Purification of Leachate from Simultaneous Leaching of Galena

Article pubs.acs.org/IECR

Purification of Leachate from Simultaneous Leaching of Galena Concentrate and Pyrolusite and Preparation of PbSO4 and Mn3O4 Wenqing Qin,* Fen Jiao, Benjun Xu, and Hui Liu School of Minerals Processing and Bioengineering, Central South University, Changsha 410083, China ABSTRACT: This paper presents the results of investigation on simultaneous leaching of galena concentrate and pyrolusite in acid medium. It has been confirmed that the leached products as the raw materialsPbCl2, MnCl2 for the PbSO4 and Mn3O4 preparations could be obtained during the leaching processes. Special emphasizes are placed on the purification of the simultaneous leaching solution including removals of Fe2+, Pb2+, Zn2+, Ca2+, and preparations. The factors affecting the purification, such as pH, temperature, agitation rate, and dosages of Cl− and Na2S are investigated in leaching cells. The results show that, at appropriate conditions, iron is removed in the form of FeOOH, Pb2+ and Zn2+ are vulcanized by Na2S, and Ca2+ is removed in the form of CaF2 precipitation. After purification and crystallizing-out of PbCl2 (97.17%), the effects of pH, temperature, reaction time, concentration, stirring rate, and drying methods on preparations are examined. The mechanisms are addressed by theoretical calculations and experiments. The results indicate that the content of PbSO4 is 99.46% characterized by elemental analysis, XRD, and SEM. The purity of Mn3O4 is identified by TG-DSC, chemical analysis, and XRD, and the Mn in the Mn3O4 product accounts for 71.7%.

1. INTRODUCTION Maintenance-free (MF)1 lead-acid batteries are of present interest for electric vehicles (EV), especially as there is a growing public awareness of the environmental benefit of replacing the gasoline-powered vehicles with low-emission electric vehicles. L. T. Lam2 pointed out that the lead-acid battery has clear advantages, including a proud history of service, a well-established network of manufacturers and distributors, a competitive cost, and efficient recycling facilities, Robert A. Huggins3 also demonstrated the reasons for the widespread use of lead-acid batteries, such as the relatively low cost, ease of manufacture, and favorable electrochemical characteristics with rapid kinetics and good cycle life, so lead-acid batteries are still widely used in most of the commercially available EVs. The battery cathode is the control electrode of the lead/acid battery capacity, and the addition of positive active material can improve the specific energy. The traditional technology of electrode material processing is the oxidation of lead powder, but this has many disadvantages such as uneven quality of the plate, relative long production cycle, higher cost, and contamination and so on. Nowadays, to improve lead-acid battery, PbSO4 is used as the active material of the electrode instead of lead powder. Previous studies4−6 show that it is completely feasible to use PbSO4 as the electrode material. Mn3O4 is the main raw material of Mn−Zn ferrite, and spinel Mn3O4 is used to synthesize anode material for lithium batteries.7,8 Meanwhile, Mn3O4 powder is the catalyst in treating automobile exhaust (CO, CO2, NO2, etc.) and in selective catalytic reduction of nitrobenzene. In addition, Mn3O4 can be used as the pigment for some coating or paint. Because of the superior performance and wide application, Mn3O4 attracts more and more attention recently.9−14 Galena PbS is the principal source from which lead is commercially produced. Besides the pyrometallurgical methods where environmental pollution resulting from emission of SO2 © 2012 American Chemical Society

gas is a real problem, hydrometallurgical recovery of lead from galena may be a promising process with environmentally inert elemental sulfur being formed instead of sulfur dioxide. Hydrometallurgical leaching processes have the advantages of being environmentally friendly and having low energy consumption. Numerous leaching processes have been advocated for recovering lead values from galena concentrate. 15−20 FeCl3,21−23 Fe2(SO4)3,24,25 or H2O226,27 is used as oxidant in the leaching processes. The disadvantages of these technologies, they are expensive leaching reagents, they are difficult to recover and utilize, and they impose high requirements on the equipment, can be improved through the choice of an appropriate leaching reagent. The direct extraction technology of pyrolusite is to use a reducing agent to reduce pyrolusite in a liquid of lixiviant,28,29 and the leaching process is relatively complex with low leaching efficiency. Therefore, a cheap, efficient reducing agent will be a great benefit to the processing industry of manganese. Recently, some very interesting studies have been reported pertaining to the leaching of MnO2−sulfide mineral couples: simultaneous leaching of pyrite and pyrolusite,30 silver and manganese,31 manganese dioxide and gold,32 chalcopyrite,33−38 sphalerate and MnO2,39 and preconcentration of lead by MnO2.40 The application of the two minerals simultaneous leaching in the sulfide minerals provides some foundations for the simultaneous leaching of MnO2−galena process. Previous works have defined condition experiments of leaching technological factors and established mechanisms.41 The main advantages of galena−pyrolusite simultaneous leaching are moderate temperature, cheap leaching oxidant, simultaneous Received: Revised: Accepted: Published: 5596

October 27, 2011 March 22, 2012 March 26, 2012 March 26, 2012 dx.doi.org/10.1021/ie2024678 | Ind. Eng. Chem. Res. 2012, 51, 5596−5607

Industrial & Engineering Chemistry Research

Article

Table 1. Elemental Analysis (%) galena concentrate pyrolusite

Fe

SiO2

CaO

Al2O3

MgO

Zn

S

Pb

12.52 7.90

1.64 26.50

0.10 0.50

0.15 0.22

0.12 0.32

0.41

25.54

58.80

MnO2

Others

40.44

0.72 23.45

the advanced purification for manganese solution II, the solution is called manganese solution II. 2.2. Purification Experiment. The galena concentrate was leached according to the optimal conditions,41 and in the later period of leaching, the solid−liquid separation was carried out by adjusting the pH of leachate. The filtrate was diluted and cooled to ambient temperature (25 °C), and then filtered to obtain the PbCl2 solid. The separation efficiency was represented by precipitation rate of lead (η%) which was the mass ratio of Pb contained in PbCl2 solid to the original filtrate. The removal of impurities (Pb2+, Zn2+, Ca2+) in Mn2+ solution was investigated. Na2S was added into the Mn2+ solution, and the reaction was conducted under the ambient temperature (25 °C) with the stirring speed of 200 rpm for 30 min and filtered after 24 h. 2.3. PbSO4 Preparation and Mn3O4 Preparation. After the purification, PbCl2 solid was mixed with deionized water, and then sulfuric acid was added into the suspension. The reaction conducted in the constant water bath. As soon as the reaction completed, the solution was filtered and the filter cake was dried in the blast oven. Finally, the samples were analyzed. The manganese solution was added into the three-neck flask. Ammonia solution was then added into the manganese solution. The reaction temperature was maintained constantly in the water trough, the solution was filtered, and the filter cake was dried as soon as the reaction completed. Finally, the samples were analyzed. 2.4. Analytical Methods. The obtained samples were characterized on a Rigaku Dmax-2000 X-ray powder diffractometer (XRD) with Cu Kα radiation (λ = 1.5418 Å). The operation voltage and current were kept at 40 kV and 40 mA, respectively. Ions’ concentrations in solution were measured with an atomic absorption spectrophotometer (Z-8000 PE). The size and morphology of the products were determined at 20 kV by a XL30 S-FEG scanning electron microscope (SEM).

leaching of galena and pyrolusite, and avoidance of volatile plumbous compounds. In this work, a method to prepare electrode material PbSO4 and Mn3O4 from minerals directly by simultaneous leaching of galena concentrate and MnO2 is proposed. The objectives of this work are to test the optimum conditions required for the purification of leachate from simultaneous leaching of galena concentrate and pyrolusite, and to test the factors affecting PbSO4 preparation and Mn3O4 preparation. The phase, morphology, and purity of the products are identified. The mechanisms of the purification and preparation are also discussed.

2. EXPERIMENTAL SECTION 2.1. Materials. The galena concentrate was acquired from Fan Kou Mill in Guang Dong province of China. The pyrolusite powder, with a percentage of 96% for −0.14 mm, was industrial grade. The chemical composition is shown in Table 1. The flow sheet of the direct preparation of PbSO4 and Mn3O4 from simultaneous leaching of galena concentrate and pyrolusite is shown in Figure 1.

3. RESULTS AND DISCUSSION 3.1. Purification of Leachate. The condition experiments of simultaneous leaching technological factors and mechanisms were based on previous work.41 The multifactor experiments are investigated. The results are shown in Table 2. As can be seen from Table 2, when the NaCl dosage is 700 g, the average value is the biggest one, hence the optimal test condition is 700 g. Likewise, the optimal test conditions of the other factors are shown in Table 2. The optimum conditions are as follows: PbS concentrate/pyrolusite material/NaCl = 1:1:3.5 (wt %); temperature, 70 °C; reaction time, 90 min; HCl, 500 mL; and HCl addition twice (V1/V2 = 4:1). Recoveries of 95% for Pb and 96% for Mn were achieved under these conditions. Table 3 shows the impurities in the leachate, and the lead exists in the form of coordination forms. 3.1.1. Removal of Iron in the Leachate. The iron contained in galena and pyrolusite is leached into the leachate. Most of iron exists in the form of Fe2+ in this study, because there is always superfluous pyrolusite in the leaching process, by adjusting the pH value using NaOH in the later period of

Figure 1. The flow sheet of the direct preparations of PbSO4 and Mn3O4 from simultaneous leaching of galena concentrate and pyrolusite.

Remarks: After the removal of iron and PbCl2 precipitation from leachate, the solution is called manganese solution I; after 5597

dx.doi.org/10.1021/ie2024678 | Ind. Eng. Chem. Res. 2012, 51, 5596−5607


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Table 2. Leaching Experiment Results test no.

a

NaCl (g)

pyrolusite (g)

temperature (°C)

time (min)

HCla (mL)

adding times of HCl

Pb recovery (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

200 200 200 400 400 400 700 700 700 200 200 200 400 400 400 700 700 700

200 300 400 200 300 400 200 300 400 200 300 400 200 300 400 200 300 400

50 60 70 60 70 50 50 60 70 70 50 60 70 50 60 60 70 50

20 50 90 50 90 20 90 20 50 50 90 20 20 50 90 90 20 50

200 300 500 200 300 500 300 500 200 500 200 300 300 500 200 500 200 300

1 2 3 3 1 2 3 1 2 1 2 3 2 3 1 2 3 1

7.95 27.37 35.00 33.29 30.90 26.25 66.67 56.01 59.83 33.60 29.36 17.60 33.35 30.73 38.80 85.75 70.94 42.50

mean 1 mean 2 mean 3 optimal conditions

25.15 33.22 63.62 700

43.44 40.89 36.66 200

33.91 43.14 43.94 70

35.35 37.89 47.74 90

40.03 36.40 44.56 500

34.96 43.65 42.37 2

95

Remarks: The concentration of HCl is 1.5 mol/L.

Table 3. Chemical Analysis (g/L) leachate manganese solution I manganese solution II

Mn2+

Pb2+

Zn2+

Ca2+

Fe(mg/L)

Na+

21.6 20.1 18.24

83.67 1.61 3.7 × 10−5

1.43 0.85 3.1 × 10−5

0.58 0.47 2.5 × 10−5

14.8 3 × 10−3 12 × 10−5

37.8 36.0 35.9

Figure 3. Relationships between log[Men+] and pH in the hydrolysis: (dashed line) 25 °C, (solid line) 70 °C. (1) Fe3+; (2) Zn2+; (3) Fe2+; (4) Pb2+; (5) Mn2+.

Figure 2. Effects of pH values on the iron removal ratio.

leaching, Fe2++ is oxidized to goethite precipitate and then removed. The reaction is shown in eq 1: 2Fe2 + + MnO2 + 2OH− → 2FeOOH↓ + Mn 2 +

It can be seen from Figure 2 that the iron concentration decreases from 1.19 × 10−3 mol/L to 9.29 × 10−5 mol/L when the pH value changes from 1.9 to 3.8. There is no significant decrease in the iron concentration when the pH varies from 3.8 to 5.4. And higher pH is not beneficial to the recoveries of Pb2+ and Mn2+, because Pb2+ and Mn2+ will precipitate in the forms

(1)

The relationship between the pH and iron concentration is shown in Figure 2. 5598



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Figure 4. XRD of the residues.

1 log αFe2 + 2

pH25 = 6.65 − pH70 = 5.6 −

1 log αFe2 + 2

Fe3 + + 2H2O ↔ FeOOH + 3H+

pH25 = 0.3154 − pH70 = 0.807 −

1 log αFe3 + 3 1 log αFe3 + 3

Zn 2 + + 2H2O ↔ Zn(OH)2 + 2H+

Figure 5. Effects of dilution on precipitation rate of PbCl2 crystal. Leachate, 1000 mL; T, 25 °C.

Men + + nH2O → Me(OH)n + nH+ 1 pH = pH − log αMe n ϕ

pH25 = 5.85 −

1 log αZn 2 + 2

pH70 = 4.91 −

1 log αZn 2 + 2

Pb2 + + 2H2O ↔ Pb(OH)2 + 2H+

of Pb(OH)2 and Mn(OH)2, respectively, misreporting in residue at higher pH. 3.1.2. Mechanism of Hydrolytic Purification. The pH is the key factor in the metal hydrolyzation. The hydrolysis reaction of metal ions in the definite pH is shown in eqs 2 and 3.

pH25 = 6.82 −

1 log αPb2 + 2

pH70 = 6.18 −

1 log αPb2 + 2

(2)

Mn 2 + + 2H2O ↔ Mn(OH)2 + 2H+

⎛ ϕ ⎞ ⎜⎜where pH ϕ = − ΔG ⎟⎟ 1364n ⎠ ⎝

pH25 = 7.655 −

(3) Φ

pH70 = 6.55 −

pH is the standard pH, when all of the activity coefficients for reactants and resultants equal to 1 under the standard state. pHΦ is the important sign of hydrolysis degree of metallic ions. When pH > pHΦ, the metallic ions will hydrolyze and the metal hydroxides will precipitate. When pH < pHΦ, the activity coefficients for reactants and resultants are greater than 1; the hydroxides will dissolve. Equation 3 represents the equilibrium conditions; “n” is the valence state. The relation formulas between pH and activity coefficients are shown in eq 4 to eq 8 when temperature is 25 and 70 °C, respectively. Fe2 + + 2H2O ↔ Fe(OH)2 + 2H+

(5)

(6)

(7)

(8)

1 log αMn 2 + 2 1 log αMn 2 + 2

The relationship between log[Men+] and pH is shown in Figure 3. It can be seen that, at 70 °C, the lowest hydrolysis pH of Fe2+, Fe3+, Zn2+, Pb2+, and Mn2+ is 5.6, 0.807, 4.91, 6.18, and 6.55, respectively. The initial hydrolysis pH of Fe2+ is higher than that of Fe3+, meanwhile, the hydrolysis pH decreases with temperature increases. The XRD of hydrolysis reaction product in the residue at different pH values is shown in Figure 4. It can be seen that iron is oxidized into FeOOH precipitation by superfluous pyrolusite, and then removed.

(4) 5599



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Table 4. Chemical Analysis (%) PbCl2 PbSO4 Mn3O4

Pb

Fe

Ca

Zn

Si

72.36 68.2 0.0047

0.0004

0.0002

0.0007

0.00005

0.0014

0.015

SO42−

Cl−

31.51

0.23

Mn

K

Na

Mg

71.7

0.00036

0.00035

0.0018

Figure 6. XRD of the PbCl2. Diluted filtrate, 1500 mL; T, 25 °C, NaCl, 10 g; H2SO4, 3 mL; pH, 5.

Figure 7. SEM of PbCl2 and PbSO4. (A) Diluted filtrate, 1500 mL; T, 25 °C; (B) PbCl2, 10 g; H2SO4, 3 mL; T, 37 °C; reaction time, 60 min; pH, 3.

Figure 8. Effects of Na2S dosage on precipitations of Pb2+ and Zn2+.

will increase with higher temperature and Cl− concentration. Especially, the effect of temperature on PbCl2 solubility is obvious. Therefore, cooling and diluting methods are employed to precipitate PbCl2, and to achieve Pb2+ separation from Mn2+. Meanwhile, other ions (Ca2+, Zn2+, etc.) cannot incorporate

3.1.3. Crystallization of PbCl2. After iron removal, PbCl2 crystal can be obtained by diluting the leachate and cooling to 25 °C. Pb2+ and Cl− exist in the different coordination forms with changes in temperature,41 and the possible reactions are from eq 9 to eq 12. The solubility of PbCl2 in NaCl solution 5600



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Figure 9. Effects of temperature and initial pH on concentrations of Pb2+ and Zn2+. Manganese solution I, 100 mL; stirring speed, 200 rpm; reaction time, 30 min; [Na2S] = 0.075 mol/L.

decreases with increasing pH, and accordingly, the effective concentration and effective dosage of S2− are improved. The fact that the Pb2+ concentration is lower than that of Zn2+ is due to the different solubility products of PbS and ZnS (Ksp(PbS) < Ksp(ZnS)). (3) Removal of Ca2+. Ca2+ in the leachate can be removed in the form of CaF2 by NH4F as precipitator. The effect of NH4F concentration on Ca2+ concentration is investigated. It can be seen from Figure 10 that there is no significant reduction in

into the PbCl2 crystallization, so it is beneficial to PbCl2 purification. Pb2 + + Cl− ↔ PbCl+

(9)

PbCl+ + Cl− ↔ PbCl2

(10)

PbCl2 + Cl− ↔ PbCl− 3

(11)

2− − PbCl− 3 + Cl ↔ PbCl 4

(12)

The effect of concentrate of diluted filtrate on precipitation rate of PbCl2 crystal is shown in Figure 5. The chemical analysis of PbCl2 solid is shown in Table 4. As Figure 6 demonstrated, the characteristic peak of PbCl2 is obvious. As can be seen from Figure 6 and Table 4, the content of PbCl2 is 97.17%. The SEM (Figure 7) shows that PbCl2 is an acicular crystal with a length of 100−200 μm. 3.1.4. Removal of Impurities from Mn2+ Solution. After the removal of iron and after PbCl2 precipitation, the chemical analysis of manganese solution I is shown in Table 3. Pb2+, Zn2+, and Fe2+ can be removed by adding sodium sulfide. In this section, effects of Na2S dosage, temperature, and pH on the purification are conducted. Ca2+ is removed in the form of CaF2 precipitation. (1) Effects of Na2S dosage on the Pb2+, Zn2+ precipitations. The effects of Na2S dosage on precipitations of Pb2+, Zn2+ in the solution are shown in Figure 8. The concentrations of Pb2+, Zn2+ decrease when the Na2S dosage increases, and Zn2+ precipitation is seriously affected by Na2S: when the Na2S dosage changes from 0.025 mol/L to 0.125 mol/L, the Zn2+ concentration decreases from 9.00 × 10−4 mol/L to 1.28× 10−6 mol/L, while Pb2+ concentration decreases from 5.00 × 10−5 mol/L to 7.73 × 10−7 mol/L, which is due to the differences between Pb2+ and Zn2+ concentration, and different solubility products of ZnS and PbS. (2) Effects of temperature and initial pH on the Pb2+, Zn2+ precipitations. The effects of temperature on concentrations of Pb2+ and Zn2+ are indicated in Figure 9. It can be seen that the concentrations of Pb2+, Zn2+ increase with the temperature increasing, so it is not beneficial for the removal of Pb2+, Zn2+ ions when the temperature increases. So in order to reduce the Na2S dosage and remove the Pb2+ and Zn2+, the temperature should not be too high. The concentrations of Pb2+ and Zn2+ decline when the pH increases, and the concentration of Pb2+ is lower than that of Zn2+. That is because the H+ concentration in solution

Figure 10. The effect of NH4F concentration on Ca2+ concentration. Manganese solution I, 100 mL; stirring speed, 200 rpm; T, 25 °C.

Ca2+ concentration when the NH4F concentration is more than 0.054 mol/L. 3.1.5. Theory of Sulfidation Process. The heavy metal ions are sulfurized by sulfur and then removed. The ionizing dissolution of the sulfide which is formed by heavy metal ion and S2− is Me2Sn ↔ 2Men + + nS2 −

(13)

K sp = [Men +][S2 −]n 5601



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Besides, the solubility product constant of PbSO4 (Ksp = 1.7 × 10−8) is much smaller than that of PbCl 2 (Ksp = 1.7 × 10 −5), therefore, the solubility product theory also indicates the reaction can be conducted. In this section, factors affecting the content and average particle size of PbSO4 are investigated, such as reaction time, temperature, initial pH, and dosage of H2SO4, and the results are shown in Figure 12. As indicated in Figure 12-1, the formation speed of PbSO4 is fast comparatively at the beginning stage and becomes constant after 45 min (99.62%). There is a peak where average particle size is 445 nm at 35 min. The crystallization process is divided into three steps: nuclei formation, nuclei growth, grain dissolution, and adsorption equilibrium. So the reaction time has a great influence on particle size and crystallization processes. When the reaction time is short, the content is low and the particle size is small, which occur because the reaction between PbCl2 and H2SO4 is not complete, and the PbSO4 nuclei cannot grow. As shown in Figure 12-2, higher content with a small average particle size of PbSO4 can be obtained at high temperature. In Figure 12-3, when pH increases, it can be seen that the content of PbSO4 decreases and the average particle size becomes large, which can be attributed to the great effects of pH on the surface energy, absorbability of grains, and formation of Pb(OH)2 precipitation. Results show (Figure 12-4) that the PbSO4 content and average particle size change to be steady when the dosage of H2SO4 is more than 4 mL. The PbSO4 was obtained under the experimental conditions: PbCl2, 10 g; H2SO4, 3 mL; temperature, 37 °C; reaction time, 60 min; and pH value, 3. The purity of PbSO4 is characterized by elemental analysis (Table 4) and XRD (Figure 13), and the morphology of the as-synthesized PbSO4 is detected by SEM. The results show that most of the PbCl2 has been transformed into PbSO4, and the content of PbSO4 is 99.46%. As the SEM (Figure 7) shows, the product has been transformed from acicular crystal into granular crystal. 3.2.2. Mn3O4 Preparation. The raw material for Mn3O4 preparation is manganese chloride, and the composition analysis is shown in Table 3. Mn3O4 is prepared in aqueous ammonia, and Eh−pH of Mn−O−H2O at 60 °C is shown in Figure 14. The main reactions are as follows:

S2− concentration is determined by the pH because the solution is weakly acidic in the test and can be calculated by H2S(g) ↔ 2H+ + S2 −

K = K1K2 =

(14)

[H+][HS−] [H+][S2 −] [H+]2 [S2 −] × = pH S [HS−] pH S 2

2

1 1 1 − log[H+] = log K sp − log(K × pH S ) − log[Men +] 2 2n 2 n (15)

At 25 °C, n = 2 as for divalent metal; pH2S = 1.01 × 105 Pa, KH2S = K1K2 = 10−8 × 10−12.9 = 10−20.9, −log Ksp(FeS) = 18.1; −log Ksp(ZnS) = 24.7; −log Ksp(PbS) = 27.5; −log Ksp(MnS) = 10.5. The relationships between pH and metal ion concentration can be obtained by putting the data into eq 15 and are shown in eq 16 to eq 19. pH = −

1 log[Fe2 +] + 5.92 2

(16)

pH = −

1 log[Zn 2 +] + 4.275 2

(17)

pH = −

1 log[Pb2 +] + 3.575 2

(18)

pH = −

1 log[Mn 2 +] + 7.825 2

(19)

It can be seen that pH is proportional to the negative logarithm of metal ion concentration and the relationship between log[Men+] and pH is shown in Figure 11.

Mn 2 + + NH3· H2O → Mn(OH)2 ↓ + NH+ 4

(21)

6Mn(OH)2 + O2 → 2Mn3O4 ↓ + 6H2O

(22)

2+

The oxidations of Mn into Mn3O4 indicated in Figure 14 are as follows: Mn3O4 + 2H+ + 2e− ⇔ 3MnO + H2O

(23)

φ25 = 0.2232 − 0.05916pH Figure 11. Relationships between log[Me ] and pH (25 °C). n+

φ50 = 0.2185 − 0.06412pH

It can be concluded that precedence order of sulfide precipitated at the same pH is

2Mn3O4 + 2H2O ⇔ 3Mn2O3 + 2H+ + 2e−

Pb2 + > Zn 2 + > Fe2 + > Mn 2 +

φ25 = 0.8318 − 0.05916pH

3.2. Preparations of PbSO4 and Mn3O4. 3.2.1. PbSO4 Preparation. The reaction between PbCl2 and H2SO4 is PbCl2 + H2SO4 → PbSO4 + 2HCl

(24)

φ50 = 0.8215 − 0.06412pH

The compounds of Mn2+ oxidation are MnO2, Mn2O3, MnO, and Mn3O4 at alkaline pH; Mn3O4 is relatively stable in the black area. The potation of Mn3O4 is narrow (0.3 V), so pH and potential are the key factors affecting the purity of Mn3O4.

(20)

The ΔrGmΘ of this reaction is −17.68 KJ/mol < 0, which indicates this reaction can be carried out in the standard state. 5602



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Figure 12. Effects of factors on the content and average particle size of PbSO4. (1) PbCl2, 10 g; H2SO4, 3 mL; temperature, 25 °C; pH, 3. (2) PbCl2, 10 g; H2SO4, 3 mL; reaction time, 60 min; pH, 3. (3) PbCl2, 10 g; H2SO4, 3 mL; reaction time, 60 min; temperature, 25 °C. (4) PbCl2, 10 g; reaction time, 60 min; pH, 3; temperature, 25 °C.

Figure 13. XRD of PbSO4: PbCl2, 10 g; H2SO4, 3 mL; T, 37 °C; reaction time, 60 min; pH, 3.

The multifactor experiments including pH, reaction time, temperature, initial concentration of Mn2+ are investigated by orthogonal test L4×3. The results are shown in Table 5. The mean analysis in Table 6 indicates the best experimental conditions, and pH, reaction temperature, initial concentration of Mn2+, and reaction time are, from high to low, influencing the recovery of Mn3O4. The product Mn3O4 is prepared under selected experimental conditions: pure Mn2+ solution of 200 mL, temperature at 50 °C, pH of 9, dropping rate of ammonia of 6 mL/min, stirring rate of 400 ± 25 rpm, reaction time of 2 h, drying in vacuum for 9 h with a temperature of 90 °C.

Aqueous ammonia is a pH regulator in this test, the pH of solution is 9 ± 0.2, at the same time, the product of NH4Cl can stabilize the potential in the preparation process. The addition of alkali resulted in a quick precipitation, because the reaction rate of manganese precipitation is fast. At alkaline pH, Mn3O4 easily precipitates to MnOOH, with the dropping rate of ammonia, reaction time, and drying method having great effects on the content of Mn3O4. Figure 15 and Figure 16 show the XRD of the oxidation products in MnCl2 solution at different reaction times and different drying methods, respectively. It is obvious that the characteristic peaks of MnOOH appear when the reaction time is 4 h and when the oxidation product dries in aeration for 9 h. 5603



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Table 5. Results of Multifactor Experiments for Mn3O4 Preparation

Figure 17 shows the TG (thermo-gravimetry)-DSC (differential scanning calorimetry) of Mn3O4. There is an exothermic peak at 648 °C and an endothermic peak at 989 °C in the DSC curve. In the TG curve, meanwhile, a mass-gain peak is observed at 648 °C (2.47%), and at 989 °C (−3.33%) a massloss peak appears. This is in agreement with the TG-DSC

test no.

pH

time (h)

T (°C)

initial concentration (mol/L)

recovery (%)

1 2 3 4 5 6 7 8 9

7.5 7.5 7.5 8.2 8.2 8.2 9 9 9

1 2 3 1 2 3 1 2 3

30 60 50 60 50 30 50 30 60

1.45 0.725 0.363 0.363 1.45 0.725 0.725 0.363 1.45

34.616 63.57 48.07 87.11 73.64 71.83 91.77 87.19 89.39

Table 6. Means and Ranges of Mn3O4 Recovery and Optimal Conditions

Figure 14. Eh−pH of Mn−O−H2O (60 °C).

results

pH

mean 1 mean 2 mean 3 range optimal conditions

48.75 77.527 89.450 40.7 9

time (h) T (°C) 71.163 74.8 69.763 5.037 2

64.543 80.023 71.16 15.48 50

initial concentration (mol/L) 65.88 75.723 74.123 9.843 0.725

Figure 15. XRD of the oxidation products with different reaction times: Mn2+ solution II, 200 mL; T, 50 °C; pH, 9; dropping rate of ammonia, 6 mL/min; stirring rate, 400 ± 25 rpm; drying in vacuum for 9 h with the temperature of 90 °C.

Figure 16. XRD of the oxidation products with different drying methods: Mn2+ solution II, 200 mL; T, 50 °C; pH, 9; dropping rate of ammonia, 6 mL/min; reaction time, 2 h; stirring rate, 400 ± 25 rpm. 5604



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Figure 17. TG (thermo-gravimetry)−DSC (differential scanning calorimetry) of Mn3O4.

Figure 18. XRD of the Mn3O4.

curves of Mn3O4(R) in the open air reported by V. Berbenni.42 The reactions at 648 and 989 °C are presented as follows: 4Mn3O4 + O2 → 6Mn2O3

(25)

6Mn2O3 → 4Mn3O4 + O2

(26)

sulfide, and Ca2+ was removed in the form of CaF2 precipitation. The relationship between the negative logarithm of metal ion concentration and pH indicates that at the same pH, the precedence order of sulfide precipitated is as follows: Pb2+ > Zn2+ > Fe2+ > Mn2+. After the precipitation of PbCl2, PbSO4 was prepared by mixing H2SO4 with PbCl2. The results showed that when the reaction time was short, the content of PbSO4 was low and particle size was small, which was due to the inadequate reaction between PbCl2 and H2SO4, and the nuclei PbSO4 could not grow. PbSO4 with a higher content and small average particle sizes could be obtained at high temperature. pH had great effects on the surface energy and absorbability of the grains. Mn3O4 was prepared in aqueous ammonia, and pH and potential were key factors affecting the purity of Mn3O4. Aqueous ammonia was a pH regulator in this test; the pH of solution was 9 ± 0.2, and at the same time, the product of NH4Cl could stabilize the potential in the preparation process. The results show that longer reaction time and drying methods influence the purity of Mn3O4, because the reaction rate of manganese precipitation was fast. MnOOH appeared when the reaction time was 4 h and when the oxidation product drying in

The theoretical mass-gain and mass-loss of eqs 25 and 26 are 3.49% and 3.33%, respectively, which equal the TG results. The purity of Mn3O4 is identified by the chemical analysis (Table 4) and XRD (Figure 18).

4. CONCLUSIONS The raw materials for PbSO4 and Mn3O4 preparations were obtained directly from simultaneous leaching of galena concentrate and pyrolusite. In the purification of the leachate, the iron was removed in the form of FeOOH, the pH was the key factor in the metal hydrolyzation, and the appropriate pH was 4.0. To separate Pb2+ from Mn2+, cooling and diluting methods were employed in the precipitation of PbCl2, because the solubility of PbCl2 in NaCl solution increased with temperature and Cl− dosage. After the removal of iron and PbCl2 crystal, Pb2+, Zn2+, and Fe2+ in the manganese solution were removed by sodium 5605



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(16) Silva, G. D. Kinetics and Mechanism of the Bacterial and Ferric Sulphate Oxidation of Galena. Hydrometallurgy 2004, 75, 99−110. (17) Jiang, L.; Zhou, H. Y.; Peng, X. T.; Ding, Z. H. Bio-oxidation of Galena Particles by Acidithiobacillus Ferrooxidans. Particuology 2008, 6, 99−105. (18) Bang, S. S.; Deshpande, S. S.; Han, K. N. The Oxidation of Galena Using Thiobacillus Ferrooxidans. Hydrometallurgy 1995, 37, 181−192. (19) Makita, M.; Esperon, M.; Pereyra, B.; Lopez, A.; Orrantia, E. Reduction of Arsenic Content in a Complex Galena Concentrate by Acidithiobacillus Ferrooxidans. BMC Biotechnol. 2004, 4, 22. (20) Pacholewska, M.; Cwalina, B. The Role of Bacteria in the Metals Leaching from Lead-Zinc Flotation Tailings. J. Biotechnol. 2007, 131, S264−S264. (21) Morin, D.; Gaunand, A.; Renon, H. Representation of the Kinetics of Leaching of Galena by Ferric Chloride in Concentrated Sodium Chloride Solutions by a Modified Mixed Kinetics Model. Metall. Mater. Trans. B. 1985, 16, 31−39. (22) Wang, S.; Fang, Z.; Wang, Y.; Chen, Y. Electrogenerative Leaching of Galena with Ferric Chloride. Miner. Eng. 2003, 16, 869− 872. (23) Wang, S. F.; Fang, Z. Mechanism of Influence of Ferric Ion on Electrogenerative Leaching of Sulfide Minerals with FeCl3. Trans. Nonferrous Met. Soc. 2006, 16, 473−476. (24) Gabriel, D. S. Kinetics and Mechanism of the Bacterial and Ferric Sulphate Oxidation of Galena. Hydrometallurgy 2004, 75, 99− 110. (25) Dutrizac, J.; Chen, T. The Leaching of Galena in Ferric Sulfate Media. Metall. Mater. Trans. B 1995, 26, 219−227. (26) Aydogan, S.; Aras, A.; Ucar, G.; Erdemoglu, M. Dissolution Kinetics of Galena in Acetic Acid Solutions with Hydrogen Peroxide. Hydrometallurgy 2007, 89, 189−195. (27) Aydogan, S.; Erdemogiu, M.; Ucar, G.; Aras, A. Kinetics of Galena Dissolution in Nitric Acid Solutions with Hydrogen Peroxide. Hydrometallurgy 2007, 88, 52−57. (28) Kholmogorov, A. G.; Zhyzhaev, A. M.; Kononov, U. S.; Moiseeva, G. A.; Pashkov, G. L. The Production of Manganese Dioxide from Manganese Ores of Some Deposits of the Siberian Region of Russia. Hydrometallurgy 2000, 56, 1−11. (29) Nayak, B. B.; Mishra, K. G.; Paramguru, R. K. Kinetics and Mechanism of MnO2 Dissolution in H2SO4 in the Presence of Pyrite. J. Appl. Electrochem. 1999, 29, 191−200. (30) Paramguru, R. K.; Mishra, K. G.; Kanungo, S. B. Electrochemical Phenomena in MnO2−FeS2 Leaching in Dilute HCl. Part 2: Studies on Polarization Measurements. Can. Metall. Q. 1998, 37, 395−403. (31) Jiang, T.; Yang, Y. B.; Huang, Z. C.; Qiu, G. Z. Simultaneous Leaching of Manganese and Silver from Manganese−Silver Ores at Room Temperature. Hydrometallurgy 2003, 69, 177−186. (32) Liu, G. W.; Xia, X.; Pan, B.; Zeng, Z. G.; Xiao, S. W. Study on Extracting Gold and Manganese by Combination Leaching Gold Concentrates and Pyrolusite. Min. Metall. Eng. 2010, 30, 81−84. (33) Madhuchhanda, M.; Devi, N. B.; Rath, P. C.; Rao, K. S.; Paramguru, R. K. Dissolution of Base-Metal Sulfide Minerals in Hydrochloric Acid in the Presence of Manganese Dioxide. Miner. Metall. Process 2001, 18, 106−109. (34) Feng, D.; Van Deventer, J. S. J. Interactions between Sulphides and Manganese Dioxide in Thiosulphate Leaching of Gold Ores. Miner. Eng. 2007, 20, 533−540. (35) Devi, N. B.; Madhuchhanda, M.; Rao, K. S.; Rath, P. C.; Paramguru, R. K. Oxidation of Chalcopyrite in the Presence of Manganese Dioxide in Hydrochloric Acid Medium. Hydrometallurgy 2000, 57, 57−76. (36) Devi, N. B.; Madhuchhanda, M.; Rath, P. C.; Rao, K. S.; Paramguru, R. K. Simultaneous Leaching of a Deep-Sea Manganese Nodule and Chalcopyrite in Hydrochloric Acid. Metall. Mater. Trans. B 2001, 32, 777−784. (37) Gantayat, B. P.; Rath, P. C.; Paramguru, R. K.; Rao, S. B. Galvanic Interaction between Chalcopyrite and Manganese Dioxide in Sulfuric Acid Medium. Metall. Mater. Trans. B 2000, 31, 55−61.

aeration for 9 h. Shortening the reaction time and drying in vacuum could hamper the formation of MnOOH. The purity of PbSO4 was characterized by elemental analysis and XRD, and morphology of the PbSO4 was detected by SEM; the content of PbSO4 was 99.46%. The purity of Mn3O4 was identified by TG-DSC, chemical analysis, and XRD, and the Mn in the Mn3O4 product accounted for 71.7%.

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AUTHOR INFORMATION

Corresponding Author

*Tel.: +86 731 8830884. Fax: +86 731 8710804. E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors gratefully acknowledge the financial support of this research by National Basic Research Program of China (Grant 2010CB630905).

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REFERENCES

(1) Vinod, M. P.; Vijayamohanan, K. Effect of Gelling on the Impedance Parameters of Pb/PbSO4 Electrode in Maintenance-Free Lead-Acid Batteries. J. Power Sources 2000, 89, 88−92. (2) Lam, L. T.; Newnham, R. H.; Ozgun, H.; Fleming, F. A. Advanced Design of Valve-Regulated Lead-Acid Battery for Hybrid Electric Vehicles. J. Power Sources 2000, 88, 92−97. (3) Huggins, R. A. Lead-Acid Batteries. Energy Storage 2010, 237− 259. (4) Toussaint, G.; Torcheux, L.; Alzieu, J.; Camps, J. C.; Livigni, D.; Sarrau, J. F.; Vaurijoux, J. P.; Benchetrite, D.; Gauthier, V.; Vilasi, M. Effect of Additives in Compressed Lead-Acid Batteries. J. Power Sources 2005, 144, 546−551. (5) Torcheux, L.; Rouvet, C.; Vaurijoux, J. P. Effect of a Special Additive on the Performance of Standby Valve-Regulated Lead Acid Batteries. J. Power Sources 1999, 78, 147−155. (6) Nelson, R. The Basic Chemistry of Gas Recombination in LeadAcid Batteries. JOM 2001, 53, 28−33. (7) Wang, H. L.; Cui, L. F.; Yang, Y.; Casalongue, H. S.; Robinson, J. T.; Liang, Y. Y; Dai, H. J. Mn3O4−Graphene Hybrid as a High-Capacity Anode Material for Lithium Ion Batteries. J. Am. Chem. Soc. 2010, 132, 13978−13980. (8) Gao, J.; Lowe, M. A.; Abruna, H. D. Spongelike Nanosized Mn3O4 as a High-Capacity Anode Material for Rechargeable Lithium Batteries. Chem. Mater. 2011, 23, 3223−3227. (9) Lee, K. J.; Iguchi, A.; Iguchi, E. Polaronic Conduction of Electrons in [Alpha]-Mn3O4 Slightly Doped with Li. J. Phys. Chem. Solids 1993, 54, 975−981. (10) Yang, Y. N.; Huang, R. L.; Chen, L.; Zhang, J. Y. Redox Behavior of Trimanganese Tetroxide Catalysts. Appl. Catal., A 1993, 101, 233−252. (11) Izumiya, K.; Akiyama, E.; Habazaki, H.; Kawashima, A.; Asami, K.; Hashimoto, K.; Kumagai, N. Surface Activation of Manganese Oxide Electrode for Oxygen Evolution from Seawater. J. Appl. Electrochem. 1997, 27, 1362−1368. (12) Synthesis of Spherical Assembly Composed of Mn 3 O4 Nanoparticles and Its Adsorption Behavior for Alizarin Red in Water. Mater. Sci. Forum 2011, 694, 165-169. (13) Wang, W. M.; Yang, Y. N.; Zhang, J. Y. Selective Reduction of Nitrobenzene to Nitrosobenzene over Different Kinds of Trimanganese Tetroxide Catalysts. Appl. Catal., A 1995, 133, 81−93. (14) Li, P.; Nan, C. Y.; Wei, Z.; Lu, J.; Peng, Q.; Li, Y. D. Mn3O4 Nanocrystals: Facile Synthesis, Controlled Assembly, And Application. Chem. Mater. 2010, 22, 4232−4236. (15) Dutrizac, J. E.; Chen, T. T. The Leaching of Galena in Ferric Sulfate Media. Metall. Mater. Trans. B 1995, 26, 219−227. 5606



Article

(38) Wang, S. F.; Fang, Z.; Tan, Y. F. Simultaneous Electrogenerative Leaching of Chalcopyrite Concentrate and MnO2. J. Cent. South Univ. (Sci. Technol.) 2006, 13, 49−52. (39) Wang, S. F.; Fang, Z.; Tan, Y. F. Application of Thermoelectrochemistry to Simultaneous Leaching of Sphalerite and MnO2. J. Therm. Anal. Calorim. 2006, 85, 741−743. (40) Burnett, J. L.; Croudace, I. W.; Warwick, P. E. Lead Preconcentration Using a Novel Manganese Dioxide Resin. Environ. Earth Sci. 2011. (41) Long, H. Z.; Chai, L. Y.; Qin, W. Q. Galena-Pyrolusite Coextraction in Sodium Chloride Solution and Its Electrochemical Analysis. Trans. Nonferrous Metal. Soc. 2010, 20, 897−902. (42) Berbenni, V.; Marini, A. Oxidation Behaviour of Mechanically Activated Mn3O4 by TGA/DSC/XRPD. Mater. Res. Bull. 2003, 38, 1859−1866.

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Purification of Leachate from Simultaneous Leaching of Galena

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