Hydrolysis of Ferric Sulfate in the Presence of Zinc Sulfate at 200Â° C

Ind. Eng. Chem. Res. 2004, 43, 6299-6308

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Hydrolysis of Ferric Sulfate in the Presence of Zinc Sulfate at 200 °C: Precipitation Kinetics and Product Characterization Terry C. Cheng and George P. Demopoulos* Department of Metals and Materials Engineering, McGill University, Montreal, Quebec H3A 2B2, Canada

The hydrolysis of ferric sulfate was studied in a batch reactor at 200 °C with the initial Fe(III) concentration ranging from 0.10 to 0.80 M in the presence/absence of zinc sulfate (1.2 M ZnSO4) with 3 h retention time. Regardless of the presence or not of ZnSO4, the hydrolytic precipitation kinetics at 200 °C was determined to be first-order with respect to Fe(III), and expressed by the following rate equation, rhyd ) khyd(CFe(III) - CFe(III),eq), where the apparent kinetic constant, khyd, was determined to be 10-2 min-1. Two reaction stoichiometries were established: one leading to the formation of kinetically favored basic ferric sulfate at CFe(III),initial g 0.4 M in the absence of ZnSO4 or at CFe(III),initial g 0.7 M in the presence of ZnSO4, while the other led to the production of hematite. The hydrothermally produced hematite material possessed very high specific surface area (50-80 m2/g) and contained around 1% S (as SO4 via chemisorption), 0.2% Zn (when ZnSO4 was present), and about 4.5% H2O/OH. After excluding the SO4 and Zn content, the following stoichiometric formula for hematite was determined: Fe2O3-x(OH)2x‚yH2O where x ) 0.08 and y ) 0.29. 1. Introduction

steps: (i) oxidation of ferrous sulfate to ferric sulfate10

The hydrolysis of ferric sulfate at elevated temperatures and pressures is of interest to those involved, either from operational or research points of view, with the hydrometallurgical processing of ores or concentrates in autoclaves. For example, hydrolytic iron precipitation is a critical reaction in the pressure leaching of zinc sulfide concentrates,1 pressure oxidation of refractory gold ores,2 and the high-pressure acid leaching of laterite ores.3 The level of iron concentration in solution and the type and properties of iron(III) precipitate forming impact directly on the performance of the process. But of particular importance is process of the hydrolysis of ferric sulfate to the hematite, which is practiced in the zinc industry.4,5 This process, which involves production of hematite out of iron-rich zinc sulfate solutions at 200 °C, is the only commercial iron precipitation process6 that does not yield a waste material. The hematite product is relatively clean and as such finds applications in the cement-manufacturing industry.7 The hematite process, however, does not compete well economically with the other iron removal technologies.8 Presently, Akita Zinc Co. of Japan9 is the sole industrial operator of such a process. For the process to become widely accepted, its retention time (currently standing at around 3 h)9 needs to be shortened in order to lower its capital cost. For this to be achieved, the kinetics of the process has to be established. This is, indeed, the object of this paper. The hematite process involves oxydrolysis of ferrous sulfate at an oxygen partial pressure of 150-500 kPa and a temperature of 200 °C:4,7

2FeSO4(aq) + 1/2O2(g) + H2SO4(aq) f Fe2(SO4)3(aq) + H2O(l) (2)

2FeSO4(aq) + 1/2O2(g) + 2H2O(l) f Fe2O3(s) + 2H2SO4(aq) (1) The oxydrolysis reaction is the sum of two reaction * To whom correspondence should be addressed. E-mail: [email protected].

and (ii) hydrolysis of ferric sulfate to produce hematite11,12

Fe2(SO4)3(aq) + 3H2O(l) f Fe2O3(s) + 3H2SO4(aq) (3) Alternatively to reaction 3, the hydrolysis of ferric sulfate may lead to the formation of basic ferric sulfate (BFS),13,14 which is undesirable15 as a final product:

Fe2(SO4)3(aq) + 2H2O(l) f 2Fe(OH)SO4(s) + H2SO4(aq) (4) The study of the hydrothermal precipitation equilibria of the Fe2O3-SO3-H2O system has attracted the attention of several researchers. Thus, following the pioneering work of Posnjak and Merwin,16 the Japanese researchers Sasaki, Tozawa, and Umetsu have perhaps carried out the most extensive equilibrium studies.13,17-21 According to these studies, the stability region of hematite is defined by a critical value of free acid concentration, [H2SO4]cr, below which hematite is stable and above which basic ferric sulfate is stable. Umetsu et al.13 and Sasaki et al.19-21 determined the [H2SO4]cr values in the temperature range between 150 and 220 °C in the presence of a variety of metal sulfate salts (ZnSO4, CuSO4, MgSO4, and Na2SO4). They concluded that the stability region of hematite as a function of free acid concentration is enlarged with temperature elevation and/or with the addition of one of the metal sulfate salts stated above, except for Na2SO4, which leads to the formation of sodium jarosite instead.21 At 200 °C, [H2SO4]cr was found to be 65 g/L (0.66 M) in the absence of any metal sulfate salt and 99 g/L (1.0 M) in the presence of zinc sulfate (72-75 g/L Zn2+).13 The positive effect that the sulfate salts have on enlarging the

10.1021/ie030711g CCC: $27.50 © 2004 American Chemical Society Published on Web 09/02/2004

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stability of hematite was explained by making reference to the SO4/HSO4- equilibrium.17,18 The addition of sulfate salts decreases the free H+ concentration (or better the proton activity), which governs the solubility of hematite, thus resulting in the enlargement of the hematite stability region. However, Voigt and Go¨bler14 found that in sulfate media containing 1.0 M initial Fe(III) concentration and zero initial free acid, hydronium jarosite [H3OFe3(SO4)2(OH)6] precipitated preferentially at temperatures up to 200 °C, above which the formation of basic ferric sulfate was favored. The formation of hematite was observed (as minor phase) only at a temperature g 250 °C (12-35 wt % depending on the temperature and reaction time). Apparently, both hydronium jarosite and basic ferric sulfate are metastable in nature relative to hematite, as they were found to convert to the latter with an increase in temperature and/or reaction time in the following order:14 >150 °C

hydronium jarosite 98 >200 °C

basic ferric sulphate 98 hematite (5) Another interesting observation from Voigt and Go¨bler’s study14 is that the precipitation yield was relatively low; e.g., only 51% iron precipitation yield was achieved after 24 h at 200 °C, yet the final solid product was merely basic ferric sulfate. On the basis of the stoichiometry of reaction 4, the final free acid concentration is estimated to be 24.5 g/L (0.25 M), which is well within the stability region of hematite. Similar observations can be made in the case of the work of Dutrizac and Chen,12 who conducted hydrolysis tests with various initial Fe(III) concentrations at 225 °C for 4 h. These researchers found hematite to form without basic ferric sulfate coprecipitation only at initial Fe(III) concentrations below 0.25 M; at higher concentrations basic ferric sulfate formed and became the principal phase when the initial Fe(III) was above 0.40 M. Once more, an estimation of the final acid concentration shows that the formation of basic ferric sulfate occurred within the stability region of hematite. This is a good manifestation of Stranski’s rule (or Ostwald’s step rule).22,23 According to this rule, when homogeneous nucleation occurs at high supersaturation, i.e., high ferric sulfate concentration, the formation of the least stable phase (in this case, basic ferric sulfate) is favored: the latter upon time prolongation transforms gradually to the most stable phase (hematite). Despite the numerous equilibrium studies13,14,16,19-21 on the hydrolytic precipitation of Fe(III), no information on the kinetics of the process has been published, especially in the presence of zinc sulfate. This work focuses on investigating the hydrolysis of ferric sulfate in terms of its precipitation kinetics and solid product characterization so as to shed light on the overall mechanism of the hematite precipitation process. The study was conducted in the presence/absence of zinc sulfate at 200 °C, the operating temperature of the industrial hematite process.4,5 2. Experimental Section 2.1. Experimental Procedure. A 2.0 L titanium (ASTM grade 2) stirred pressure reactor manufactured by Parr Instrument Co. was used. Known amounts of reagent-grade ferric sulfate pentahydrate were dissolved into acidified deionized water. The initial Fe(III) concentration used varied from 0.10 to 0.80 M (5.6-45 g/L)

while the initial free sulfuric acid concentration was 5 g/L for all experiments. In experiments where zinc sulfate was present, 1.2 M reagent-grade zinc sulfate heptahydrate (i.e., 80 g/L Zn) was used. Both the initial free acid concentration and zinc sulfate concentration chosen were such to simulate the respective composition in the industrial feed solution.4 After the salts were completely dissolved, 1.0 L of the prepared solution was transferred into the autoclave glass liner. After assembly of the autoclave, the content was heated to 200 °C with no agitation applied. As soon as the temperature reached 200 °C (i.e., t ) 0), agitation was applied at 500 rpm. The standard retention time of each experiment was set at 180 min, except in a series of experiments conducted with the initial Fe(III) concentration of 0.7 M, where retention times of 0, 40, 120, and 180 min were used. During the course of each experiment, a single sample of 8 cm3 was taken at a designated time interval (0, 5, 10, 20, 40, 120, and 180 min). All samples were withdrawn from the vessel via a dip tube connected to a sample collector (∼10 cm3) immersed in a cold bath after systematic flushing of the sample collector. After each sample was collected, it was immediately filtered by a 13 mm syringe-type filter cap of 0.2 µm pore size. The solution sample was analyzed for iron concentration using a classical titration method for total iron determination24 as well as free acid concentration using the standard EDTA-assisted titration method.25 At the designated total retention time, agitation was stopped, and the slurry was allowed to cool naturally to approximately 50 °C. The cooled slurry was filtered by using a filter press with a membrane of 0.2 µm pore size. The filter cake was washed four times with hot (∼80 °C) deionized water with a volume equal to the amount of filtrate collected and dried in an oven at 50 °C for 24 h; there was no further weight loss with longer drying time. The solid products were characterized by using an array of instrumental analyses26,27 starting with the X-ray powder diffraction method (XRD) for phase identification and electron microprobe analysis for distribution of major elements (i.e., Fe, O, and S), followed by Fourier transform infrared (FTIR) spectroscopy for SO4 grouping analysis. The composition of the solid products in terms of Fe and S were determined using a X-ray fluorescence spectrometer (XRF) and a Leco combustion analyzer, respectively. In addition, the specific surface area of the solid products was measured by the N2-BET method.28 3. Results and Discussion 3.1. Analysis of Kinetic Data. It was observed during the course of this study that a significant amount of the initial Fe(III) precipitated out of the solution during the heat-up period (∼50 min), regardless of the presence or not of zinc sulfate. The extent of iron precipitation during the heat-up period was more profound at lower initial ferric concentration. For example, in the presence of 80 g/L of zinc as zinc sulfate, the extent of precipitation during the heat-up period was over 96% for CFe(III),initial ) 0.10 M, but only around 20% for CFe(III),initial ) 0.80 M. To reduce the uncertainty arising from the low-end Fe(III) concentration experiments, it was decided to exclude the experiments done at CFe(III),initial ) 0.10 M from the kinetic analysis. Figure 1 shows a set of Fe(III) concentration-time profiles for various initial Fe(III) concentrations as a function of time obtained at 200 °C. By defining time

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Figure 1. Fe(III) concentration-time profiles resulting from the hydrolysis of ferric sulfate solutions at 200 °C in the presence of ZnSO4. The curves were obtained by fitting the experimental data into eq 6, where the fitting parameters are given in Table 1.

zero (t ) 0) at the moment when the content in the autoclave reached 200 °C, the Fe(III) concentration in solution profiles resulting from hydrolysis can be expressed as

CFe(III)(t) ) A exp(-Bt) + CFe(III),eq,∞

(6)

where CFe(III),eq,∞ is a constant corresponding to the equilibrium concentration at infinite time. CFe(III),eq,∞ was estimated by using the final free sulfuric acid concentration (after 3 h retention time) via the following empirical correlations (valid at 200 °C) that were derived from Umetsu et al.’s data.13 In the absence of zinc sulfate,

log CFe(III),eq ) 3.73 log CH2SO4 - 0.819

(7)

and in the presence of zinc sulfate (68-101 g/L Zn),

log CFe(III),eq ) 5.34 log CH2SO4 - 1.20

(8)

where the concentration terms are in moles per liter. Using the least-squares regression method,29 the parameters A and B of eq 6 were determined for each Fe(III) concentration profile and are reported in Table 1. In analogy to the power law expression used in precipitation/crystallization kinetics,30 the hydrolytic reaction rate of ferric sulfate is expressed here as

rhyd ) khyd(CFe(III) - CFe(III),eq)β

(9)

where rhyd is in mol‚L-1‚min-1, while khyd (in min-1) and β are the apparent kinetic constant and reaction order, respectively. The experimental data were evaluated by using the differential method31 with the initial rates, rhyd,0, obtained from the first derivative of eq 6:

rhyd,0 ) -

|

dCFe(III) dt

t)0

) AB

(10)

By taking logarithms on both sides of eq 9 and evaluating at t ) 0, eq 11 is obtained

log rhyd,0 ) log khyd + β log(CFe(III),0 - CFe(III),eq,0) (11)

Figure 2. log-log kinetic plots of the hydrolysis of Fe2(SO4)3 at 200 °C in the presence/absence of ZnSO4. Table 1. Obtained Fitting Parameters, A and B, in Equation 6 for the Hydrolysis Experiments Conducted in the Presence/Absence of Zinc Sulfate with Various Initial Fe(III) Concentrations CFe(III),initial (M) 0.2 0.3 0.4 0.5 0.7 0.8 0.2 0.3 0.4 0.5 0.7 0.8

A In the Absence of Zinc Sulfate 2.08 × 10-2 1.10 × 10-1 1.57 × 10-1 1.66 × 10-1 2.01 × 10-1 7.29 × 10-1

B 1.36 × 10-2 7.31 × 10-3 5.19 × 10-3 5.99 × 10-3 6.38 × 10-3 4.34 × 10-3

In the Presence of Zinc Sulfate (80 g/L Zn) 1.54 × 10-2 2.83 × 10-2 6.00 × 10-2 3.33 × 10-2 1.29 × 10-1 2.62 × 10-2 2.61 × 10-1 2.58 × 10-2 4.76 × 10-1 2.22 × 10-2 5.94 × 10-1 2.33 × 10-2

where CFe(III),0 and CFe(III),eq,0 are the respective ferric concentration in solution and the equilibrium concentration at t ) 0. By plotting log rhyd,0 versus log(CFe(III),0 - CFe(III),eq,0) for T ) 200 °C, a straight line is obtained with the slope equal to the reaction order, β, while its y intercept is the logarithm of the apparent kinetic constant, log khyd. 3.2. Precipitation Rate Equation. With the use of eqs 10 and 11, the reaction order, β, for the hydrolysis of ferric sulfate in the presence or not of zinc sulfate was determined at 200 °C; the log-log plots are shown in Figure 2. From the experiments done in the presence of zinc sulfate, β was found to be 0.917 with a coefficient of determination, R2, equal to 0.994, while in the absence of zinc sulfate, β was found to be 0.878 with R2 equal to 0.890. The higher R2 value from the former indicated a better regression fit of the experimental data, yet for the sake of simplicity, the β value of unity was adopted. (Since the equilibrium concentration term, CFe(III),eq, in eq 9 varies with time, a simple integration solution is not available to verify the reaction order.) By fixing the reaction order, the apparent kinetic constants, khyd, for the hydrolysis reaction in the presence and absence of zinc sulfate were determined, respectively. The logarithmic values of the apparent

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Figure 3. Relationship between final Fe(III) and free acid concentrations obtained from the hydrolysis experiments in the presence/absence of ZnSO4 at 200 °C for a retention time of 180 min; the solubility data (*) are extracted from Umetsu et al.13

kinetic constant, log khyd, were found to be -1.59 ( 1.06 in the presence of zinc sulfate (80 g/L Zn) and -2.11 ( 1.20 in the absence of zinc sulfate. Because of the error associated with the determination of khyd, a single value, khyd ) 10-2 min-1, was adopted to be included in the following rate equation irrespective of the presence or not of zinc sulfate:

rhyd[mol‚L-1‚min-1] ) 10-2(CFe(III) - CFe(III),eq) (12) 3.3. Evaluation of Reaction Stoichiometries. As it would be expected from the iron(III) hydrolysis reactions (reactions 3 and 4), acid is generated as a result of the hydrolysis of ferric sulfate. The relationship between the final Fe(III) and the final free sulfuric acid concentrations observed in the 3 h hydrolysis experiments is given in Figure 3. Figure 3 shows that in both series of experiments (in the presence/absence of zinc sulfate), the final Fe(III) concentration in solution was always above the respective equilibrium curves (i.e., equilibrium conditions had not been reached in the employed retention time of 180 min). The final Fe(III) and acid concentration data of all hydrolysis tests were analyzed according to the stoichiometries of the two hydrolysis reactions (reactions 3 and 4) by calculating the moles of iron precipitated, XFe(III)(s), and acid generated, CH2SO4,generated as shown in Figure 4. Despite the fact that no true equilibrium had been reached, the observed ratio CH2SO4,generated/XFe(III)(s) (the slope) was close to 1.5 for the experiments done in the presence of ZnSO4 with initial Fe(III) concentration less than 0.7 M; however, as the initial Fe(III) concentration increased above 0.7 M, a change in slope to 0.5 was observed, hence indicating the formation of basic ferric sulfate, in lieu of hematite. This observation is in accordance with the product characterization data reported in the next section. In the experiments done

Figure 4. Relationship between the amount of Fe(III) precipitated and acid generated.

in the absence of ZnSO4, the slope of 0.5 was observed at lower initial Fe(III) concentration (g0.4 M). On Figure 4, the critical H2SO4 concentration levels marking the stability region of hematite and basic ferric sulfate at 200 °C (in the absence or presence of ZnSO4) as determined by Umetsu et al.13 are indicated as well. The results from this work behave in general as predicted by Umetsu et al.’s work, with the exception that the shift from one region of stability to the other one occurs at a bit lower [H2SO4] than the respective [H2SO4]cr. This implies that within the relatively short time of the present experiments (3 h), some basic ferric sulfate did not convert to the more stable hematite product observed in the long equilibrium experiments of Umetsu et al.13 The relationship between the amount of iron precipitated (% Fe removal) and the amount of acid generated as a function of retention time was followed as well for the case of an initial Fe(III) concentration of 0.70 M (39 g/L). The results are shown in Figure 5. According to these data, after 50 min of heat-up period, about 25% of the initial iron had precipitated out of solution; as retention time increased, both the % Fe removal and acid concentration were found to increase. In the same figure, the molar ratio between the amount of acid generated and the amount of iron(III) precipitated are plotted as a function of time. The ratio appeared to increase from around 0.6 at t ) 0 to nearly 1.5 at t > 120 min. This suggests that initially basic ferric sulfate forms (because of its apparently favorable homogeneous nucleation kinetics) followed by the conversion of it to hematite. This observation is in agreement with the comments made in the previous paragraph with reference to [H2SO4]cr and is further validated by the evolution of the solid product composition with time as reported in the next section. 3.4. Characterization of Solid Products. The composition of the solid products obtained from the 3 h hydrolysis tests in the presence/absence of zinc sulfate


Figure 5. Effect of retention time on the amount of Fe(III) precipitated and acid generated from the hydrolysis of Fe2(SO4)3 in the presence of ZnSO4 at 200 °C with initial Fe(III) concentration of 0.7 M. Table 2. Composition of the Solid Products Produced from the 3 h Hydrolysis Tests in the Presence/Absence of Zinc Sulfate with Various Initial Fe(III) Concentrations CFe(III),initial (M)

% Fe

%S

% Zn

% H2O/OH

0.2 0.3 0.4 0.5 0.7 0.8

In the Absence of Zinc Sulfate 63.3 1.3 63.3 1.7 62.5 2.4 54.4 12.2 37.3 15.4 29.5 17.1

4.6 4.3 4.9 8.5 9.3 14.7

0.1 0.2 0.3 0.4 0.5 0.7 0.8

In the Presence of Zinc Sulfate (80 g/L Zn) 64.8 1.3 0.2 65.9 0.8 0.3 64.9 0.9 0.2 64.4 1.2 0.2 63.8 1.3 0.2 62.0 2.0 0.2 62.0 3.8 0.3

4.2 4.8 4.2 4.5 4.1 4.7 5.3

with various initial Fe(III) concentrations is reported in Table 2. The results show that the precipitates produced in the presence of zinc sulfate contained 0.2-0.3% Zn and around 4-5% H2O/OH content irrespective of the initial Fe(III) concentration. On the other hand, the H2O/OH content of the precipitates produced in the absence of zinc sulfate increased dramatically (from 5 to 9%) as the initial Fe(III) concentration increased from 0.4 to 0.5 M. As discussed by Schwertmann and Cornell26,27 and Dutrizac and Chen,32 at least part of the “water” is OH contained in the structure of hematite as O2- substituent. The average composition of hematite obtained from solution containing ZnSO4 and CFe(III),initial from 0.1 to 0.4 M was estimated to be 65 ( 0.6% Fe, 1.1 ( 0.2% S, 0.2 ( 0.05% Zn, and 4.4 ( 0.3% H2O/OH. By excluding the S and Zn contents, this gives 96.6% hematite content. On the 100% hematite basis, the elemental composition was determined to be 67.3% Fe, 28.1% O, and 4.58% H2O/OH. The theoretical composition of R-Fe2O3 is 69.9% Fe and 30.1% O. By assigning part of the H2O content as OH (this determined on the

Figure 6. XRD analysis of ferric sulfate hydrolysis products obtained at various initial Fe(III) concentrations in the absence of zinc sulfate.

basis of charge balance considerations) and the rest as crystallization water, the following chemical formula was deduced: Fe2O3-x(OH)2x‚yH2O, where x ) 0.08 and y ) 0.29. In concordance with the data reported in the previous section, XRD analysis showed that the production of both hematite (R-Fe2O3) and basic ferric sulfate or the mixture of the two depends on the presence or not of zinc sulfate and the level of the initial ferric iron concentration. The formation of hematite was found at low initial ferric iron concentration and in the presence of zinc sulfate. Qualitative XRD evidence of the transition from the production of hematite to basic ferric sulfate was detected at CFe(III),initial g 0.4 M in the absence of zinc sulfate and at CFe(III),initial g 0.7 M in the presence of zinc sulfate (refer to Figures 6 and 7). Chemical analysis of the solid products in terms of % Fe and % SO4 confirmed the trends indicated by XRD analysis. Thus, as can be seen in Figure 8 in the case of zinc sulfate being absent, a drastic change in solid product composition occurred at CFe(III),initial between 0.4 and 0.5 M (iron content decreased from 62.5 to 54.4% while sulfate content increased from 7.19 to 36.5%), indicating an abrupt transition from a predominantly hematite production regime to a predominantly basic ferric sulfate production regime. Similarly, in the presence of zinc sulfate, the solid product composition results further confirmed what the XRD data showed, that the precipitate produced with CFe(III),initial e 0.7 M was predominately hematite (% Fe > 62% and % SO4 < 6%). Included in Figure 8 are also the experimental data of Dutrizac and Chen,12 who studied the hydrolysis of ferric sulfate at 225 °C for 4 h (240 min) with no zinc sulfate present and zero initial free acid. The results of the present study (in the absence of zinc sulfate) are in good agreement with the general trend of the results of Dutrizac and Chen.12

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Figure 7. XRD analysis of ferric sulfate hydrolysis products obtained at various initial Fe(III) concentrations in the presence of zinc sulfate.

Figure 9. % S content of the precipitates produced in 3 h hydrolysis tests at 200 °C in the presence/absence of ZnSO4 (80 g/L Zn) as compared to the respective stability regions (defined by the dash lines) of hematite and basic ferric sulfate given by Umetsu et al.13

Figure 10. Effect of retention time on the composition of the hydrolysis products (same experimental conditions as those reported in Figure 5). Figure 8. Effect of initial Fe(III) concentration on the solids composition along with the respective data (*) extracted from Dutrizac and Chen.12

Next, the sulfur content (% S) of the precipitates obtained in the present study after 3 h hydrolysis is plotted as a function of final free acid concentration ([H2SO4]final) and compared to the hematite/basic ferric sulfate stability regions (equilibrium data) reported by Umetsu et al.13 (see Figure 9). The precipitates obtained in the present study, regardless of the presence or not

of zinc sulfate, were found to correspond to the solid products obtained at equilibrium despite that the final Fe(III) concentration in solution data (see Figure 3) showed that equilibrium was not reached after 3 h of hydrolysis. In other words, it appears that while stable hematite has been produced, the remaining solution is still supersaturated with reference to hematite. It is interesting to consider the evolution of product composition with time for the tests conducted with 0.7 M initial Fe(III) concentration in the presence of zinc


basic ferric sulfate.) As retention time increased, the % Fe content of the precipitate increased while the % SO4 content was reduced gradually to 10% after 2 h and 6% after 3 h. It is apparent that basic ferric sulfate (BFS) formed in the early stages of hydrolysis is metastable in nature. The formation of metastable BFS under the present conditions is believed to be the result of Stranski’s rule,22,23 which stipulates that under conditions of homogeneous nucleation (or equivalently high supersaturation), metastable products form ahead of the stable one. With the progress of the reaction, the metastable basic ferric sulfate converts to hematite. Such transformation (reaction 13) has been observed as well by Voigt and Go¨bler.14 However, in their case the conversion kinetics was slower, reflecting apparently the higher initial ferric iron content (1 M) and consequently higher acid concentration buildup:

Fe(OH)SO4(s) + H2O(l) f Fe2O3(s) + 2H2SO4(aq) (13) Figure 11. FTIR spectra of hematite precipitates obtained from (a) hydrolysis of Fe2(SO4)3 (CFe(III),initial ) 0.2 M) and from (b) hydrolysis of Fe(NO3)3 following the procedure prescribed by Shang and Van Weert (CFe(III),initial ) 0.8 M);33 (c) the spectrum of basic ferric sulfate precipitate (produced from hydrolysis of Fe2(SO4)3 in the absence of zinc sulfate; CFe(III),initial ) 0.8 M).

sulfate. The respective composition data are shown in Figure 10. As can be seen at the end of the heat-up period (t ) 0), the precipitate had a composition of 52.3% Fe and 18% SO4; i.e., clearly a mixture of hematite and basic ferric sulfate. (A simple mass balance calculation suggests this mixture to consist of 2/3 hematite and 1/3

As the previous data showed, the hematite products contained always some background level of sulfur in the range of 1-2% (see Figure 9). It was confirmed by dissolution of the hematite products in HCl followed by the addition of BaCl2 to precipitate BaSO4 that all the sulfur in the hematite precipitates was present as SO4. The incorporation of SO4 in the hematite precipitates was further characterized by FTIR spectroscopy. Figure 11 shows the FTIR spectrum of the hematite precipitate produced with 0.2 M initial Fe(III) which contained 1.3% S (or 3.92% SO4) and that of the precipitate produced

Figure 12. SEM pictures of the solid products obtained from hydrolysis of Fe2(SO4)3 either in the absence of zinc sulfate, (a) CFe(III),initial ) 0.2 M and (b) CFe(III),initial ) 0.5 M, or in the presence of zinc sulfate, (c) CFe(III),initial ) 0.2 M and (d) CFe(III),initial ) 0.7 M.

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Figure 13. SEM pictures of the hematite products obtained from hydrolysis of Fe2(SO4)3 either (a) in the absence of zinc sulfate (CFe(III),initial ) 0.2 M), or (b) in the presence of zinc sulfate (CFe(III),initial ) 0.2 M), compared to (c) basic ferric sulfate, and (d) hematite produced from hydrolysis of Fe(NO3)3.

in a nitrate medium following the procedure prescribed by Shang and Van Weert.33 Also included in Figure 11 is the spectrum of basic ferric sulfate. The appearance of the four infrared bands in the ν(SO4) region (S-O stretch) between 900 and 1300 cm-1 clearly distinguishes the hematite product obtained in this work from the SO4-free hematite material as well as basic ferric sulfate. Furthermore, the positions of the four infrared bands are at the same wavenumbers as when SO4 is chemisorbed on hematite (i.e. direct coordination of the sulfate anion to surface iron sites); the splitting of triply degenerate ν3 bands (C2v) indicates the formation of a bidentate surface complex.34-36 No evidence by XRD (Figures 6 and 7)salbeit this may be because BFS content is below the detection limit (4-5%)sor by FTIR analysis (Figure 11) was found to indicate this small amount of sulfate to be present as basic ferric sulfate. The adsorption of SO4 on hydrothermally produced hematite has been postulated as well by Dutrizac and Chen.12 The solid products were characterized as well in terms of morphology and specific surface area. Two product morphologies are evident: (i) aggregated particles consisting of small crystallites clustered together (Figure 12a,c) and (ii) acicular (needle-shaped) crystallites (Figures 12b,d). By referring to Figure 8, it is clear that the former morphology is characteristic of the material with low sulfate content, i.e., hematite, while the latter morphology is characteristic of the material with high sulfate content, i.e., basic ferric sulfate (BFS). The distinctively different morphologies of hematite and

Figure 14. Effect of the initial Fe(III) concentration on the specific surface area of precipitates produced from the hydrolysis of Fe2(SO4)3 in the presence/absence of ZnSO4.

basic ferric sulfate are better seen in Figure 13. Interestingly, the hematite produced in the absence of zinc sulfate was found to possess a lower specific surface area than that produced in the presence of zinc sulfate (see Figure 14). Moreover, the product containing BFS was found to possess very low specific surface (∼1 m2/g). In the presence of zinc sulfate, an increase in initial Fe-


(III) concentration caused the specific surface area to increase, as would be expected from crystallization theory, since homogeneous nucleation rate increases with supersaturation exponentially, leading to the production of ultrafine particles with a high degree of aggregation and impurity uptake.30,37,38 4. Summary and Conclusions Independent of the presence or not of zinc sulfate, the hydrolytic precipitation kinetics of ferric sulfate at 200 °C was determined to be first-order with respect to Fe(III), whereas the apparent kinetic constant was found to be 10-2 min-1. In the absence of zinc sulfate, the formation of metastable basic ferric sulfate (BFS) was found to be favored at high initial Fe(III) concentration (CFe(III),initial g 0.4 M). On the other hand, in the presence of zinc sulfate not only was the stability region of hematite enlarged as previously reported but also the formation of the metastable BFS was suppressed such that only hematite was produced at CFe(III),initial e 0.7 M. The solution and solid composition analyses concluded that hydrolysis of ferric sulfate at 200 °C proceeds initially via the formation of basic ferric sulfate, which converts gradually to hematite. The degree of completion of conversion depends on initial Fe(III) concentration, zinc sulfate content, and reaction time. The hydrothermally produced hematite materials were found to possess very high specific surface areas (50-80 m2/g), this being apparently due to the fact that precipitation proceeds via homogeneous nucleation. As a result of this high specific surface area, a certain amount of sulfate (3-5%) was found to contaminate hematite via surface complexation (chemisorption). After excluding the low sulfate and zinc content, the produced hematite was determined to have the following stoichiometric formula: Fe2O3-x(OH)2x‚yH2O, where x ) 0.08 and y ) 0.29. Acknowledgment This work was funded by the Natural Sciences and Engineering Research Council of Canada (NSERC) and the Akita Zinc Co. of Japan. Literature Cited (1) Baldwin, S. A.; Demopoulos, G. P.; Papangelakis, V. G. Mathematical Modelling of the Zinc Pressure Leach Process. Metall. Mater. Trans. B 1995, 26B, 1035. (2) Demopoulos, G. P.; Papangelakis, V. G. Recent Advances in Refractory Gold Processing. CIM Bull. 1989, Nov, 85. (3) Papangelakis, V. G.; Liu, H.; Rubisov, D. H. Solution Chemistry and Reactor Moldelling of the PAL Process: Successes and Challenges. In International Laterite Nickel Symposium 2004; Imrie, W. P., Lane, D. M., Eds.; TMS: Warrendale, PA, 2004; pp 289-305. (4) Cheng, T. C.; Demopoulos, G. P.; Shibachi, Y.; Masuda, H. The Precipitation Chemistry and Performance of the Akita Hematite ProcesssAn Integrated Laboratory and Industrial Scale Study. In Hydrometallurgy 2003; Young, C., et al., Eds.; TMS: Warrendale, PA, 2003; Vol. 2, pp 1657-1674. (5) Cheng, T. C.; Demopoulos, G. P. The Hematite Processs New Concepts for Increased Throughput and Clean Hematite Production. In Pressure Hydrometallurgy 2004; Papangelakis, V. G., Collins, M. J., Eds.; CIM: Montreal, Canada, 2004; in press. (6) Dutrizac, J. E. An Overview of Iron Precipitation in Hydrometallurgy. In Crystallization and Precipitation; Strathdee, G. I., Klein, M. O., Melis, L. A., Eds.; CIM: Montreal, Canada, 1987; pp 259-284.

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Received for review September 17, 2003 Revised manuscript received July 9, 2004 Accepted July 17, 2004 IE030711G

Hydrolysis of Ferric Sulfate in the Presence of Zinc Sulfate at 200Â° C

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