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A mixed surface reaction kinetic model for the reductive leaching of manganese dioxide with acidic sulfur dioxide
Hydrometallurgy 73 (2004) 215 – 224 www.elsevier.com/locate/hydromet
A mixed surface reaction kinetic model for the reductive leaching of manganese dioxide with acidic sulfur dioxide Gamini Senanayake * Department of Mineral Science and Extractive Metallurgy, Murdoch University, Perth, WA 6150, Australia Received 21 November 2002; received in revised form 16 September 2003; accepted 25 October 2003
Abstract Previous researchers have rationalised the kinetics of dissolution of manganese dioxide in acidic sulfur dioxide solutions on the basis of either the adsorption of reactants followed by reactions with HSO 3 and SO2 or a two-electron surface + electrochemical reaction with H+ and HSO 3 , the species involved in the equilibrium: SO2(aq) + H2O = H + HSO3 . This paper revisits the reported rate data for the dissolution of MnO2 in the form of plates [Herring and Ravitz, Trans. SME/ AIME 231 (1965) 191] or particles [Miller and Wan, Hydrometallurgy 10 (1983) 219] to show that the rate of dissolution can be rationalised on the basis of a fast redox reaction : (i) MnO2(s) + SO2(aq) + H2O = MnOOH(s) + HSO3(aq) followed by one of the two rate determining redox reactions: (ii) MnOOH(s) + HSO3(aq) = MnSO04(aq) + H2O or Mn(OH)HSO4(s), (iii) MnOOH(s) + SO2(aq) = MnO(s) + HSO3(aq) depending on the pH of the medium. Whilst the rate equation for (iii) agrees with the electrochemical kinetic model reported previously, the dissolution of the solid Mn(OH)HSO4 or the direct reaction: (iv) MnO2(s) + SO2(aq) = MnSO04(aq) also appear to be rate controlling. The two rate constants k1 and k2 based on a shrinking sphere kinetic model: 1 (1 X)1/3 = k1K 0.5[SO2]0.5(t/rq) + k2[SO2](t/rq), for the leaching of monosized pyrolusite and electrolytically prepared MnO2 particles, are in reasonable agreement with k1 and k2 based on the dissolution kinetics of MnO2 plates. D 2003 Elsevier B.V. All rights reserved. Keywords: Mixed surface reaction kinetic model; Manganese dioxide; Acidic sulfur dioxide
1. Introduction The reductive leaching of manganese dioxide with sulfur dioxide in acid media has been the subject of a number of recent research publications mainly due to the fact that SO2(aq) offers fast leaching kinetics and the choice of selective leaching of metals from limo-
* Tel.: +61-8-9360-2833; fax: +61-8-9360-6343. E-mail address: [email protected] (G. Senanayake). 0304-386X/$ - see front matter D 2003 Elsevier B.V. All rights reserved. doi:10.1016/j.hydromet.2003.10.010
nitic laterite, low-grade manganiferrous ores and deep-sea nodules (Das et al., 1997; Senanayake and Das, 2003; Abburzzese, 1990; Grimanelis et al., 1992; Kanungo and Das, 1988). Therefore a proper understanding of the reaction mechanism is useful for the optimisation of selective leaching of manganese from these materials. Herring and Ravitz (1965) rationalised the dissolution kinetics of MnO2 plates on the basis of reactions with SO2 and HSO3, the two sulfur species in equilibrium with H+ in Eq. (1). More recently Miller and Wan (1983) carried out a detailed kinetic analysis based on electrochemical and leach-
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
ing studies to propose a two-electron reaction mechanism, which involves H+ and HSO3 (Eqs. (2) –(4)). SO2 þ H2 O ¼ Hþ þ HSO 3
ð1Þ
pMnO2 ðsÞ þ 4Hþ þ 2e ¼ Mn2þ þ 2H2 O ðcathodic reactionÞ
ð2Þ
þ 2 HSO 3 þ H2 O¼3H þ SO4 þ 2e ðanodic reactionÞ
ð3Þ pMnO2 ðsÞ þ SO2 ðaqÞ ¼ Mn2þ ðaqÞ þ SO2 4 ðaqÞ ðoverall reactionÞ
ð4Þ
However, the equilibrium constants (K) reported in Table 1 show that Mn2 + forms aqueous complex species with both SO42 and SO32 ions and the magnitude of log K for MnSO30 (aq) is higher than that for MnSO40 (aq). This highlights the need to consider the distribution of manganese and sulfur species MnSO40, MnSO30, Mn2 +, HSO3, SO2, HSO4, and SO42 of different oxidation numbers Mn(II) – S(IV) –S(VI) shown in Figs. 1 and 2. Percentage of SO42 and MnSO40 (aq) increases with increasing pH
Table 1 Equilibrium constants (log K ) for reactions of S(IV) and S(VI) species with H+ and Mn2 + Reaction
Ionic strength (I )
Log K at 25 jCa
H+ + SO42 = HSO 4 H+ + HSO 3 = H2O + SO2
0 1 0
H+ + SO32 = HSO 3 Mn2 + SO42 = MnSO04(aq)
1 1.2 0 0 0 0
1.99 1.07 1.77b (1.94b at 45 jC) 1.37 6.34 2.26 (3.0 at 45 jC) 3.0 0.27c 3.61c
0 0
3.34c 5.11c
Mn2 + + SO23 = MnSO03(aq) + 0 Mn2 + HSO 4 = MnSO4(aq) + H 0 MnSO3(aq)+ HSO4 = MnSO04(aq) + HSO 3 + 0 Mn2 + + HSO 3 = MnSO3(aq) + H 2+ Mn + SO2 + H2O = MnSO03(aq) + 2H+ a
Data from Sillen and Martell (1964). Herring and Ravitz (1965) used these values of K in Eq. (9). c Derived values from the data in this table. b
Fig. 1. Effect of pH on percentage distribution of species in Mn(II) – sulfate system at 25 jC, [Mn(II)] = 0.001 M, [S(VI)] = 0.1 M, based on thermodynamic data in Table 1.
and therefore MnSO40 (aq) is the predominant species in solutions of pH > 0 and HSO4 is the predominant S(VI) species at pH < 1.5 (Fig. 1). In contrast, SO32 is unstable in acidic solutions of S(IV) and therefore the complex species MnSO30 (aq) becomes the predominant species only at pH values greater than 4; Mn2 + is predominant at low pH values (Fig. 2). The log K values derived for the reaction between Mn2 + and HSO4, SO2 or HSO3 shown in Table 1 reveal that MnSO40 is thermodynamically more favourable than MnSO30. Therefore Mn(II) in an acidic leach liquor produced by the reduction of MnO2 with SO2 is likely to be in the complex form MnSO40 (aq) rather than the uncomplexed Mn2 + (aq). Moreover, SO2 can be oxidised from the oxidation state S(IV) to the two oxidation states S(V) and S(VI) to produce S2O62 (dithionate) and SO42 ions,
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
MnðIVÞ þ SðIVÞ ¼ MnðIIÞ þ SðVIÞ
217
ð7Þ
2MnO2 ðsÞ þ 3SO2 ðaqÞ ¼ MnSO4 ðaqÞþ MnS2 O6 ðaqÞ ð8Þ
Fig. 2. Effect of pH on percentage distribution of species in Mn(II) – sulfite system at 25 jC, [Mn(II)] = 0.001 M, [S(IV)] = 0.1 M, based on thermodynamic data in Table 1.
respectively. Bassett and Parker (1951) noted that dithionate may be formed when MnO2 is reduced in one-electron steps (Eqs. (5) and (6)) whereas sulfate may be formed by a two-electron reduction mechanism (Eq. (7)), where Mn(IV) = MnO2, Mn(III) = Mn 2 O 3 (or MnOOH), S(IV) = SO 2 or HSO 3, S(V) = HSO3, H2S2O6 or S2O62 and S(VI) = HSO4 or SO42 . These two modes of reactions led to the overall reaction (Eq. (8)) between MnO2 and SO2 reported in the literature by Meyer and Schramm (Miller and Wan, 1983). MnðIVÞ þ SðIVÞ ¼ MnðIIIÞ þ SðVÞ
MnðIIIÞ þ SðIVÞ ¼ MnðIIÞ þ SðVÞ
ð5Þ
ð6Þ
The extent to which the two modes of oxidation will occur depends on the Mn/O molar ratio of the starting oxide material, i.e., the lattice structure of MnO2 (Bassett and Parker, 1951). For example, pyrolusite and MnO2 produced by the thermal decomposition of Mn(NO3)2 react with SO2 to produce mainly dithionate (95 –97%). In contrast, MnO2 prepared electrolytically (EMD) or by wet methods and dried at a moderate temperature (96 jC) produces 74– 79% dithionate. Freshly precipitated amorphous or partly crystalline MnO2, used in the wet state, produces low and variable percentages of dithionate. The rapid dissolution of amorphous MnO2 in SO2 (aq) led to the suggestion that the reduction can take place on the surface as well as in solution (Bassett and Parker, 1951). Therefore it is important to establish the mechanism of the surface reaction with respect to both one-electron and two-electron reaction mechanisms. This paper makes use of the results for the dissolution of EMD plates of known surface area (1 –3 cm2) in stirred solutions (Herring and Ravitz, 1965) to establish the mechanism of the reaction between MnO2 and SO2 over a wide range of conditions: pH 1– 4.5, total SO2 concentration 0.03 – 0.46 mol L 1 and temperature 25 –45 jC. Three different rate determining steps are considered to rationalise the effect of pH and total SO2 concentration on the initial rate of dissolution of MnO2. Two of them are used to rationalise the reported leaching results for pyrolusite and EMD in acidic sulfur dioxide in the concentration range 0.03– 1 mol L 1 at pH 2 (Miller and Wan, 1983).
2. Discussion 2.1. Effect of SO2 2.1.1. Dissolution of EMD plates Table 2 lists the different conditions used by Herring and Ravitz (1965) in three sets of experiments
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Table 2 Materials and conditions used in kinetic studies and fraction of manganese (X ) leached Material
Test Temperature [SO2(t)]a (jC)
EMD platesb EMD platesb EMD platesb EMD particlese
A
25
0.03 – 0.46 1.8 – 1.1c
B
45
0.03 – 0.27 1.9 – 1.3c
C
25
0.06 – 0.43 4.7 – 1.0d
D
26
0.07
30
0.30 0.69 0.03
Pyrolusite E particlese
0.27 0.52 0.98
pH
Time X (s)
2d
300
0.35
2d
300 300 480
0.68 0.91 0.15
480 480 480
0.41 0.64 0.76
Total SO2 mol L 1. Dissolution of plates of EMD of constant surface area, at a stirring speed of 500 rpm (Herring and Ravitz, 1965). c Natural pH. d pH adjusted using H2SO4 or NaOH. e Monosized (357 Am) particle leaching at 0.2% solids and 650 rpm (Miller and Wan, 1983). a
b
curves in Fig. 4, becomes less significant with the increase in [SO2(t)] and pH, particularly at high pH. At a given value of [SO2(t)] = 0.2 mol L 1 (for example), the rate decreases with the increase in pH from 1.1 to 4.4. This shows the effect of the decrease in [H+] as well as the change in speciation summarized in Figs. 1 and 2 on the rate of dissolution of MnO2. 2.1.2. Dissolution of pyrolusite and EMD particles Miller and Wan (1983) reported the leaching results of monosized pyrolusite and EMD particles in the concentration range of SO2 0.03 –1 mol L 1 at 26 –30 jC and pH 2. Table 2 lists selected results from their work (Sets D and E) to show that the fraction leached (X) at a given time interval increases with increasing [SO2(t)]. At a total SO2 concentration of 0.3 mol L 1 at 26 jC, 68% of manganese is dissolved from EMD in 5 min. At a comparable total SO2 concentration of 0.27 mol L 1 at 30 jC, the percentage manganese dissolved from pyrolusite is 41%, indicating a slower reaction of SO2 with pyrolusite
(A – C ) using EMD. They used the equilibrium constant for Eq. (1): K ¼ ½Hþ HSO 3 =½SO2
ð9Þ
reported in Table 1 to calculate the natural pH and the concentration of HSO3 and ‘‘free’’ sulfur dioxide at equilibrium denoted by [SO2], in the first set of experiments (A) at 25 jC in solutions of different total concentration of sulfur dioxide [SO2(t)]. A higher temperature of 45 jC was maintained in the second set (B). In the third set (C), the pH adjustments were made by adding H2SO4 or NaOH to solutions containing different [SO2(t)]. Fig. 3 shows the equilibrium concentration of sulfur species, pH and the initial rate as a function of [SO2(t)] in Set A. The general trend is that the rate increases with increasing concentration of HSO3 and free SO2 and decreasing pH. Due to the fact that [H+], [HSO3] and [SO2] are interrelated according to Eq. (9), Fig. 4 shows the rate data in Sets A and C at four different pH values as a function of [SO2(t)]. It is clear that the dependence of reaction rate on [SO2(t)], shown by the slope of the
Fig. 3. Variation of natural pH, [HSO 3 ], [SO2] and initial rate of manganese dissolution at 25 jC (results for EMD plates from Herring and Ravitz, 1965).
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
219
(Eq. (12)) the observed slope in Fig. 5 is 0.71– 0.73 for the results in Sets A and B and at high values of [SO2] in the case of Set C. At low values of [SO2] in Set C, representing the high pH data in Fig. 4, the rate is relatively low and independent of [SO2]. These facts indicate the possibility of other surface reactions which contribute to the rate-limiting step(s). Three surface reaction mechanisms may be used to rationalise these experimental data. 2.2. Reaction schemes and mechanisms 2.2.1. Reaction scheme I at high pH (>3) Higginson and Marshall (1957) proposed the formation of HSO3 (a free radical) in a fast one-electron oxidation step in order to rationalise the kinetics of the reaction between SO2 and Fe(III) in acid media. Thus it is possible that a fast surface reaction between MnO2 and SO2 which produces MnOOH and HSO3
Fig. 4. Variation of initial rate of manganese dissolution with total SO2 concentration and adjusted pH at 25 jC (results for EMD plates from Herring and Ravitz, 1965).
compared to EMD. Miller and Wan (1983) analysed the initial rate data on the basis of the redox reactions described by Eqs. (2) and (3) and proposed the following rate equation: 0:5 0:5 R ¼ d½MnðIIÞ=dt ¼ k ½Hþ HSO ð10Þ 3 When combined with Eq. (9) this may be rewritten in the forms: R ¼ kK 0:5 ½SO2 0:5
ð11Þ
log R ¼ 0:5 log½SO2 þ log kK 0:5
ð12Þ
The general change in rate with [SO2] in Figs. 3 and 4 is consistent with Eq. (11). Although a log – log plot of R vs. [SO2] is expected to have a slope of 0.5
Fig. 5. Log – log plot of rate vs. [SO2], data from Figs. 3 and 4 and Herring and Ravitz (1965).
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
is followed by the rate determining step (RDS) in Eq. (14) to produce MnSO40 (aq) at high pH (>3).
HSO 3 ðaqÞ ¼ HSO3 ðaqÞ þ e ðanodic reactionÞ ð18Þ
pMnOOHðsÞ þ Hþ þ HSO 3
pMnO2 ðsÞ þ SO2 ðaqÞ þ H2 O ¼ pMnOOHðsÞ þ HSO3 ðaqÞðfast redox reactionÞ
¼ MnOðsÞ þ HSO3 ðaqÞðRDS 1Þ
ð19Þ
ð13Þ 2HSO3 ðaqÞ ¼ H2 S2 O6 ðaqÞ ðdimerisation reactionÞ
pMnOOHðsÞ þ HSO3 ðaqÞ ¼ MnSO04 ðaqÞ þ H2 OðRDSÞ
ð20Þ ð14Þ
Eq. (14) does not involve H+, HSO3 or SO2 and thus explains why the rate of dissolution of MnO2 remains relatively independent of [SO2] at high pH (Figs. 4 and 5), despite the fact that sulfur dioxide is predominantly in the form of HSO3 at pH values greater than 3 (Fig. 2). However, in more acidic solutions of pH < 3 it is possible to write two reaction schemes to describe the formation of dithionate and sulfate. 2.2.2. Reaction scheme II at low pH: formation of dithionate In solutions of pH less than 3, a fast surface reaction between MnO2, H+ and HSO3 produces MnOOH and HSO3 (Eqs. (15) and (16)). This is followed by one-electron redox reactions of MnOOH and HSO3 to produce MnO and HSO3 leading to the rate determining step represented in Eq. (19) (RDS-1). The dimerization of HSO3 to H2S2O6 and the subsequent reaction of H2S2O6 with MnO(s) produce MnS2O6 (aq) leading to the overall reaction in Eq. (22): pMnO2 ðsÞ þ Hþ þ HSO 3 ¼ pMnO2 Hþ HSO 3 ðfast equilibrationÞ
ð15Þ
pMnO2 Hþ HSO 3 ¼ pMnOOHðsÞ þ HSO3 ðaqÞðfast redox reactionÞ ð16Þ pMnOOHðsÞ þ Hþ þ e ¼ pMnOðsÞ þ H2 O ðcathodic reactionÞ
ð17Þ
pMnOðsÞ þ 2HSO3 ðaqÞ ðor H2 S2 O6 Þ ¼ MnS2 O6 ðaqÞ þ H2 O ðfastÞ
ð21Þ
pMnO2 ðsÞ þ 2SO2 ðaqÞ ¼ MnS2 O6 ðaqÞ ðoverall reactionÞ
ð22Þ
The rate of dissolution of manganese according to this reaction scheme is controlled by RDS-1 for which the electrochemical kinetic model (Nicol and Lazaro, 2002) described in Appendix A can be used to write the rate equation: 0:5
R1 ¼ k1 ½Hþ
HSO 3
0:5
ð23Þ
2.2.3. Reaction scheme III at low pH: formation of sulfate Instead of producing Mn(III) and S(V) via a oneelectron mechanism (Eq. (13) or (16)), MnO2 may be directly reduced to Mn(II) producing S(VI) according to Eqs. (2) and (3) (Miller and Wan, 1983). This reduction may also take place via the solid product Mn(OH)HSO4(s) instead of MnO(s) considered previously in reaction scheme II. The solid product formed in Eq. (24) is solubilised by SO2(aq) according to Eq. (25) (RDS-2). This leads to the overall reaction described by Eq. (26) (sum of (Eqs. (15), (16), (24) and (25)), which is equivalent to Eq. (4) and agrees with the species distribution discussed with reference to Fig. 1. The rate of dissolution of manganese according to this reaction scheme depends on RDS-2a and/or RDS-2b. The resulting rate equation (Eq. (27)) for both RDS-2a and RDS-2b indicates that
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
221
this reaction scheme is favoured in solutions of high concentrations of free SO2.
the basis of the two rate determining steps in Eqs. (19), (25) and (26) and may be expressed in the form R = R1 + R2;
pMnOOHðsÞ þ HSO3 ðaqÞ
R ¼ k1 ½Hþ ½HSO3 0:5 þk2 ½SO2
0:5
¼ pMnðOHÞHSO4 ðsÞ
ð24Þ
pMnðOHÞHSO4 ðsÞ þ SO2 ðaqÞ ¼
MnSO04 ðaqÞ
þH
þ
HSO 3 ðRDS
2aÞ
ð25Þ
pMnO2 ðsÞ þ SO2 ðaqÞ ¼ MnSO04 ðaqÞ ðoverall reaction or RDS 2bÞ ð26Þ R2 ¼ k2 ½SO2
ð27Þ
2.3. Overall rate equation at pH<3
ð29Þ
when combined with Eq. (9): R ¼ k1 K 0:5 ½SO2 0:5 þk2 ½SO2
ð30Þ
R½SO2 1 ¼ k1 K 0:5 ½SO2 0:5 þk2
ð31Þ
Fig. 6 shows a plot of R[SO2] 1 vs. [SO2] 0.5 for the data from Figs. 3 and 4. The two sets A and B correspond to a good linear correlation and some of the data points from set C lie close to the linear relationship for set A. The data points from set C which show the largest deviations are the same points which show rates close to 6.5 10 8 mol cm 2 s 1 and relatively independent of [SO2] and pH in Fig. 4. The slopes and intercepts of Fig. 6 along with the
The reaction schemes II and III explain the formation of dithionate and sulfate, respectively, at pH < 3. Eqs. (22) and (26) added together give Eq. (8) for the overall reaction, whilst dithionate may also decompose to sulfate: MnS2 O6 ðaqÞ ¼ MnSO04 ðaqÞ þ SO2 ðaqÞ
ð28Þ
The extent of oxidation of SO2 to sulfate or dithionate will depend on the relative rates of RDS-1, RDS2a and RDS-2b in Eqs. (19), (25) and (26). Solutions of lower pH values correspond to high [SO2] (Figs. 2 and 3) and favour reaction scheme III (Eqs. (25) and (26)) over II (Eq. (19)). This is consistent with the observation reported in the literature (Miller and Wan, 1983) that the dithionate formation decreased as the pH was lowered. Furthermore, the observation that the dithionate formation decreased rapidly with the increased rate of agitation (Miller and Wan, 1983) indicates that higher agitation rates would favour reaction schemes I and III which do not form dithionate. The mass transfer of uncharged species HSO3 and SO2 involved in Eqs. (14) and (26) to the reaction surface is likely to be facilitated by the high rates of agitation. However, the contribution of reaction scheme I to the overall rate may be ignored at pH values lower than 3. Thus the overall rate equation for the dissolution of MnO2 at pH < 3 may be derived on
Fig. 6. Correlation between R[SO2] 1 vs. [SO2] 0.5: testing the validity of Eq. (31), data from Figs. 3 – 5.
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
calculated rate constants k1, k2 (see Eq. (31)) and the ratio k1/k2 listed in Table 3 show that the rate constant k1 is twice as large as k2. The activation energy for RDS-1 is 24 kJ mol 1 compared to 23 kJ mol 1 for RDS-2. These values are slightly larger than the reported value of 19 kJ mol 1 (Herring and Ravitz, 1965). 2.4. Application to leaching of monosized particles
the overall rate of leaching. This may be further examined by combining the overall rate equation (Eq. (30)) with the general equation for shrinking sphere model: 1 ð1 X Þ1=3 ¼ k ½SO2 ðt=rqÞ
where r (cm) is the initial particle radius and q (mol cm 3) is the concentration of MnO2 in the solid. The combined equation:
Miller and Wan (1983) showed that their results for the leaching of pyrolusite and EMD obeyed the rate equation corresponding to a shrinking sphere model:
1 ð1 X Þ1=3 ¼ k1 K 0:5 ½SO2 0:5 ðt=rqÞ
1 ð1 X Þ1=3 ¼ k ½SO2 n t
may be rearranged:
ð32Þ
where X = fraction of Mn reacted at time t, n = reaction order with respect to SO2 and k = apparent rate constant. This supports the view that the kinetics of leaching is controlled by the surface reactions alone and the rate decreases with decreasing surface area of the spherical particles with no product layer to retard the diffusion of reactants or products. However, the fact that their results at pH 2 showed a reaction order of 0.58 or 0.59 for the pyrolusite and EMD, respectively, indicates that both RDS-1 and RDS-2 control
ð33Þ
þ k2 ½SO2 ðt=rqÞ
ð34Þ
f1 ð1 X Þ1=3 g½SO2 1 t 1 ¼ k1 K 0:5 ðrqÞ1 ½SO2 0:5 þk2 ðrqÞ1
ð35Þ
Fig. 7 shows a plot of {1 (1 X)1/3}[SO2] 1 t 1 vs. [SO2] 0.5 and confirms the validity of this mixed kinetic model for the leaching of monosized pyrolusite and EMD. The slopes and intercepts depend on the value of K at the appropriate ionic strength (I) used to calculate [SO2]. However, good linear correlations with R2 > 0.99 are observed with the values of K at
Table 3 Rate constants (k1 and k2) obtained for different materials and conditions calculated from slopes and intercepts of Figs. 6 and 7 using K at ionic strengths 0 and 1 Material/test
T (jC)
Slope
Intercept
25 45
0.74e 1.11e
2.67e 4.76e
2.5 10 3g (3.5 10 3)
2.1 10 3g (4.1 10 3)
9.0 10 4g (12 10 4)
8.0 10 4g (16 10 4)
104 rqa (mol cm 2)
103 k1b (cm s 1)
103 k2c (cm s 1)
k1/k2
5.7 10
2.7 4.8
2.1 2.2
9.4
18 (16)
2.0 (3.9)
9.0 (4.1)
9.3
6.4 (5.4)
0.7 (1.5)
9.1 (3.6)
d
EMD plates A B
EMD particlesf D 26
Pyrolusite particlesf E 30
q = wt.% MnO2 density of the material (g cm 3)/molar mass of MnO2 (g mol 1) (Miller and Wan, 1983). Plates: k1 = 10 3 K 0.5 slope; particles: k1 = 103 K 0.5 rq slope; K 0.5 = 7.67 (I = 0, 25 – 30 jC), 9.33 (I = 0, 45 jC), 4.84 (I = 1, 25 – 30 jC); values in parentheses are based on K at I = 1 (see Table 1). c Plates: k2 = 10 3 intercept; particles k2 = 103 rq intercept. d See footnote in Table 2. e Fig. 6 (Eq. (31)). f See footnote in Table 2. g Fig. 7 (Eq. (35)). a
b
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
223
nature of MnO2. Nevertheless, the comparison of k values for the dissolution of MnO2 plates and particles provides useful insight into the rate-limiting step(s).
3. Conclusion
Fig. 7. Correlation between {1 (1 X)1/3}[SO2] 1 t 1 vs. [SO2] 0.5: testing the validity of Eq. (35); conditions and data listed in Table 2.
both ionic strengths I = 0 and 1. Table 3 lists and compares the values of k1 and k2 obtained from the slopes and intercepts of Fig. 7. It is clear that the use of K values at different ionic strengths does not make a significant difference to the rate constants. The k1 value for pyrolusite leaching at 30 jC shows good agreement with that for the dissolution of EMD plates at 25 jC; but k1 for EMD leaching is nearly three times larger than that for pyrolusite. The values for k2 are also in good agreement for the three cases within a factor of 1 – 4 at 25 – 30 jC. The ratio k1/k2 for leaching is in the range 3.6– 9.1 compared to 2.1 for the dissolution of EMD plates. These small differences in rate constants k1 and k2 support the view of Bassett and Parker (1951) that the extent of reduction of MnO2 with SO2 in acid media (pH < 3) according to the two reaction schemes II and III forming dithionate and sulfate depends on the crystalline
Kinetics of reductive dissolution of manganese dioxide in acidic sulfur dioxide media (pH < 3) is controlled by different surface reaction schemes depending on the conditions. In the first scheme a fast one-electron surface reaction produces MnOOH and HSO3. This is followed by another one-electron reaction leading to the production of MnS2O6 and the RDS corresponds to a reaction order of 0.5 with respect to H+ and HSO3. However, MnS2O6 may decompose to MnSO4 and SO2. In the second reaction scheme, SO2 may directly react with MnO2 or two one-electron reactions may produce an insoluble product Mn(OH)HSO4 on the surface. The rate of dissolution of manganese in these cases appears to be first order with respect to SO2 in equilibrium with H+ and HSO3. Thus the overall rate equation for leaching can be expressed by a mixed surface reaction kinetic model combined with a shrinking sphere model which incorporate the two rate constants k1 and k2 : 1 (1 X)1/3= k1K0.5[SO2]0.5 (t/rq)+ k2[SO2] (t/rq).
Appendix A The rate of an electrochemical reaction can be related to the current i, the number of electrons n involved in the reaction and the Faraday constant F: R ¼ i=nF
ð36Þ
According to the electrochemical model (Nicol and La´zaro, 2002) the rate of the two electrochemical reactions in Eqs. (17) and (18) can thus be represented by the cathodic and anodic currents ic and ia : ic ¼ kc ½Hþ expðac FEm =RT Þ
ð37Þ
ia ¼ ka HSO 3 expðaa FEm =RT Þ
ð38Þ
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
where ac and aa are the transfer coefficients, and kc and ka are electrochemical rate constants per unit surface area, and Em is the mixed potential. At the mixed potential Em, jicj = ia = imix and since ac = ac = 0.5, the two equations can be combined: ic ia ¼ ðimix Þ2 ¼ ka kc ½Hþ HSO 3
ð39Þ
to express the rate of dissolution of MnO2 per unit area according to the one-electron mechanism (N = 1 in eq. (36)): 0:5 1 R1 ¼ imix F 1 ¼ ðka kc ½Hþ HSO F 3 Þ
R1 ¼ k1 ½Hþ
0:5
HSO 3
0:5
ð40Þ
ð23Þ
where k1=(kakc)0.5 F 1 represents the rate constant of RDS-1 (Eq. (19)).
References Abbruzzese, C., 1990. Percolation leaching of manganese ore by aqueous sulfur dioxide. Hydrometallurgy 25, 85 – 97.
Bassett, H., Parker, W.G., 1951. The oxidation of sulphurous acid. J. Chem. Soc., 1540 – 1560. Das, G.K, Anand, S., Das, R.P., Muir, D.M., Senanayake, G., Singh, P., Hefter, G., 1997. Acid leaching of nickel laterites in the presence of sulfur dioxide at atmospheric pressure. In: Cooper, W.C., Mihaylov, I. (Eds.), Hydrometallurgy and Refining of Nickel and Cobalt, vol. 1. Canadian Inst. Min. Metall. and Petroleum, Montreal, pp. 471 – 488. Grimanelis, D., Neou-Syngouna, P., Vazarlis, H., 1992. Leaching of a rich Greek manganese ore by aqueous solution of sulphur dioxide. Hydrometallurgy 31, 139 – 146. Herring, A.P., Ravitz, S.F., 1965. Rate of dissolution of manganese dioxide in sulfurous acid. Trans. SME/AIME 231, 191 – 195. Higginson, W.C.E, Marshall, J.W., 1957. Equivalence change in oxidation – reduction reactions in solutions: some aspects of the oxidation of sulphurous acid. J. Chem. Soc., 447 – 458. Kanungo, S.B., Das, R.P., 1988. Extraction of metals from manganese nodules of the Indian ocean by leaching in aqueous solutions of sulphur dioxide. Hydrometallurgy 20, 135 – 146. Miller, J.D., Wan, R-Y., 1983. Reaction kinetics for the leaching of MnO2 by sulfur dioxide. Hydrometallurgy 10, 219 – 242 (and references therein). Nicol, M.J., La´zaro, I., 2002. The role of EH measurements in the interpretation of the kinetics and mechanisms of the oxidation and leaching of sulphide minerals. Hydrometallurgy 63, 15 – 22. Senanayake, G., Das, G.K., 2003. A comparative study of leaching kinetics of limonitic laterite and synthetic iron oxides in sulfuric acid containing sulfur dioxide. Hydrometallurgy, 229 – 251. Sillen, L.G., Martell, A.E., 1964. Stability constants of metal complexes. Spec. Publ. Chem. Soc. 17.
A mixed surface reaction kinetic model for the reductive leaching of manganese dioxide with acidic sulfur dioxide Gamini Senanayake * Department of Mineral Science and Extractive Metallurgy, Murdoch University, Perth, WA 6150, Australia Received 21 November 2002; received in revised form 16 September 2003; accepted 25 October 2003
Abstract Previous researchers have rationalised the kinetics of dissolution of manganese dioxide in acidic sulfur dioxide solutions on the basis of either the adsorption of reactants followed by reactions with HSO 3 and SO2 or a two-electron surface + electrochemical reaction with H+ and HSO 3 , the species involved in the equilibrium: SO2(aq) + H2O = H + HSO3 . This paper revisits the reported rate data for the dissolution of MnO2 in the form of plates [Herring and Ravitz, Trans. SME/ AIME 231 (1965) 191] or particles [Miller and Wan, Hydrometallurgy 10 (1983) 219] to show that the rate of dissolution can be rationalised on the basis of a fast redox reaction : (i) MnO2(s) + SO2(aq) + H2O = MnOOH(s) + HSO3(aq) followed by one of the two rate determining redox reactions: (ii) MnOOH(s) + HSO3(aq) = MnSO04(aq) + H2O or Mn(OH)HSO4(s), (iii) MnOOH(s) + SO2(aq) = MnO(s) + HSO3(aq) depending on the pH of the medium. Whilst the rate equation for (iii) agrees with the electrochemical kinetic model reported previously, the dissolution of the solid Mn(OH)HSO4 or the direct reaction: (iv) MnO2(s) + SO2(aq) = MnSO04(aq) also appear to be rate controlling. The two rate constants k1 and k2 based on a shrinking sphere kinetic model: 1 (1 X)1/3 = k1K 0.5[SO2]0.5(t/rq) + k2[SO2](t/rq), for the leaching of monosized pyrolusite and electrolytically prepared MnO2 particles, are in reasonable agreement with k1 and k2 based on the dissolution kinetics of MnO2 plates. D 2003 Elsevier B.V. All rights reserved. Keywords: Mixed surface reaction kinetic model; Manganese dioxide; Acidic sulfur dioxide
1. Introduction The reductive leaching of manganese dioxide with sulfur dioxide in acid media has been the subject of a number of recent research publications mainly due to the fact that SO2(aq) offers fast leaching kinetics and the choice of selective leaching of metals from limo-
* Tel.: +61-8-9360-2833; fax: +61-8-9360-6343. E-mail address: [email protected] (G. Senanayake). 0304-386X/$ - see front matter D 2003 Elsevier B.V. All rights reserved. doi:10.1016/j.hydromet.2003.10.010
nitic laterite, low-grade manganiferrous ores and deep-sea nodules (Das et al., 1997; Senanayake and Das, 2003; Abburzzese, 1990; Grimanelis et al., 1992; Kanungo and Das, 1988). Therefore a proper understanding of the reaction mechanism is useful for the optimisation of selective leaching of manganese from these materials. Herring and Ravitz (1965) rationalised the dissolution kinetics of MnO2 plates on the basis of reactions with SO2 and HSO3, the two sulfur species in equilibrium with H+ in Eq. (1). More recently Miller and Wan (1983) carried out a detailed kinetic analysis based on electrochemical and leach-
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
ing studies to propose a two-electron reaction mechanism, which involves H+ and HSO3 (Eqs. (2) –(4)). SO2 þ H2 O ¼ Hþ þ HSO 3
ð1Þ
pMnO2 ðsÞ þ 4Hþ þ 2e ¼ Mn2þ þ 2H2 O ðcathodic reactionÞ
ð2Þ
þ 2 HSO 3 þ H2 O¼3H þ SO4 þ 2e ðanodic reactionÞ
ð3Þ pMnO2 ðsÞ þ SO2 ðaqÞ ¼ Mn2þ ðaqÞ þ SO2 4 ðaqÞ ðoverall reactionÞ
ð4Þ
However, the equilibrium constants (K) reported in Table 1 show that Mn2 + forms aqueous complex species with both SO42 and SO32 ions and the magnitude of log K for MnSO30 (aq) is higher than that for MnSO40 (aq). This highlights the need to consider the distribution of manganese and sulfur species MnSO40, MnSO30, Mn2 +, HSO3, SO2, HSO4, and SO42 of different oxidation numbers Mn(II) – S(IV) –S(VI) shown in Figs. 1 and 2. Percentage of SO42 and MnSO40 (aq) increases with increasing pH
Table 1 Equilibrium constants (log K ) for reactions of S(IV) and S(VI) species with H+ and Mn2 + Reaction
Ionic strength (I )
Log K at 25 jCa
H+ + SO42 = HSO 4 H+ + HSO 3 = H2O + SO2
0 1 0
H+ + SO32 = HSO 3 Mn2 + SO42 = MnSO04(aq)
1 1.2 0 0 0 0
1.99 1.07 1.77b (1.94b at 45 jC) 1.37 6.34 2.26 (3.0 at 45 jC) 3.0 0.27c 3.61c
0 0
3.34c 5.11c
Mn2 + + SO23 = MnSO03(aq) + 0 Mn2 + HSO 4 = MnSO4(aq) + H 0 MnSO3(aq)+ HSO4 = MnSO04(aq) + HSO 3 + 0 Mn2 + + HSO 3 = MnSO3(aq) + H 2+ Mn + SO2 + H2O = MnSO03(aq) + 2H+ a
Data from Sillen and Martell (1964). Herring and Ravitz (1965) used these values of K in Eq. (9). c Derived values from the data in this table. b
Fig. 1. Effect of pH on percentage distribution of species in Mn(II) – sulfate system at 25 jC, [Mn(II)] = 0.001 M, [S(VI)] = 0.1 M, based on thermodynamic data in Table 1.
and therefore MnSO40 (aq) is the predominant species in solutions of pH > 0 and HSO4 is the predominant S(VI) species at pH < 1.5 (Fig. 1). In contrast, SO32 is unstable in acidic solutions of S(IV) and therefore the complex species MnSO30 (aq) becomes the predominant species only at pH values greater than 4; Mn2 + is predominant at low pH values (Fig. 2). The log K values derived for the reaction between Mn2 + and HSO4, SO2 or HSO3 shown in Table 1 reveal that MnSO40 is thermodynamically more favourable than MnSO30. Therefore Mn(II) in an acidic leach liquor produced by the reduction of MnO2 with SO2 is likely to be in the complex form MnSO40 (aq) rather than the uncomplexed Mn2 + (aq). Moreover, SO2 can be oxidised from the oxidation state S(IV) to the two oxidation states S(V) and S(VI) to produce S2O62 (dithionate) and SO42 ions,
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
MnðIVÞ þ SðIVÞ ¼ MnðIIÞ þ SðVIÞ
217
ð7Þ
2MnO2 ðsÞ þ 3SO2 ðaqÞ ¼ MnSO4 ðaqÞþ MnS2 O6 ðaqÞ ð8Þ
Fig. 2. Effect of pH on percentage distribution of species in Mn(II) – sulfite system at 25 jC, [Mn(II)] = 0.001 M, [S(IV)] = 0.1 M, based on thermodynamic data in Table 1.
respectively. Bassett and Parker (1951) noted that dithionate may be formed when MnO2 is reduced in one-electron steps (Eqs. (5) and (6)) whereas sulfate may be formed by a two-electron reduction mechanism (Eq. (7)), where Mn(IV) = MnO2, Mn(III) = Mn 2 O 3 (or MnOOH), S(IV) = SO 2 or HSO 3, S(V) = HSO3, H2S2O6 or S2O62 and S(VI) = HSO4 or SO42 . These two modes of reactions led to the overall reaction (Eq. (8)) between MnO2 and SO2 reported in the literature by Meyer and Schramm (Miller and Wan, 1983). MnðIVÞ þ SðIVÞ ¼ MnðIIIÞ þ SðVÞ
MnðIIIÞ þ SðIVÞ ¼ MnðIIÞ þ SðVÞ
ð5Þ
ð6Þ
The extent to which the two modes of oxidation will occur depends on the Mn/O molar ratio of the starting oxide material, i.e., the lattice structure of MnO2 (Bassett and Parker, 1951). For example, pyrolusite and MnO2 produced by the thermal decomposition of Mn(NO3)2 react with SO2 to produce mainly dithionate (95 –97%). In contrast, MnO2 prepared electrolytically (EMD) or by wet methods and dried at a moderate temperature (96 jC) produces 74– 79% dithionate. Freshly precipitated amorphous or partly crystalline MnO2, used in the wet state, produces low and variable percentages of dithionate. The rapid dissolution of amorphous MnO2 in SO2 (aq) led to the suggestion that the reduction can take place on the surface as well as in solution (Bassett and Parker, 1951). Therefore it is important to establish the mechanism of the surface reaction with respect to both one-electron and two-electron reaction mechanisms. This paper makes use of the results for the dissolution of EMD plates of known surface area (1 –3 cm2) in stirred solutions (Herring and Ravitz, 1965) to establish the mechanism of the reaction between MnO2 and SO2 over a wide range of conditions: pH 1– 4.5, total SO2 concentration 0.03 – 0.46 mol L 1 and temperature 25 –45 jC. Three different rate determining steps are considered to rationalise the effect of pH and total SO2 concentration on the initial rate of dissolution of MnO2. Two of them are used to rationalise the reported leaching results for pyrolusite and EMD in acidic sulfur dioxide in the concentration range 0.03– 1 mol L 1 at pH 2 (Miller and Wan, 1983).
2. Discussion 2.1. Effect of SO2 2.1.1. Dissolution of EMD plates Table 2 lists the different conditions used by Herring and Ravitz (1965) in three sets of experiments
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
Table 2 Materials and conditions used in kinetic studies and fraction of manganese (X ) leached Material
Test Temperature [SO2(t)]a (jC)
EMD platesb EMD platesb EMD platesb EMD particlese
A
25
0.03 – 0.46 1.8 – 1.1c
B
45
0.03 – 0.27 1.9 – 1.3c
C
25
0.06 – 0.43 4.7 – 1.0d
D
26
0.07
30
0.30 0.69 0.03
Pyrolusite E particlese
0.27 0.52 0.98
pH
Time X (s)
2d
300
0.35
2d
300 300 480
0.68 0.91 0.15
480 480 480
0.41 0.64 0.76
Total SO2 mol L 1. Dissolution of plates of EMD of constant surface area, at a stirring speed of 500 rpm (Herring and Ravitz, 1965). c Natural pH. d pH adjusted using H2SO4 or NaOH. e Monosized (357 Am) particle leaching at 0.2% solids and 650 rpm (Miller and Wan, 1983). a
b
curves in Fig. 4, becomes less significant with the increase in [SO2(t)] and pH, particularly at high pH. At a given value of [SO2(t)] = 0.2 mol L 1 (for example), the rate decreases with the increase in pH from 1.1 to 4.4. This shows the effect of the decrease in [H+] as well as the change in speciation summarized in Figs. 1 and 2 on the rate of dissolution of MnO2. 2.1.2. Dissolution of pyrolusite and EMD particles Miller and Wan (1983) reported the leaching results of monosized pyrolusite and EMD particles in the concentration range of SO2 0.03 –1 mol L 1 at 26 –30 jC and pH 2. Table 2 lists selected results from their work (Sets D and E) to show that the fraction leached (X) at a given time interval increases with increasing [SO2(t)]. At a total SO2 concentration of 0.3 mol L 1 at 26 jC, 68% of manganese is dissolved from EMD in 5 min. At a comparable total SO2 concentration of 0.27 mol L 1 at 30 jC, the percentage manganese dissolved from pyrolusite is 41%, indicating a slower reaction of SO2 with pyrolusite
(A – C ) using EMD. They used the equilibrium constant for Eq. (1): K ¼ ½Hþ HSO 3 =½SO2
ð9Þ
reported in Table 1 to calculate the natural pH and the concentration of HSO3 and ‘‘free’’ sulfur dioxide at equilibrium denoted by [SO2], in the first set of experiments (A) at 25 jC in solutions of different total concentration of sulfur dioxide [SO2(t)]. A higher temperature of 45 jC was maintained in the second set (B). In the third set (C), the pH adjustments were made by adding H2SO4 or NaOH to solutions containing different [SO2(t)]. Fig. 3 shows the equilibrium concentration of sulfur species, pH and the initial rate as a function of [SO2(t)] in Set A. The general trend is that the rate increases with increasing concentration of HSO3 and free SO2 and decreasing pH. Due to the fact that [H+], [HSO3] and [SO2] are interrelated according to Eq. (9), Fig. 4 shows the rate data in Sets A and C at four different pH values as a function of [SO2(t)]. It is clear that the dependence of reaction rate on [SO2(t)], shown by the slope of the
Fig. 3. Variation of natural pH, [HSO 3 ], [SO2] and initial rate of manganese dissolution at 25 jC (results for EMD plates from Herring and Ravitz, 1965).
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
219
(Eq. (12)) the observed slope in Fig. 5 is 0.71– 0.73 for the results in Sets A and B and at high values of [SO2] in the case of Set C. At low values of [SO2] in Set C, representing the high pH data in Fig. 4, the rate is relatively low and independent of [SO2]. These facts indicate the possibility of other surface reactions which contribute to the rate-limiting step(s). Three surface reaction mechanisms may be used to rationalise these experimental data. 2.2. Reaction schemes and mechanisms 2.2.1. Reaction scheme I at high pH (>3) Higginson and Marshall (1957) proposed the formation of HSO3 (a free radical) in a fast one-electron oxidation step in order to rationalise the kinetics of the reaction between SO2 and Fe(III) in acid media. Thus it is possible that a fast surface reaction between MnO2 and SO2 which produces MnOOH and HSO3
Fig. 4. Variation of initial rate of manganese dissolution with total SO2 concentration and adjusted pH at 25 jC (results for EMD plates from Herring and Ravitz, 1965).
compared to EMD. Miller and Wan (1983) analysed the initial rate data on the basis of the redox reactions described by Eqs. (2) and (3) and proposed the following rate equation: 0:5 0:5 R ¼ d½MnðIIÞ=dt ¼ k ½Hþ HSO ð10Þ 3 When combined with Eq. (9) this may be rewritten in the forms: R ¼ kK 0:5 ½SO2 0:5
ð11Þ
log R ¼ 0:5 log½SO2 þ log kK 0:5
ð12Þ
The general change in rate with [SO2] in Figs. 3 and 4 is consistent with Eq. (11). Although a log – log plot of R vs. [SO2] is expected to have a slope of 0.5
Fig. 5. Log – log plot of rate vs. [SO2], data from Figs. 3 and 4 and Herring and Ravitz (1965).
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
is followed by the rate determining step (RDS) in Eq. (14) to produce MnSO40 (aq) at high pH (>3).
HSO 3 ðaqÞ ¼ HSO3 ðaqÞ þ e ðanodic reactionÞ ð18Þ
pMnOOHðsÞ þ Hþ þ HSO 3
pMnO2 ðsÞ þ SO2 ðaqÞ þ H2 O ¼ pMnOOHðsÞ þ HSO3 ðaqÞðfast redox reactionÞ
¼ MnOðsÞ þ HSO3 ðaqÞðRDS 1Þ
ð19Þ
ð13Þ 2HSO3 ðaqÞ ¼ H2 S2 O6 ðaqÞ ðdimerisation reactionÞ
pMnOOHðsÞ þ HSO3 ðaqÞ ¼ MnSO04 ðaqÞ þ H2 OðRDSÞ
ð20Þ ð14Þ
Eq. (14) does not involve H+, HSO3 or SO2 and thus explains why the rate of dissolution of MnO2 remains relatively independent of [SO2] at high pH (Figs. 4 and 5), despite the fact that sulfur dioxide is predominantly in the form of HSO3 at pH values greater than 3 (Fig. 2). However, in more acidic solutions of pH < 3 it is possible to write two reaction schemes to describe the formation of dithionate and sulfate. 2.2.2. Reaction scheme II at low pH: formation of dithionate In solutions of pH less than 3, a fast surface reaction between MnO2, H+ and HSO3 produces MnOOH and HSO3 (Eqs. (15) and (16)). This is followed by one-electron redox reactions of MnOOH and HSO3 to produce MnO and HSO3 leading to the rate determining step represented in Eq. (19) (RDS-1). The dimerization of HSO3 to H2S2O6 and the subsequent reaction of H2S2O6 with MnO(s) produce MnS2O6 (aq) leading to the overall reaction in Eq. (22): pMnO2 ðsÞ þ Hþ þ HSO 3 ¼ pMnO2 Hþ HSO 3 ðfast equilibrationÞ
ð15Þ
pMnO2 Hþ HSO 3 ¼ pMnOOHðsÞ þ HSO3 ðaqÞðfast redox reactionÞ ð16Þ pMnOOHðsÞ þ Hþ þ e ¼ pMnOðsÞ þ H2 O ðcathodic reactionÞ
ð17Þ
pMnOðsÞ þ 2HSO3 ðaqÞ ðor H2 S2 O6 Þ ¼ MnS2 O6 ðaqÞ þ H2 O ðfastÞ
ð21Þ
pMnO2 ðsÞ þ 2SO2 ðaqÞ ¼ MnS2 O6 ðaqÞ ðoverall reactionÞ
ð22Þ
The rate of dissolution of manganese according to this reaction scheme is controlled by RDS-1 for which the electrochemical kinetic model (Nicol and Lazaro, 2002) described in Appendix A can be used to write the rate equation: 0:5
R1 ¼ k1 ½Hþ
HSO 3
0:5
ð23Þ
2.2.3. Reaction scheme III at low pH: formation of sulfate Instead of producing Mn(III) and S(V) via a oneelectron mechanism (Eq. (13) or (16)), MnO2 may be directly reduced to Mn(II) producing S(VI) according to Eqs. (2) and (3) (Miller and Wan, 1983). This reduction may also take place via the solid product Mn(OH)HSO4(s) instead of MnO(s) considered previously in reaction scheme II. The solid product formed in Eq. (24) is solubilised by SO2(aq) according to Eq. (25) (RDS-2). This leads to the overall reaction described by Eq. (26) (sum of (Eqs. (15), (16), (24) and (25)), which is equivalent to Eq. (4) and agrees with the species distribution discussed with reference to Fig. 1. The rate of dissolution of manganese according to this reaction scheme depends on RDS-2a and/or RDS-2b. The resulting rate equation (Eq. (27)) for both RDS-2a and RDS-2b indicates that
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
221
this reaction scheme is favoured in solutions of high concentrations of free SO2.
the basis of the two rate determining steps in Eqs. (19), (25) and (26) and may be expressed in the form R = R1 + R2;
pMnOOHðsÞ þ HSO3 ðaqÞ
R ¼ k1 ½Hþ ½HSO3 0:5 þk2 ½SO2
0:5
¼ pMnðOHÞHSO4 ðsÞ
ð24Þ
pMnðOHÞHSO4 ðsÞ þ SO2 ðaqÞ ¼
MnSO04 ðaqÞ
þH
þ
HSO 3 ðRDS
2aÞ
ð25Þ
pMnO2 ðsÞ þ SO2 ðaqÞ ¼ MnSO04 ðaqÞ ðoverall reaction or RDS 2bÞ ð26Þ R2 ¼ k2 ½SO2
ð27Þ
2.3. Overall rate equation at pH<3
ð29Þ
when combined with Eq. (9): R ¼ k1 K 0:5 ½SO2 0:5 þk2 ½SO2
ð30Þ
R½SO2 1 ¼ k1 K 0:5 ½SO2 0:5 þk2
ð31Þ
Fig. 6 shows a plot of R[SO2] 1 vs. [SO2] 0.5 for the data from Figs. 3 and 4. The two sets A and B correspond to a good linear correlation and some of the data points from set C lie close to the linear relationship for set A. The data points from set C which show the largest deviations are the same points which show rates close to 6.5 10 8 mol cm 2 s 1 and relatively independent of [SO2] and pH in Fig. 4. The slopes and intercepts of Fig. 6 along with the
The reaction schemes II and III explain the formation of dithionate and sulfate, respectively, at pH < 3. Eqs. (22) and (26) added together give Eq. (8) for the overall reaction, whilst dithionate may also decompose to sulfate: MnS2 O6 ðaqÞ ¼ MnSO04 ðaqÞ þ SO2 ðaqÞ
ð28Þ
The extent of oxidation of SO2 to sulfate or dithionate will depend on the relative rates of RDS-1, RDS2a and RDS-2b in Eqs. (19), (25) and (26). Solutions of lower pH values correspond to high [SO2] (Figs. 2 and 3) and favour reaction scheme III (Eqs. (25) and (26)) over II (Eq. (19)). This is consistent with the observation reported in the literature (Miller and Wan, 1983) that the dithionate formation decreased as the pH was lowered. Furthermore, the observation that the dithionate formation decreased rapidly with the increased rate of agitation (Miller and Wan, 1983) indicates that higher agitation rates would favour reaction schemes I and III which do not form dithionate. The mass transfer of uncharged species HSO3 and SO2 involved in Eqs. (14) and (26) to the reaction surface is likely to be facilitated by the high rates of agitation. However, the contribution of reaction scheme I to the overall rate may be ignored at pH values lower than 3. Thus the overall rate equation for the dissolution of MnO2 at pH < 3 may be derived on
Fig. 6. Correlation between R[SO2] 1 vs. [SO2] 0.5: testing the validity of Eq. (31), data from Figs. 3 – 5.
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G. Senanayake / Hydrometallurgy 73 (2004) 215–224
calculated rate constants k1, k2 (see Eq. (31)) and the ratio k1/k2 listed in Table 3 show that the rate constant k1 is twice as large as k2. The activation energy for RDS-1 is 24 kJ mol 1 compared to 23 kJ mol 1 for RDS-2. These values are slightly larger than the reported value of 19 kJ mol 1 (Herring and Ravitz, 1965). 2.4. Application to leaching of monosized particles
the overall rate of leaching. This may be further examined by combining the overall rate equation (Eq. (30)) with the general equation for shrinking sphere model: 1 ð1 X Þ1=3 ¼ k ½SO2 ðt=rqÞ
where r (cm) is the initial particle radius and q (mol cm 3) is the concentration of MnO2 in the solid. The combined equation:
Miller and Wan (1983) showed that their results for the leaching of pyrolusite and EMD obeyed the rate equation corresponding to a shrinking sphere model:
1 ð1 X Þ1=3 ¼ k1 K 0:5 ½SO2 0:5 ðt=rqÞ
1 ð1 X Þ1=3 ¼ k ½SO2 n t
may be rearranged:
ð32Þ
where X = fraction of Mn reacted at time t, n = reaction order with respect to SO2 and k = apparent rate constant. This supports the view that the kinetics of leaching is controlled by the surface reactions alone and the rate decreases with decreasing surface area of the spherical particles with no product layer to retard the diffusion of reactants or products. However, the fact that their results at pH 2 showed a reaction order of 0.58 or 0.59 for the pyrolusite and EMD, respectively, indicates that both RDS-1 and RDS-2 control
ð33Þ
þ k2 ½SO2 ðt=rqÞ
ð34Þ
f1 ð1 X Þ1=3 g½SO2 1 t 1 ¼ k1 K 0:5 ðrqÞ1 ½SO2 0:5 þk2 ðrqÞ1
ð35Þ
Fig. 7 shows a plot of {1 (1 X)1/3}[SO2] 1 t 1 vs. [SO2] 0.5 and confirms the validity of this mixed kinetic model for the leaching of monosized pyrolusite and EMD. The slopes and intercepts depend on the value of K at the appropriate ionic strength (I) used to calculate [SO2]. However, good linear correlations with R2 > 0.99 are observed with the values of K at
Table 3 Rate constants (k1 and k2) obtained for different materials and conditions calculated from slopes and intercepts of Figs. 6 and 7 using K at ionic strengths 0 and 1 Material/test
T (jC)
Slope
Intercept
25 45
0.74e 1.11e
2.67e 4.76e
2.5 10 3g (3.5 10 3)
2.1 10 3g (4.1 10 3)
9.0 10 4g (12 10 4)
8.0 10 4g (16 10 4)
104 rqa (mol cm 2)
103 k1b (cm s 1)
103 k2c (cm s 1)
k1/k2
5.7 10
2.7 4.8
2.1 2.2
9.4
18 (16)
2.0 (3.9)
9.0 (4.1)
9.3
6.4 (5.4)
0.7 (1.5)
9.1 (3.6)
d
EMD plates A B
EMD particlesf D 26
Pyrolusite particlesf E 30
q = wt.% MnO2 density of the material (g cm 3)/molar mass of MnO2 (g mol 1) (Miller and Wan, 1983). Plates: k1 = 10 3 K 0.5 slope; particles: k1 = 103 K 0.5 rq slope; K 0.5 = 7.67 (I = 0, 25 – 30 jC), 9.33 (I = 0, 45 jC), 4.84 (I = 1, 25 – 30 jC); values in parentheses are based on K at I = 1 (see Table 1). c Plates: k2 = 10 3 intercept; particles k2 = 103 rq intercept. d See footnote in Table 2. e Fig. 6 (Eq. (31)). f See footnote in Table 2. g Fig. 7 (Eq. (35)). a
b
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
223
nature of MnO2. Nevertheless, the comparison of k values for the dissolution of MnO2 plates and particles provides useful insight into the rate-limiting step(s).
3. Conclusion
Fig. 7. Correlation between {1 (1 X)1/3}[SO2] 1 t 1 vs. [SO2] 0.5: testing the validity of Eq. (35); conditions and data listed in Table 2.
both ionic strengths I = 0 and 1. Table 3 lists and compares the values of k1 and k2 obtained from the slopes and intercepts of Fig. 7. It is clear that the use of K values at different ionic strengths does not make a significant difference to the rate constants. The k1 value for pyrolusite leaching at 30 jC shows good agreement with that for the dissolution of EMD plates at 25 jC; but k1 for EMD leaching is nearly three times larger than that for pyrolusite. The values for k2 are also in good agreement for the three cases within a factor of 1 – 4 at 25 – 30 jC. The ratio k1/k2 for leaching is in the range 3.6– 9.1 compared to 2.1 for the dissolution of EMD plates. These small differences in rate constants k1 and k2 support the view of Bassett and Parker (1951) that the extent of reduction of MnO2 with SO2 in acid media (pH < 3) according to the two reaction schemes II and III forming dithionate and sulfate depends on the crystalline
Kinetics of reductive dissolution of manganese dioxide in acidic sulfur dioxide media (pH < 3) is controlled by different surface reaction schemes depending on the conditions. In the first scheme a fast one-electron surface reaction produces MnOOH and HSO3. This is followed by another one-electron reaction leading to the production of MnS2O6 and the RDS corresponds to a reaction order of 0.5 with respect to H+ and HSO3. However, MnS2O6 may decompose to MnSO4 and SO2. In the second reaction scheme, SO2 may directly react with MnO2 or two one-electron reactions may produce an insoluble product Mn(OH)HSO4 on the surface. The rate of dissolution of manganese in these cases appears to be first order with respect to SO2 in equilibrium with H+ and HSO3. Thus the overall rate equation for leaching can be expressed by a mixed surface reaction kinetic model combined with a shrinking sphere model which incorporate the two rate constants k1 and k2 : 1 (1 X)1/3= k1K0.5[SO2]0.5 (t/rq)+ k2[SO2] (t/rq).
Appendix A The rate of an electrochemical reaction can be related to the current i, the number of electrons n involved in the reaction and the Faraday constant F: R ¼ i=nF
ð36Þ
According to the electrochemical model (Nicol and La´zaro, 2002) the rate of the two electrochemical reactions in Eqs. (17) and (18) can thus be represented by the cathodic and anodic currents ic and ia : ic ¼ kc ½Hþ expðac FEm =RT Þ
ð37Þ
ia ¼ ka HSO 3 expðaa FEm =RT Þ
ð38Þ
224
G. Senanayake / Hydrometallurgy 73 (2004) 215–224
where ac and aa are the transfer coefficients, and kc and ka are electrochemical rate constants per unit surface area, and Em is the mixed potential. At the mixed potential Em, jicj = ia = imix and since ac = ac = 0.5, the two equations can be combined: ic ia ¼ ðimix Þ2 ¼ ka kc ½Hþ HSO 3
ð39Þ
to express the rate of dissolution of MnO2 per unit area according to the one-electron mechanism (N = 1 in eq. (36)): 0:5 1 R1 ¼ imix F 1 ¼ ðka kc ½Hþ HSO F 3 Þ
R1 ¼ k1 ½Hþ
0:5
HSO 3
0:5
ð40Þ
ð23Þ
where k1=(kakc)0.5 F 1 represents the rate constant of RDS-1 (Eq. (19)).
References Abbruzzese, C., 1990. Percolation leaching of manganese ore by aqueous sulfur dioxide. Hydrometallurgy 25, 85 – 97.
Bassett, H., Parker, W.G., 1951. The oxidation of sulphurous acid. J. Chem. Soc., 1540 – 1560. Das, G.K, Anand, S., Das, R.P., Muir, D.M., Senanayake, G., Singh, P., Hefter, G., 1997. Acid leaching of nickel laterites in the presence of sulfur dioxide at atmospheric pressure. In: Cooper, W.C., Mihaylov, I. (Eds.), Hydrometallurgy and Refining of Nickel and Cobalt, vol. 1. Canadian Inst. Min. Metall. and Petroleum, Montreal, pp. 471 – 488. Grimanelis, D., Neou-Syngouna, P., Vazarlis, H., 1992. Leaching of a rich Greek manganese ore by aqueous solution of sulphur dioxide. Hydrometallurgy 31, 139 – 146. Herring, A.P., Ravitz, S.F., 1965. Rate of dissolution of manganese dioxide in sulfurous acid. Trans. SME/AIME 231, 191 – 195. Higginson, W.C.E, Marshall, J.W., 1957. Equivalence change in oxidation – reduction reactions in solutions: some aspects of the oxidation of sulphurous acid. J. Chem. Soc., 447 – 458. Kanungo, S.B., Das, R.P., 1988. Extraction of metals from manganese nodules of the Indian ocean by leaching in aqueous solutions of sulphur dioxide. Hydrometallurgy 20, 135 – 146. Miller, J.D., Wan, R-Y., 1983. Reaction kinetics for the leaching of MnO2 by sulfur dioxide. Hydrometallurgy 10, 219 – 242 (and references therein). Nicol, M.J., La´zaro, I., 2002. The role of EH measurements in the interpretation of the kinetics and mechanisms of the oxidation and leaching of sulphide minerals. Hydrometallurgy 63, 15 – 22. Senanayake, G., Das, G.K., 2003. A comparative study of leaching kinetics of limonitic laterite and synthetic iron oxides in sulfuric acid containing sulfur dioxide. Hydrometallurgy, 229 – 251. Sillen, L.G., Martell, A.E., 1964. Stability constants of metal complexes. Spec. Publ. Chem. Soc. 17.