessary to distinguish between total and free chloride. Of the 40 terms in eq 4, most can immediately be ruled out because they lack the proper form to be consistent with the empirical rate law (eqs l a and lb). Those consistent with the observed concentration dependence must be proportional to [Cu+][Fe3+]or [C1-I2[Cu+][Fe3+]/[H+].There is only one term of the first type, for direct reaction of Cu+ and Fe3+. This reaction thus accounts for the constant term in eq lb, and c1 may be identified with k(Cu+,Fe"+.Likewise, there is only one term with the proper combination of [C1-I2and l/[H+] dependence, corresponding to reaction between CuC12- and FeOH+, and the constant c2 of eq l b may be interpreted as (2.74 X 105)(2.00X 10-3)k(CuCh-,~e~~+). The Cu2C12- + FeOH2+term of eq 4 can be seen to have the wrong dependence on the Cu+ and C1- concentrations to account for eq lb. Likewise, 37 other terms in eq 4 cannot account for the observed concentration dependencies. The forward rate constant of eq 1b may be then expressed as
and the overall rate is rate = k[Cu+][Fe3+] (6) A series of speciation calculations was made by using the model of Kimura et al. (1984) to determine whether there would be significant shifts in species distributions during reaction. The component concentrations considered were 1.0 mol/dm3 HC1 and total Fe and total Cu of 0.44 and 16.4 mmol/dm3, respectively. The extent of reaction, or frac-
1182 Ind. Eng. Chem. Res., Vol. 29, No. 7, 1990
tion of Cu(1) oxidized, was increased from 0.0 to 0.995. The to concentration of free Cu+ decreased from 0.12 X mmol/dm3 and that of free Fe3+from 0.40 X 0.62 X lo-’ to 0.20 X 10” mmol/dm3 as the extent of reaction increased over the range specified. However, the fraction of monovalent copper in the Cu+ form remained virtually constant at 2.8 X lo4; similarly, the fraction of Fe(II1) present as Fe3+was 0.092 throughout the reaction. Thus, the same form of the rate expression, though not the same values for the constants, would be expected whether [Cu’] and [Fe3+]or [Cu(I)] and [Fe(III)]are used as reactant concentrations. The mechanism suggested by these considerations is the following: First, the CuC12-and FeOH2+species must form; this is assumed to occur by reversible reactions c u + + 2c1- F= cuc1,(7) Fe3+ + H 2 0 + FeOH2++ H+ (8) Other complex species of course are also present but are not involved to a significant extent in the redox reaction. There are then parallel redox reactions Cu+ Fe3+ Fez+ Cu2+ (9) and CuC1,- + FeOH2+ FeOH+ + CuC120 (loa) or CuC1; + FeOH2+ FeClz0+ CuOH+ (lob) Equation loa, involving only an electron transfer, is an outer-sphere mechanism. The inner-sphere mechanism of eq 10b involves exchange of both an electron and the ligands. While there is no direct experimental evidence to discriminate between the two possibilities, the latter requires transfer of two chlorides and one hydroxide in addition to an electron, with formation of CuOH+, and therefore seems unlikely. If it occurs, it is presumably not an elementary reaction. Equation 10a therefore is more plausible. As a further test of this mechanism, experimental values for the additive terms in
-
+
+
-
-
I
rate = 9.7
eq 12 becomes (Orth, 1988)
k = 2.3
X
lo4 + 1.6 X 105/[H+]
(13)
which can be directly compared with eq 2 of this work. Except for the absence of a ligand concentration in eq 13, the two equations have the same form. Furthermore, the constants are of the same magnitude, and (7.0 X 1@)/(9.7 X lo4) zz 7 . Chloride ion does appear to enhance the rate, and this is in accord with conclusions from other kinetic studies (Orth, 1988). Conclusions The reaction Cu(1) + Fe(II1) == Cu(I1) + Fe(I1) has been studied in chloride media by using a recently developed electrochemical technique that enables overlapping currents at the ring electrode of a rotating ring disk assembly to be assigned to their individual half-reactions. The reaction is first order in both copper(1) and iron(II1). The reaction rate constant depends inversely on the H+ concentration and on the square of the total chloride concentration; in addition, there is a contribution to the rate constant that is independent of these parameters. The empirical rate law, eq 1, was explained in terms of a mechanism involving parallel reactions between Cu+ and Fe3+,accounting for the constant term, and CuC12- and FeOH2+,responsible for the chloride and proton dependence. The activation energy of the overall reaction was found to be 86.4 kJ/mol. Ionic strength has no significant effect on the rate of reaction up to approximately 1 mol/dm3. The mechanism and rate law are in good agreement with those previously reported for perchlorate solution by Parker and Espenson (1969). The chloride ligand, however, does enhance the overall reaction rate. Acknowledgment This work was supported by the National Science Foundation under Grant MSM-84-00233. Registry No. Cu, 17493-86-6; Fe, 20074-52-6; HCl, 7647-01-0.
I
X
lo4 + 7.0
[Cl-],
X
lo5-
[H+l
[Cu(I)I[Fe(III)I
(11) were combined with calculated Cu+, Fe3+, CuC12-, and FeOH2+species concentrations to estimate lqCu+,+3+) and k(CuCliS”eO~) for representative conditions. The computed values of the individual rate constants are of reasonable magnitude when compared with constants reported for various other redox reactions (Orth, 1988). Parker and Espenson (1969) conducted a somewhat similar study which is a useful point of comparison. Two methodological differences do need to be mentioned: the earlier work was done in perchlorate media and used the stopped-flow technique. Parker and Espenson reported that the rate constant was a function of l/[H+]. The mechanism proposed by Parker and Espenson involves a single pathway with reaction between FeOH2+and Cu+. From the error analysis, Parker and Espenson concluded that if the rate constant had a non-zero intercept, i.e., if
k = cl’
+C2’
(12)
[H+l the ratio c2//cIf would be 27; c2/ was given as 1.6 X lo5 s-l. The fact that the second term in the sum is much larger than the first made it impossible to determine clf accurately. If, however, the lower limit of 7 is taken for c;/clf,
Literature Cited Abdullin, R. S.; Kal’chev, V. P.; Pen’kov, I. N. Investigation of Copper Minerals by NQR Crystallochemistry,Electronic Structure, Lattice Dynamics. Phys. Chem. Miner. 1987,14,25&263. Armstrong, A. M.; Halpern, J. Kinetics of the Oxidation of Mercury (I) by Thallium (111) in Aqueous Solution. Can. J . Chem. 1957, 35, 1020-1030. Basolo, F.; Pearson, R. G. Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution;John Wiley and Sons: New York, 1967;p 466. Bennett, L. E.; Sheppard, J. C. The Kinetics of the Iron (1I)aq-Cobalt (1II)aq Reaction. J . Phys. Chem. 1962,66, 1275-1279. Daugherty, N. A.; Newton, T. W. The Kinetics of the Reaction Between Vanadium (V) and Iron (11). J. Phys. Chem. 1963,67, 1090-1093. Dub,G.; Sutin, N. The Kinetics of the Oxidation of Iron (11) and ita Substituted tris(1,lO-Phenanthroline)Complexes by Cerium (IV). Znorg. Chem. 1963,2,917-921. Harkness, A. C.; Halpern, J. Kinetics of the Oxidation of Uranium (IV) by Thallium (111). J.Am. Chem. SOC.1959,81,3526-3529. Kimura, R. T.; Haunschild, P. A.; Liddell, K. C. A Mathematical Model for Calculation of Equilibrium Solution Speciations for the FeCl,-FeC12-CuClz-CuCl-HCl-NaCl-H20System at 25 “C. Metall. Trans. B 1984, 15B,213-219. Nakaj, I.; Sugitani, Y.; Nagashima, K.; Niwa, Y. X-ray Photoelectron Spectroscopic Study of Copper Minerals. J . Inorg. Nucl. Chem. 1978,40,789-791. Newton, T. W.; Baker, F. B. The Kinetics of the Reaction Between Plutonium (VI) and Iron (11). J.Phys. Chem. 1963,67,1425-1432. Orth, R. J. Rotating Ring Disk Electrodes: A Technique for Evaluating the Kinetics for Systems with Overlapping Ring Currents.
Znd. Eng. Chem. Res. 1990, 29, 1183-1189 Ph.D. DiMertation, Washington State University, Pullman,1988. Orth, R. J.; Parikh, R. S.; Liddell, K. C. Application of the Rotating Ring-Disk Electrode in Determining the Second Order Rate Constant for the Reaction Between Cu(1) and Fe(II1) in 1.0 mol/dm3 HC1. J . Electrochem. SOC.1989,136, 2924-2930. Parker, 0. J.; Espenson, J. H. Reactions Involving Copper (I) in Perchlorate Solution. A Kinetic Study of the Reduction of Iron (111) by Copper (I). Inorg. Chem. 1969,8, 1523-1526. Peters, E.; Swinkels, G. M.; Vizsolyi, A. Copper Recovery from Sulfide Concentrates by the UBC-Cominco Ferric Chloride Leach
1183
Route. In Process and Fundamental Considerations of Selected Hydrometallurgical Systems; Kuhn, M. C., Ed.; Society of Mining Engineers: New York, 1981;pp 71-81. Schweitzer, F. W.; Livingston, R. W. Duval's CLEAR Hydrometallurgical Process. In Chloride Electrometallurgy; Parker, P. D., Ed.; The Metallurgical Society of AIME: Warrendale, PA, 1982; pp 221-227.
Received for review April 3, 1989 Accepted February 12,1990
Demetalation Chemistry: Control of Vanadium on Fluid Cracking Catalyst Jin S. Yoo,* Emmitt H. Burk, Jr., John A. Karch, and Andrew P. Voss Mail Station 0-4,Amoco Chemical Company, Amoco Research Center, P.O. Box 3011, Naperuille, Illinois 6&66
Some fraction (- 13%) of the vanadium deposited on the regenerated equilibrium FCC catalyst in the FCC unit was readily removed from the catalyst matrix by a water wash. The vanadium removal was improved to -28% by applying the reductive wash with an aqueous sulfur dioxide solution followed by the oxidative wash with an aqueous hydrogen peroxide solution to the same regenerated catalyst. A significant fraction of vanadium, 50-70%, was further removed by subjecting the same catalyst to the calcination step at higher temperatures (730-815 "C) before the reductive and oxidative washing scheme. The microactivity test results showed that the catalytic activity of the devanadated FCC equilibrium catalyst was remarkably rejuvenated. The improvement in the catalytic activity of the devanadated system is directly tied with the degree of vanadium removal. These results clearly suggest that the poisoning effect of the vanadium on the FCC catalyst can at least partially be reversed by removing vanadium from the metal-contaminated FCC catalyst by a simple calcination/washing process. Metal-contaminated fluid cracking catalysts have been studied by SIMS (secondary ion mass spectroscopy) (Jaeras, 1982), ESCA and atomic absorption spectrophotometry (Lars et al., 1984), EMPA (electron microprobe analysis), and DTA (Masuda et al., 1983a,b, 1985). These studies showed that nickel was homogeneously distributed throughout the catalyst surface and that vanadium was preferentially deposited on the zeolite sites and reacted destructively with zeolite. Vanadium pentoxide interacts with lanthanum oxide exchanged with the zeolite sites to form an eutectic mixture, which becomes a major cause for destruction of the zeolite and the drastic decline of catalytic activity, surface area, crystallinity, and pore volume (Occelli et al., 1985; Masuda et al., 1983a,b). The poison precursor was identified as volatile vanadic acid, H3V03,which exists in 1-10 ppm level in a typical FCC regenerator at conditions of 730 "C, 20% steam, and 2 atm of total pressure (Wormsbecher et al., 1986). V205(s)+ 3H@(v) % ~ H ~ V O ~ ( V ) The affinity of MgO, La203,and a-A1203for V205was found to be considerably higher than rare earth metal exchanged zeolite, REHY, and other metal oxides such as Zr02, Ti02,and SiOz (Masuda et al., 1985). In the FCC unit, the degree of interparticle migration of vandium directly coincided with the affinity trend of metal oxides (Masuda et al., 1988). The idea of using MgO as a vanadium scavenger has been proposed (Wormsbecher et al., 1986; Masuda et al., 1988). A large number of vanadium oxide phases have been
* Author to whom correspondence should be addressed. 0888-5885/90/2629-1183$02.50/0
identified for the vanadium-oxygen system. An incomplete phase diagram of vanadium oxide has been published in 1958 (Rostocker, 1958) and more complete one later (Stringer, 1965). The identities of five oxides such as VO1.75, VOl.so,VO1.84, V01.85,and V01.8, have been well established. These are obviously members of a homologous series of general formula VnOln+ where n equals 4-8. These were prepared by roasting the mixture of V203and V 0 2 at elevated temperature (Samoilov et al., 1980). Reduction of V205 and ammonium metavanadate in a hydrogen atmosphere at 375 "C produced Vz03 via these homologous oxide intermediates (Sato et al., 1970, 1968). V409is also shown to exist under the oxidation-reduction situation (Srivastava et al., 1982). V20,
-
V404
+
V,013
+
V02
-
Vn02n-1
-
V2O3
An equilibrium exists between V205 and Vz04 as expressed in the equation below (Samyratovet al., 1967). An
+
2V204 2 V205 V20,
increase in temperature from 200 to 600 "C displaced the equilibrium to the right, favoring the formation of pentoxide. A further increase in temperature from 600 to 1200 "C favored the V204formation by shifting the equilibrium to the left. Laser Raman spectroscopy on the structure of vanadium ions on the V204/a-AlzO3catalyst identified three different species (Roozeboom et al., 1978). Numerous approaches have been taken to remove vanadium oxide(s) from the metal-contaminated fluid cracking catalysts. Various solutions tried in the liquidphase washinglleaching processes include a dilute solution of complexing agent such as oxalic acid (Beuther and Flinn, 1963),and an ammoniacal solution containing a chelating 0 1990 American Chemical Society
