Al2O3 Catalyst with Mineral Acids

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Ind. Eng. Chem. Res. 2010, 49, 6514–6520

Dissolution of Cobalt from CoO/Al2O3 Catalyst with Mineral Acids Ikram Boukerche, Naima Habbache, Nadia Alane, Souad Djerad,* and Lakhdar Tifouti Laboratory of EnVironmental Engineering, Department of Chemical Engineering, UniVersity of Annaba, P.O. Box 12 El Hadjar, 23000 Annaba, Algeria

In this study, the dissolution of CoO from CoO/Al2O3 catalyst with HCl, H2SO4, and HNO3 solutions was investigated in a batch reactor employing parameters that were expected to affect the dissolution rate of cobalt, such as stirring speed, temperature, and acid concentration. It was found that 99.82% of cobalt was dissolved after 4 h with HCl at 2 M, 75 °C, and a liquid-to-solid ratio (l/s) of 100 mL/g, while only 31.96% and 13.57% cobalt dissolutions were reached with H2SO4 and HNO3, respectively, under the same operation conditions. The difference in dissolution rates was due to the presence of different anions (Cl-, NO3-, and HSO4-) involved in the surface reactions. Dissolution kinetic of cobalt was examined according to a heterogeneous model. It was found that the dissolution rate was controlled by surface chemical processes in all cases. 1. Introduction Nowadays, economical and environmental requirements impose the development of effective and inexpensive methods for the recovery of valuable metals from secondary sources. In fact, recovering metals comprises ∼20%-30% of the total supply, lowering the raw materials cost by ∼15%-50%. The United States Environmental Protection Agency (EPA) states that spent catalysts may be pyroforic, spontaneously combustible, and release toxic gases. Furthermore, spent catalysts have been proven to contain high amounts of heavy metals that can be leached by water.1 For purposes of establishing a recycling process for heavy metals, leaching is considered to be the process of dissolving high-value metals with an acid or base solution. The dissolution reaction of metal oxides in acidic solutions is widely utilized industrially. The leaching of bauxite ore and oxidized copper ore, the upgrading of ilmenite ore, and the pickling of rolled steel are typical examples. However, the dissolution of other metal oxide sources such as catalysts fails to be extensively studied, among them cobalt-oxide-supported catalysts. Cobalt oxide is an attractive catalyst in air pollution control for the abatement of CO,2,3 NOx,4,5 organic pollutants from effluent streams,6 and gas desulfurization.7 The high costs of the precursor metal salts and the toxicity of cobalt render the disposal of the spent catalysts uneconomic and environmentally unacceptable. In the literature, studies dealing with leaching processes generally used one acid as a leaching reagent by which different operation conditions are optimized for a maximum dissolution efficiency of the material studied. The study of the behavior of different acids in the leaching process of one or several materials under the same operation conditions may be an interesting task to compare their efficiencies and determine which one should be used. In our previous paper,8 the dissolution of copper oxide (CuO) with several acids (HCl, H2SO4, HNO3, and C6H8O7) was studied to determine the better leachant. The results indicated that copper oxide was efficiently dissolved (99.95%) at ambient temperature (25 °C) by mineral acids at 0.5 M after 14 min, while higher temperature and concentration (80 °C, 2 M) and longer reaction time (2 h) were needed for citric acid to reach the efficiency of the mineral acids. Furthermore, among the three mineral acids studied, HCl was the more-efficient leachant since higher dissolution efficiencies * To whom correspondence should be addressed. Tel.: +213 7 71 01 88 16. Fax: +213 38 87 65 60. E-mail: [email protected].

and rate constants were obtained with it, regardless of the operation conditions applied. It was concluded that leaching efficiency depends greatly on the nature of the acid used. In this study, an acid hydrothermal leaching process was used to dissolve cobalt oxide from a CoO/Al2O3 catalyst. When cobalt oxide is contacted with the acids, the initially colorless solutions get a light-colored pink tone, indicating the presence of Co2+ that is formed according to the general reaction9 CoO(s) + 2HX f CoX2 + H2O

(1)

In the case of hydrochloric acid (HCl), sulfuric acid (H2SO4), or nitric acid (HNO3), the reaction becomes CoO + 2HNO3 f Co(NO3)2 + H2O

(2)

CoO + 2HCl f CoCl2 + H2O

(3)

CoO + H2SO4 f CoSO4 + H2O

(4)

The dissolution of cobalt oxide was tested with HCl, HNO3, and H2SO4 solutions. The effects of acid concentration, stirring speed, and temperature on the dissolution rate were evaluated using the three acids. The leaching results were analyzed on the basis of heterogeneous kinetic models, and the best-fitted equation to the experimental data was determined. 2. Experimental Section The catalyst used in this study was a 10% CoO/R-Al2O3 (n/n) prepared using the incipient wetness method. In this method, a defined volume of Co(CH3COO)2 · 4H2O (99%, Carlo Erba) with a known concentration was used to impregnate a given mass of R-Al2O3. With this technique, the load of cobalt can be calculated in a simple way. The sample was dried overnight at 110 °C and calcined in an oven at 500 °C for 2 h. Cobalt composition then was checked via atomic absorption. Scanning electron microscopy (SEM) images were taken with a low vacuum (JEOL, Model JEOL-5410) that was equipped with an energy-dispersive X-ray spectroscopy (EDS) system (Oxford Link, Model Isis 300). H2SO4 (96-98%, Biochem), HNO3 (60%, Cheminova), and HCl (37%, Carlo Erba) were used as leaching reagents. Leaching experiments were conducted in a spherical glass batch reactor of 100 mL heated by a

10.1021/ie901444y  2010 American Chemical Society Published on Web 06/22/2010

Ind. Eng. Chem. Res., Vol. 49, No. 14, 2010

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Table 1. Effect of Reaction Parameters on Cobalt Dissolution with the Three Mineral Acids for Various Reaction Times Percentage of Cobalt Dissolved (%) parameter varied

5 min

10 min

15 min

30 min

60 min

90 min

120 min

135 min

150 min

240 min

T ) 75 °C, l/s ) 100 mL/g, W ) 100 rpm, Various Acid Concentrations HCl Cacid Cacid Cacid H2SO4 Cacid Cacid Cacid HNO3 Cacid Cacid Cacid

) 0.5 M )1M )2M

1.75 2.18 3.50

3.06 3.94 6.34

4.81 5.25 10.06

5.25 6.50 19.70

7 21 40.71

11.82 26.26 64.35

15.32 39.40 78.80

18.82 47.72 81.87

21.45 51.66 99.82

) 0.5 M )1M )2M

0 0 0.87

1.18 1.75 2.18

1.94 2.18 4.17

3.37 3.94 5.40

6.12 8.30 8.5

8.75 9.79 11

10.06 11.94 13.13

11.19 17.51 20.57

16.63 26.26 31.96

) 0.5 M )1M )2M

0 0 0

0 0 0

0 1.18 2.18

1.55 2.18 3.06

2.18 4 4.37

2.72 4.81 5.85

3.74 5.69 7.88

3.94 6.56 8.87

5.25 9.63 13.57

Cacid ) 2 M, T ) 75 °C, l/s ) 100 mL/g, Various Stirring Speeds HCl W) W) W) W) H2SO4 W) W) W) W) HNO3 W) W) W) W)

100 350 600 850

rpm rpm rpm rpm

3.50 4.37 4.37 4.37

6.34 8.45 8.58 8.73

10.06 13.57 15.63 16.61

19.70 27.15 27.65 28.86

40.71 58.29 59.85 60.82

64.35 72.23 74.61 76.57

78.80 81.94 84.13 90.94

100 350 600 850

rpm rpm rpm rpm

0.87 1 1.75 2.18

2.18 2.18 2.18 3.50

4.17 4.25 4.25 4.37

5.40 7.44 7.55 8.31

8.5 11.49 13.13 16.18

11 19.70 20.85 21.87

13.13 24.95 28.45 30.62

100 350 600 850

rpm rpm rpm rpm

0 0 0 0

0 0 1.75 1.75

2.18 2.25 2.50 2.71

3.06 3.40 3.54 3.65

4.37 6.56 7 7.88

5.85 7.44 9.5 10.50

7.88 10.06 13.13 14.01

Cacid ) 2 M, W ) 850 rpm, l/s ) 100 mL/g, Various Temperatures HCl T) T) T) T) H2SO4 T) T) T) T) HNO3 T) T) T) T)

50 60 75 85

°C °C °C °C

0 0 4.37 7.44

0 1.75 8.73 14.01

1.31 2.18 16.61 24.07

2.18 3.06 28.86 56.91

3.50 9.62 60.82 72.23

6.11 13.13 76.57 84.49

7.82 20.55 90.94 96.31

8.25 21.90 94.22 99.85

50 60 75 85

°C °C °C °C

0 0 2.18 2.18

0 0 3.5 3.94

0 0 4.37 4.37

0 0 8.31 12.25

0 2.08 16.18 26.24

0 4.37 21.87 38.93

1.75 5.87 30.62 49.03

2.76 6.58 34.15 54.34

50 60 75 85

°C °C °C °C

0 0 0 1.75

0 0 1.75 2.18

0 0 2.71 3.06

0 0 3.65 4.81

0 1.75 7.88 10.70

0 2.41 10.50 15.30

1.31 3.50 14.01 18.88

1.65 4 15.34 21.11

temperature-controlled water bath and equipped with a returnflow cooler (glass condenser) to minimize water losses due to evaporation. The solutions were mixed using a magnetic stirrer to eliminate the influence of mixing and mass transfer on the kinetic results. A typical experiment was conducted as follows: 0.1 g of the sample was placed into the glass flask with 10 mL of the selected acid leading to a liquid-to-solid ratio (l/s) of 100 mL/g. After the leaching process, the reaction mixture was filtered and the Co2+ species was analyzed by titration with ethylenediamine tetraacetic acid (EDTA), using murexide as an indicator. The percentage of dissolution was calculated from the following equation: dissolution (%) ) number of moles of cobalt in the solution × 100 (5) number of moles of cobalt in the catalyst leached The data presented are an average of two test replicates, with an error of 5%. Particles of the catalyst were analyzed after leaching experiments with the three acids via SEM. Before analysis, the samples

were filtered, washed several times with distilled water to remove the impurities, and dried at 110 °C. The leaching behavior of CoO was tested under reaction conditions that were characterized by a relatively high excess of acids to eliminate possible effects of the changes in lixiviant composition during individual runs on the rate of leaching. The main parameters that may affect the dissolution of cobalt oxide, such as the nature of the acid (HCl, H2SO4, and HNO3), the acid concentration (0.5, 1, and 2 M), the stirring speed (100, 350, 600, and 850 rpm), and temperature (50, 60, 75, and 85 °C), were considered, and the best conditions for the maximum recovery were established. 3. Results and Discussion 3.1. Effect of Reagents and Conditions. To determine whether cobalt oxide dissolved easily in acid solution or not, a preliminary test was conducted with a HCl solution under the following conditions: 2 M, ambient temperature, l/s ) 100 mL/g, and without agitation. No dissolution of cobalt oxide was observed after one week. After one month, 26.26%

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Figure 1. SEM image (left) and EDS analysis (right) of the CoO/Al2O3 catalyst.

of cobalt was dissolved. Another test was performed using an agitation speed of 100 rpm for 8 h at ambient temperature. No dissolution of cobalt was observed. These imposed tests chose a high temperature (75 °C) from the beginning of the study to decrease the dissolution time. To investigate the effect of acid concentration, experiments were performed with the three acids with different concentrations (0.5, 1, and 2 M) at 75 °C, l/s ) 100 mL/g, and 100 rpm (see Table 1). It was observed that the fraction of cobalt dissolved increased as the acid concentration increased in all cases. A higher dissolution efficiency was obtained with 2 M HCl, where 99.82% of cobalt was dissolved after 4 h. After the same period of time, 31.96% and 13.57% of cobalt were dissolved with H2SO4 and HNO3, respectively. The concentration of 2 M was selected to investigate the effect of the other parameters. Alumina dissolution was checked with the three acid solutions several times at different concentrations. No detectable dissolution occurred. SEM and EDS analyses of the prepared sample (see Figure 1) show particles with almost spherical shape, composed primarily of alumina, whereas cobalt is the minor component. Figures 2a-c show SEM and EDS analyses of the catalyst after 4 h of treatment at 75 °C, l/s ) 100 mL/g, and 100 rpm with HCl, H2SO4, and HNO3, respectively. The figures show that particles of the alumina support were not affected by the treatments. This is certainly due to the chemical inertia and the high stability of R-alumina against acids.10 Also note that the particles obtained after treatments are in the form of agglomerates (∼100 µm) consisting of primary particles 2-10 µm in size. The EDS analyses indicate the presence of an increased amount of cobalt remaining on the alumina support when the sample was contacted with HCl, H2SO4, and HNO3, respectively. This is consistent with the leaching results obtained. The effect of stirring speed on the dissolution rate of cobalt was investigated using different agitation speeds (100, 350, 600, and 850 rpm) at 75 °C, l/s ) 100 mL/g, and an acid concentration of 2 M. It was observed that agitation has almost no effect on cobalt dissolution with HCl (see Table 1). In the case of H2SO4 and HNO3, despite the fact that the dissolution efficiencies remained low in both cases, compared to HCl, a certain improvement in cobalt dissolution was observed at high agitation speeds. In fact, the dissolution

efficiencies increased from 13.13% and 7.88% after 120 min at 100 rpm to 30.62% and 14.01% after the same period of time at 850 rpm with H2SO4 and HNO3, respectively. An agitation speed of 850 rpm was chosen for the subsequent experiments to ensure that the influence of external mass transfer is negligible. The effect of temperature on the rate of cobalt dissolution from CoO in the three inorganic acids was investigated over a temperature range of 50-85 °C at an acid concentration of 2 M, a stirring speed of 850 rpm, and l/s ) 100 mL/g. Table 1 shows that cobalt dissolution with the three acids was highly temperature-dependent. The time needed to reach high dissolution efficiency was decreased with increasing temperature. In the case of HCl, the dissolution efficiency reached 99.82% after 4 h at 75 °C; it attained 99.85% at 85 °C after 135 min of reaction time. In the case of H2SO4, 54.34% of cobalt was dissolved after 135 min at 85 °C, while only 31.96% was reached after 4 h at 75 °C. With HNO3, 13.57% of cobalt was dissolved after 4 h at 75 °C and it increased up to 21.11% after 135 min at 85 °C. 3.2. Kinetic Model. In heterogeneous reaction systems, most reactions follow a shrinking core model (SCM).11 The particles of the solid treated in this work are almost spherical in shape, as shown by SEM. In this model, the overall leaching process may be controlled by intrinsic chemical reaction or by external mass transfer. The following expressions can be used to describe the dissolution kinetics of the process: for liquid film diffusion control: x ) kt

(6)

for film diffusion control through the ash or product layer: 1 - 3(1 - x)2/3 + 2(1 - x) ) kt (7) for surface chemical reaction control: 1 - (1 - x)1/3 ) kt (8) where x is the fractional conversion of cobalt at time t and k is the apparent rate constant (expressed in units of min-1). The overall rate of dissolution is controlled by the slowest of these sequential steps. Figure 3 shows a good linear correlation, based on eq 8. The relationship between the overall rate constant (from eq


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Figure 2. SEM images (left) and EDS analyses (right) for (a) a CoO/Al2O3 catalyst after treatment with HCl, (b) a CoO/Al2O3 catalyst after treatment with H2SO4, and (c) a CoO/Al2O3 catalyst after treatment with HNO3.

8) and the temperature can be expressed by the Arrhenius equation:

( )

k ) A exp -

Ea RT

(9)

where A is the frequency factor (min-1), Ea the activation energy (J mol-1), R the universal gas constant (R ) 8.314 J K-1 mol-1), and T the reaction temperature (K). The activation energy in the dissolution process may be characterized to predict the rate-controlling step. The activation

energy of a diffusion-controlled process is typically 4-12 kJ/ mol, whereas for chemically controlled processes, it is usually >40 kJ/mol.12 The apparent rate constants were determined from the straight lines of Figure 3 and plotted according to the Arrhenius equation. The activation energies calculated from the slopes (Ea/(RT)) for the dissolution reactions with the three acids were Ea (HCl) ) 97.97 kJ/mol, Ea (H2SO4) ) 85.72 kJ/mol, and Ea (HNO3) ) 72.74 kJ/mol.

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Figure 3. Plots of 1- (1 - x)1/3 versus leaching time at different temperatures for the dissolution reactions of cobalt with (a) HCl, (b) H2SO4, and (c) HNO3.

Thus, eq 8 may be written for the three reactions as follows: for cobalt dissolution by HCl: 1 - (1 - x)

1/3

) 1.5 × 10

12

97.97 exp t (10) RT

[ (

)]

[ (

85.72 t (11) RT

[ (

72.74 t (12) RT

for cobalt dissolution by H2SO4: 1 - (1 - x)1/3 ) 5.9 × 109 exp -

)]

for cobalt dissolution by HNO3: 1 - (1 - x)1/3 ) 2.76 × 107 exp -

)]

The activation energy values are >40 kJ/mol, which is consistent with a chemically controlled process. In our previous paper,8 it was observed that copper oxide was almost easily dissolved by mineral acids under moderate conditions (25 °C, 0.5 M, l/s ) 10 mL/g, and without agitation), compared to cobalt oxide. In fact, 37.3%, 92%, and 97.8% of the copper oxide were dissolved with HNO3, H2SO4, and HCl, respectively, after just 6 min of reaction time, while no dissolution was observed in the case of cobalt oxide under the same leaching conditions. Furthermore, higher apparent rate constants were obtained in the case of copper oxide, as shown in Table 2. Comparison of the rate constants of copper and cobalt oxides dissolved with HCl at 50 °C shows that the value obtained in the case of copper (0.23 min-1) was greater than that obtained with cobalt (0.000 min-1). Even at higher temperature (85 °C), cobalt oxide dissolved slowly with HCl

Table 2. Comparison between Rate Constants of Copper and Cobalt Oxides Dissolved with Mineral Acids acid used

CuO

CoO

HCl HCl HCl HCl

-1

K(30 °C) ) 0.132 min K(40 °C) ) 0.187 min-1 K(50 °C) ) 0.236 min-1

K(50 °C) ) 0.0002 min-1 K(60 °C) ) 0.0006 min-1 K(75 °C) ) 0.0046 min-1 K(85 °C) ) 0.0056 min-1

H2SO4 H2SO4 H2SO4

K(30 °C) ) 0.073 min-1 K(40 °C) ) 0.152 min-1 K(50 °C) ) 0.184 min-1


HNO3 HNO3 HNO3



(0.005 min-1), compared to the lowest value obtained in the case of copper (0.07 min-1) dissolved with H2SO4 at 30 °C. These results indicate that the nature of the material leached affects the results of the dissolution. 3.3. Reaction Mechanism. Strong mineral acids are completely dissolved in water, leading to the same number of moles of H+ ions in the case of HCl, H2SO4, and HNO3 when used at the same concentration. The dissolution efficiency of oxides in such acids should be the same if H+ is responsible for the dissolution reaction. In this study, HCl dissolved 99.82% of cobalt while H2SO4 and HNO3 dissolved only 31.96% and 13.57%, respectively, under the same operation conditions (reaction time ) 4 h, acid concentration ) 2 M, temperature ) 75 °C, and stirring speed ) 100 rpm). The unique difference between these acids is the presence of three different anions,

Ind. Eng. Chem. Res., Vol. 49, No. 14, 2010 Table 3. Polarizabilities of Some Ions ion

Polarizability (cm2/mol)

ClSO42NO3ClO4-

8.9 3.9 3.6 3.3

namely, Cl-, HSO4-, and NO3-. To reach the dissolution efficiency of HCl, sulfuric and nitric acids required a prolonged reaction time and/or higher temperature. This indicates that cobalt dissolution was influenced by the behavior of the anions with the surface, rather than by the overall acidity. Thus, the difference in dissolution behavior of the three acids indicates the importance of considering the role of anions. However, very little information is available in a form suitable on the mechanism involving anions in the dissolution of metal oxides. In the scientific branches dealing with the corrosion of metals or the leaching of metals from ores, researchers are generally interested on the effect of inhibitors on the protection of metals in corrosive environments13,14 or on the kinetic studies of metals leaching in acidic or basic media.15–20 The enhanced leaching rate in the presence of chlorides has been noted by several researchers. For example, in an earlier study21 conducted on the dissolution of cupric oxide in different acidic media (HCl, H2SO4, and HNO3), the authors found that the dissolution rate of cupric oxide in HCl was ∼10 times faster than in perchloric acid. They found that the sequence of the dissolution rate of CuO in the four acids coincided with the order of polarizabilities of anions, as shown in Table 3. They explained that the strong adsorption power of an anion on a Cu site is related directly to that ion’s polarizability. The adsorption may result in a reduction of the chemical bond strength between copper and oxygen, leading to cupric oxide dissolution but with different rates, depending on the nature of the anion. Senanayake9 studied the dissolution of copper oxide (CuO) in the presence of Cl- ions and explained the dissolution on the basis of surface adsorption equilibria involving the formation of chloro-complexes as adsorbed species. The author reported that the adsorption of aqueous HCl onto the oxide surface releases some of the water in the hydration sphere of HCl to the solution to establish heterogeneous (solid/surface/solution) equilibrium for the oxide: CuO(s) + HCl(aq) T Cu(OH)Cl(ad) + ∆hH2O

(13)

In addition, the adsorbed specie Cu(OH)Cl may react with HCl to complete the reaction as shown below: Cu(OH)Cl(ads/aq) + HCl(aq) T CuCl2(aq) + H2O (14) The author indicated that the high stability of a chloro-complex would increase the value of Kads for the adsorption equilibria and, thus, enhance the rate of dissolution. In another study conducted on the covellite (CuS) dissolution,22 a low rate for the chloride-free leaching of covellite in HNO3 and H2SO4 solutions at 90 °C and an enhanced leaching upon the addition of NaCl to both acids was observed. It was concluded that the leaching was due to the involvement of the anion X- and that the complexation was stronger when X- ) Cl-, compared to X- ) HSO4- or NO3-. Lu et al.23 showed that excellent leaching kinetics of chalcopyrite (CuFeS2) existed for solutions that contained chloride, while for solutions without chloride, the leaching was very slow. Furthermore, note that the nature of the metal undergoing the dissolution in an acidic medium is of great importance. The

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dissolution rate of cobalt oxide can be considered to be negligible with HCl at ambient temperature and in the absence of agitation (static leaching), when compared to copper oxide. The cause of this behavior is not easy to explain. What can be concluded is that cobalt oxide is more stable in aggressive environments than copper oxide since high temperature (75 °C) and 4 h of reaction time were needed for HCl at high concentration (2 M) to dissolve it efficiently. On the other hand, it is known that HCl in water gives H+ and Cl-. Both are susceptible to reaction with the solid surface. Thus, it is possible that H+ initiates the dissolution reaction and Cl- can then play the role of accelerator. The contrary may also occur. The study of the role of Cl- was taken globally in the literature in the framework of the HCl molecule involving the dissolution as reported by Senanayake9 and not as an anion that initiates or promotes dissolution in the presence of H+. Therefore, it is interesting to determine the mechanism for a metal oxide dissolution with respect to anions. Thus, schemes of the stepwise dissolution of cobalt oxide in chloride, sulfate, or nitrate electrolytes are fairly complicated. The dissolution efficiency depends not only on the nature of anions but rather on the combination of several effects such as the overall acidity of the media, temperature, nature of anions, their concentration, their affinity to the solid surface, and the physicochemical characteristics of the solid. Potential Use of the Resulting Cobalt Salt Solutions. Nitrate, sulfate, and chloride cobalt solutions can be used for a further recovery of cobalt or can subsequently serve as raw material for new cobalt-based catalyst preparation. 4. Conclusion From the results obtained in this study, it can be concluded that the electrolytic environment (the presence of anions) plays an important role on cobalt dissolution in the acidic media. The presence of Cl- ions seems to enhance the dissolution rate of cobalt oxide significantly, in comparison to the presence of HSO4- or NO3- ions. In fact, 99.85% of cobalt was dissolved with HCl after 135 min at 85 °C, whereas only 54.34% and 21.11% were leached with H2SO4 and HNO3, respectively, under the same operation conditions. The dissolution of cobalt with the three acids was controlled by the chemical reaction. This result was confirmed by the high activation energies (Ea > 40 kJ/mol). Cobalt oxide, compared to copper oxide, needs more drastic conditions to be efficiently dissolved with mineral acids. Thus, the nature of the metal oxide affects greatly the results of the hydrothermal treatment applied. Acknowledgment The authors acknowledge thankfully the financial assistance of the Ministry of Higher Education and Scientific Research of Algeria (Project No. J0101120070026). The authors also would like to thank Dr. Y. Hamlaoui and Dr. F. Pedraza for SEM analyses. Literature Cited (1) Angelidis, T. N.; Tourasanidis, E.; Marinou, E.; Stalidis, G. A. Selective dissolution of critical metal from diesel and naphtha spent hydrodesulfurization catalysts. Resour. ConserV. Recycl. 1995, 13, 269. (2) Mergler, Y. J.; Hoebink, J.; Nieuwenhuys, B. E. CO oxidation over a Pt/CoOx/SiO2 catalyst: A study using temporal analysis of products. J. Catal. 1997, 167, 305. (3) Jansson, J. Low-temperature CO oxidation over Co3O4/Al2O3. J. Catal. 2000, 194, 55.

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(4) Hamada, H.; Kintaichi, Y.; Inaba, M.; Tabata, M.; Yoshinari, T.; Tsuchida, H. Role of supported metals in the selective reduction of nitrogen monoxide with hydrocarbons over metal/alumina catalysts. Catal. Today 1996, 29, 53. (5) Pietrogiacomi, D.; Tuti, S.; Campa, M. C.; Indovina, V. Cobalt supported on ZrO2: Catalysts characterization and their activity for the reduction of NO with C3H6 in the presence of excess O2. Appl. Catal., B 2000, B28, 43. (6) Garbowski, E.; Guenin, M.; Marion, M. C.; Primet, M. Catalytic properties and surface states of cobalt-containing oxidation catalysts. Appl. Catal. 1990, 64, 209. (7) Chiuping, L.; Chen, Y. W.; Yang, S. J.; Wu, J. C. Hydrodesulfurization reactions of atmospheric gas oil over cobalt-molybdenum/aluminaaluminum borate catalysts. Ind. Eng. Chem. Res. 1993, 32 (8), 1573. (8) Habbache, N.; Alane, N.; Djerad, S.; Tifouti, L. Leaching of copper oxide with different acid solutions. Chem. Eng. J. 2009, 152, 503. (9) Senanayake, G. Review of theory and practice of measuring proton activity and pH in concentrated chloride solutions and application to oxide leaching. Miner. Eng. 2007, 20, 634. (10) Al-Mansi, N. M.; Monem, N. M. A. Recovery of nickel oxide from spent catalyst. Waste Manage. 2002, 22, 85. (11) Levenspiel, O. Chemical Reaction Engineering, 3rd ed.; John Wiley & Sons: New York, 1999; p 570. (12) Abdel-Aal, E. A. Kinetics of sulfuric acid leaching of low-grade zinc silicate ore. Hydrometallurgy 2000, 55, 247. (13) Scendo, M. Inhibition of copper corrosion in sodium nitrate solutions with nontoxic inhibitors. Corros. Sci. 2008, 50, 1584. (14) Datta, J.; Samanta, B.; Jana, A.; Sinha, S.; Bhattacharya, C.; Bandyopadhyay, S. Role of Cl- and NO3- ions on the corrosion behavior of 20% SiCp reinforced 6061-Al metal matrix composite: A correlation

between electrochemical studies and atomic force microscopy. Corros. Sci. 2008, 50, 2658. (15) Sarker, M. K.; Rashid, A. K. M. B.; Kurny, A. S. W. Kinetics of leaching of oxidized and reduced ilmenite in dilute hydrochloric acid solutions. Int. J. Miner. Process. 2006, 80, 223. (16) Dutra, A. J. B.; Paiva, P. R. P.; Tavares, L. M. Alkaline leaching of zinc from electric arc furnace steel dust. Miner. Eng. 2006, 19, 478. (17) Nizamettin, D.; Ku¨nku¨l, A. Dissolution kinetics of ulexite in perchloric acid solutions. Int. J. Miner. Process. 2007, 83, 76. (18) Teir, S.; Revitzer, H.; Eloneva, S.; Fogelholm, C. J.; Zevenhoven, R. Dissolution of natural serpentinite in mineral and organic acids. Int. J. Miner. Process. 2007, 83, 36. (19) Valverde, I. M.; Paulino, J. F.; Afonso, J. C. Hydrometallurgical route to recover molybdenum, nickel, cobalt and aluminum from spent hydrotreating catalysts in sulphuric acid medium. J. Hazard. Mater. 2008, 160, 310. (20) Langova´, Sˇ.; Lesˇko, J.; Maty´sek, D. Selective leaching of zinc from ferrite with hydrochloric acid. Hydrometallurgy 2009, 95, 179. (21) Majima, H.; Awakura, Y.; Yazaki, T.; Chikamori, Y. Acid dissolution of cupric oxide. Metall. Trans. B 1980, B (11), 209. (22) Senanayake, G. A review of chloride assisted copper sulfide leaching by oxygenated sulfuric acid and mechanistic considerations. Hydrometallurgy 2009, 98, 21. (23) Lu, Z. Y.; Jeffrey, M. I.; Lawson, F. The effect of chloride ions on the dissolution of chalcopyrite in acidic solutions. Hydrometallurgy 2000, 56, 189.

ReceiVed for reView September 15, 2009 ReVised manuscript receiVed June 10, 2010 Accepted June 10, 2010 IE901444Y

Al2O3 Catalyst with Mineral Acids

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