Effect of preparation conditions on physicochemical, surface and catalytic properties of cobalt ferrite prepared by coprecipitation

Journal of Alloys and Compounds 493 (2010) 415–422

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Effect of preparation conditions on physicochemical, surface and catalytic properties of cobalt ferrite prepared by coprecipitation G.A. El-Shobaky a,∗ , A.M. Turky b , N.Y. Mostafa b , S.K. Mohamed b a b

Physical Chemistry Department, National Research Center, Dokki, Cairo, Egypt Chemistry Department, Faculty of Science, Suez Canal University, Ismailia 41522, Egypt

a r t i c l e

i n f o

Article history: Received 24 June 2009 Received in revised form 14 December 2009 Accepted 18 December 2009 Available online 28 December 2009 Keywords: Cobalt ferrite Coprecipitation Specific surface area XRD EDX H2 O2 decomposition

a b s t r a c t Cobalt ferrite nanoparticles were prepared via thermal treatment of cobalt–iron mixed hydroxides at 400–600 ◦ C. The mixed hydroxides were coprecipitated from their nitrates solutions using NaOH as precipitating agent. The effects of pH and temperature of coprecipitation and calcination temperature on the physicochemical, surface and catalytic properties of the prepared ferrites were studied. The prepared systems were characterized using TG, DTG, DTA, chemical analysis, atomic absorption spectroscopy (AAS), X-ray diffraction (XRD), energy dispersive X-ray (EDX) as well as surface and texture properties based on nitrogen adsorption–desorption isotherms. The prepared cobalt ferrites were found to be mesoporous materials that have crystallite size ranges between 8 and 45 nm. The surface and catalytic properties of the produced ferrite phase were strongly dependent on coprecipitation conditions of the mixed hydroxides and on their calcination temperature. © 2010 Published by Elsevier B.V.

1. Introduction Nanomaterials based on spinel ferrites have already numerous applications, including gas sensors [1], microwave devices [2], photocatalysis [3], adsorption technologies [4], high-frequency transformer technology [5]. Currently, magnetic nanoparticles offer widespread applications in biotechnology, such as DNA and RNA purification, cell separation, drug delivery, magnetic resonance imaging [6], magnetic hyperthermia for cancer treatments [7], thermal coagulation therapy [8], biosensors[10], biomolecular recognition and cell imaging [9]. In the field of catalysis, spinel ferrites are effective catalysts for a number of industrial processes such as decomposition of alcohols and hydrogen peroxide, oxidation of CO [11], aerobic oxidation of monoterpenic alkenes [12] and cyclohexane [13], and hydrolysis of 4-nitrophenyl phosphate [14]. Among spinel ferrites, CoFe2 O4 has attracted considerable attention in recent years due to the unique physical properties such as high Curie temperature, large magnetocrystalline anisotropy, high coercivity, moderate saturation magnetization, large magnetostrictive coefficient, excellent chemical stability and mechanical hardness [15,16]. Additionally, this material exhibits a signifi-

∗ Corresponding author. Tel.: +20 237494265; fax: +2 02 3370931. E-mail address: [email protected] (G.A. El-Shobaky). 0925-8388/$ – see front matter © 2010 Published by Elsevier B.V. doi:10.1016/j.jallcom.2009.12.115

cant higher magnetostriction than metallic Fe or Ni [17]. A lot of synthetic techniques for preparing nanosized cobalt ferrite have been presented. Among these techniques, chemical coprecipitation proved to be the most economical one [18]. The present work reports the results of a study on the effects of some preparative conditions namely, coprecipitation pH, coprecipitation temperature and calcination temperature on the physicochemical, surface and catalytic properties of cobalt ferrites prepared by thermal treatment of their mixed hydroxides at 400–600 ◦ C. The mixed hydroxides were prepared by coprecipitation from their nitrates solutions using 4 M NaOH solution. 2. Experimental 2.1. Materials All chemicals employed were of analytical grade and supplied by BDH company. Cobalt ferrites CoFe2 O4 were prepared using wet chemical coprecipitation route. The nitrates of cobalt and iron were dissolved in distilled water at the designated molar ratio (Fe/Co = 2). Aqueous solution of 4 M NaOH was used as the precipitating agent. The metal nitrate solutions and the NaOH solution were added dropwise from two separate burettes into a reaction vessels containing 1 l of distilled water under mechanical stirring. The rate of addition was controlled in order to maintain a constant pH (8 or 10) during the coprecipitation process. Coprecipitation was thermostated at the desired temperature (60 or 70 ◦ C). The precipitate was kept over night for ageing at room temperature and washed till free from NO3 − and Na+ ions. It was then filtered, dried at 110 ◦ C overnight then calcined at 400, 500 or 600 ◦ C for 4 h to achieve transformation into spinel phase.

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2.2. Techniques

3. Results and discussion

Thermogravimetric (TG), differential thermogravimetric (DTG) and differential thermal analysis (DTA) curves were obtained using a Shimazdu DTG-60H in a dynamic air atmosphere (40 ml min−1 ) at a heating rate of 10 ◦ C min−1 . The stoichiometry of the prepared systems was measured using a SOLAAR AA report, spectrometer: S series 711838 v1.26. X-ray powder diffractograms of various investigated samples calcined at 400, 500 and 600 ◦ C were determined using a Brucker diffractometer (Brucker D8 advance target), the scanning rate was fixed at 8◦ in 2 min−1 for phase identification and 0.8◦ in 2 min−1 for line broadening profile analysis. The patterns were run with Cu K␣1 with secondly monochromator ( = 0.1545 nm) at 40 kV and 40 mA. The crystallite size of crystalline phases present in different solids investigated was calculated from the line broadening of the main diffraction lines of these phases using Scherrer equation [19]:

3.1. Thermal analyses

d=

K ˇ1\2 cos 

where d is the mean crystallite diameter,  the X-ray wave length, K the Scherrer constant (0.89), ˇ1\2 the full-width at half-maximum (FWHM) of the main diffraction peak of crystalline phases and  is the diffraction angle. EDX measurements were carried out on a Hitachi S-800 electron microscope with a Kevex Delta system attached. The parameters were as follows: −15 kV accelerating voltage, 100 s accumulation time, 8 ␮m window width. The surface molar composition was determined by the Asa method (Zaf-correction, Gaussian approximation). A nitrogen adsorption system (Quantachrome NOVA Automated Gas Sorption System) was employed to measure adsorption–desorption isotherms at −196 ◦ C. The samples were degassed for 3 h at a pressure of 10−5 Torr at temperature of 350 ◦ C to remove physisorbed gases prior to the measurement. The BET approach using adsorption data over the relative pressure ranging from 0.05 to 0.3 was utilized to determine specific surface area. The (BJH) approach was used to calculate pore size distribution (v/r) of the samples. The catalytic activity of various prepared catalysts was determined using liquid phase H2 O2 decomposition at different temperatures. The kinetics of the reaction was measured gasometrically using a home-made gasometer similar to that described by Deren and Haber [20]. For each of the investigated samples a mass of 200 mg was taken for each catalytic run. The arabic numbers 1–4 denote the pH and temperature of coprecipitation. Thus, cobalt–iron mixed hydroxides obtained by coprecipitation at 60 ◦ C and pH 8 was denoted 1, the mixed hydroxides obtained by coprecipitation at the same temperature but, at pH 10 was denoted 2. Coprecipitated cobalt–iron mixed hydroxides obtained at 70 ◦ C and at pH 8 and 10 were denoted as 3 and 4, respectively.

The thermograms (TG–DTG–DTA) of samples 1–4 are shown in Fig. 1. It is shown that: (i) the TG-curves show continuous weight loss. Total weight losses of 15.8, 11.9, 16.3 and 10.1 wt.% were determined for samples 1–4, respectively. (ii) The DTA-curve of sample 1 consists of five endothermic peaks with their maxima located at 58, 294, 409, 461 and 544 ◦ C. The energy changes associated with these endothermic peaks are 169, 9, 4, 0.4, 0.8 J/g, respectively. The DTA-curve of sample 2 consists of five endothermic peaks with their maxima located at 59, 252, 313, 517 and 691 ◦ C. The energy changes associated with these endothermic peaks are 39.5, 6.7, 2.4, 5.3 and 8.2 J/g, respectively. The DTA-curve of sample 3 consists of eight endothermic peaks with their maxima located at 63, 149, 183, 279, 370, 478, 588 and 631 ◦ C. The energy changes associated with these endothermic peaks are 62.4, 0.17, 0.2, 0.77, 4.5, 0.38, 0.61 and 0.07 J/g, respectively. The DTA-curve of sample 4 consists of five endothermic peaks with their maxima located at 54, 137, 189, 249 and 356 ◦ C. The energy changes associated with these endothermic peaks are 89.9, 0.05, 0.88, 1.86 and 3.92 J/g, respectively. For all the samples, the first strong endothermic peak indicates the removal of physisorbed water and water of crystallization of some of hydrated ferric oxide. For samples 1 and 2, the second peak might correspond to the complete departure of water of crystallization of hydrated ferric oxide yielding Fe2 O3 and formation of cobalt oxide Co3 O4 . For samples 3 and 4, the second and the third peaks refer to the decomposition of cobalt hydroxide and formation of its oxide while the fourth peak indicates the formation of ferric oxide. The other small peaks might correspond to the completion of the dehydration of precursor in consecutive steps along with gradual crystallization of the produced cobalt ferrite. 3.2. Elemental analysis: Fe/Co ratio The atomic ratios (Fe/Co) of 600 ◦ C-calcined samples were determined using AAS for bulk analysis and EDX for surface analysis (Table 1). It can be seen from Table 1 that the determined Fe/Co ratio was bigger than the nominated value which could be attributed to the possibility of dissolution of small portion of the precipitated

Fig. 1. Thermograms (TG–DTG–DTA) of different cobalt–iron mixed hydroxides samples.

G.A. El-Shobaky et al. / Journal of Alloys and Compounds 493 (2010) 415–422 Table 1 Relative atomic ratio (Fe/Co) of the calcined cobalt–iron mixed hydroxides. Sample

Fe/Co ratioa (bulk)

Fe/Co ratiob (surface)

Sample 1 Sample 2 Sample 3 Sample 4

2.2 2.7 2.5 2.6

2.5 2.2 2.3 2.4

a b

Atomic absorption spectroscopy. EDX analysis.

cobalt hydroxide and/or the possible presence of some of cobalt ions in solution that have not been converted to cobalt hydroxides.

417

temperature, increased the sharpness of the XRD peaks as a result of an enhanced crystallization of the produced ferrite phase. The crystallite sizes of cobalt ferrite as calculated from the Scherrer equation using the full-width at half-maximum of the main diffraction peak (3 1 1) is shown in Table 2, which reveals the following: (i) the crystallite size of cobalt ferrite ranges between 8 and 45 nm. (ii) The crystallite size increases with the pH of the precipitating medium and the precipitation temperature. However, the effect of increasing the pH was more obvious. (iii) The crystallite size increases with increasing the calcination temperatures. (iv) It can be seen that heating from 500 to 600 ◦ C led to a drastic increase in the crystallite size indicating an effective sintering of the ferrite phase.

3.3. X-ray powder diffraction 3.4. Surface and texture studies The X-ray diffractograms of cobalt–iron mixed hydroxides powders which were calcined at different temperatures (400–600 ◦ C) are shown in Fig. 2. Examination of these diffractograms shows the following: (i) diffraction lines corresponding to cobalt ferrite spinel phase (JCPDS 22-1086) were detected in all samples prepared under different conditions. (ii) Increasing the calcination

Nitrogen adsorption–desorption isotherms conducted at −196 ◦ C over cobalt ferrite solids are shown in Fig. 3. The obtained isotherms belong to type IV isotherms according to IUPAC classification [21] and exhibited hysteresis loops of different shapes and areas. These hysteresis loops are characteristic of mesoporous

Fig. 2. X-ray diffractograms of cobalt–iron mixed hydroxides calcined at different temperatures.

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Fig. 3. Nitrogen adsorption–desorption isotherms conducted at −196 ◦ C over cobalt ferrite solids.

solids. All samples calcined at 400 and 500 ◦ C show type H2 loop according to IUPAC classification which indicate that the pore structure is complex and tend to be made up of interconnected networks of pores of different size and shape. Upon heating at 600 ◦ C, the shape of the hysteresis loops is changed from H2 to H3 revealing a change in the pore structure. Type H3 loop is usually assigned to slit-shaped pores due to non-rigid aggregates of particles. SBET and St (m2 /g) as determined from the BET method and the t-method, total pore volume Vp and mean pore radius r are listed in Table 3. Table 3 reveals the following: (i) the SBET and St values for the different cobalt ferrite samples are close to each other which justifies the correct choice of the standard t-curve used for pore analysis and shows the absence of ultra micropores. (ii) The values of SBET decrease by increasing the calcination temperature. Also, it can be noticed that, for samples 1, 3, and 4, raising the calcination temperature from 500 to 600 ◦ C is accompanied by a drastic drop in the SBET values. (iii) Sample 2 shows a different behavior towards increasing the calcination temperature since the values of the measured SBET remains virtually unchanged. It seems that this sample can resist sintering due to increasing calcination tempera-

Table 2 Effect of preparation conditions and calcination temperature on crystallite size and degree of crystallinity of the produced cobalt ferrite solids. Calcination temperature (◦ C)

Crystallite size (nm)

Degree of crystallinity (a.u.)a

Sample 1

400 500 600

8 10 38

19.7 24.5 62.5

Sample 2

400 500 600

12 14 45

20.4 25.5 68.7

Sample 3

400 500 600

8 9 42

14.4 17.7 58.5

Sample 4

400 500 600

15 17 36

20.7 22.8 45.5

Sample

a The peak area of the main diffraction line was considered as a quantitative measure for the degree of crystallinity of the phase present.

ture. (iv) Coprecipitation at lower pH value (i.e. pH 8) yielded solids with comparatively bigger specific surface areas than those measured for the solids prepared by coprecipitation at higher pH value. (v) The Vp values decrease significantly by raising the calcination temperature from 500 to 600 ◦ C. (vi) The mean pore radius of the investigated cobalt ferrite samples is in the range of 5–12 nm which confirm the nature of the prepared system as mesoporous solid. The pore size distributions of the prepared cobalt ferrite solids are shown in Fig. 4. They are tri-modal. All samples show nearly the same behavior with three peaks. The first peak ranged from 0.4 to 2.2 nm, the second peak ranged from 2.2 to 3.9 nm and the last peak ranged from 3.9 to 5.2 nm. The area of the first group (0.4–2.2 nm) is relatively small as compared to the areas of the other pore sizes. The wide distribution of pore width may account for the inter-space generated from the aggregation of nanoparticles. This Figure also demonstrates the effect of calcination temperature on the pore structure of the calcined powders. It can be seen that calcination of different adsorbents up to 500 ◦ C, the sample texture was apparently preserved since the pore size distribution has not been significantly altered. After calcination at 600 ◦ C, the pore volume was remarkably reduced reflecting a sintering process. Sample 2 has a different behavior since even after raising calcination temperature up to 600 ◦ C the pore size distribution showed only a slight change in pore volume. It is worth mentioning to report that the increase in calcination temperature within 400–600 ◦ C did not much affect the pore size distribution shape. In other words, no significant shift in the pore size to higher values took place. However, the thermal treatment at 400–600 ◦ C results in a significant decrease in the number of the pre-existing pores as being evident by the significant decrease in the area of pore size distribution curves. 3.5. Catalytic decomposition of H2 O2 over the prepared cobalt ferrite system The kinetics of the catalytic decomposition of H2 O2 at 20, 30, 40 and 50 ◦ C, using the different cobalt ferrite solids, were monitored by measuring the volumes of liberated oxygen as a function of time up to (10 min) at time intervals of 1 min. The reaction follows first order kinetics in all cases. Representative first order plots of H2 O2 decomposition over the most active catalysts are given in Fig. 5. The

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Table 3 Surface characteristics of the prepared cobalt ferrite samples. Sample

Calcination temperature (◦ C)

SBET (m2 /g)

St (m2 /g)

Vp (cc/g)

r (nm)

BET-C constant

Sample 1

400 500 600

99.3 87.1 23.9

98.8 85.7 25.3

0.249 0.254 0.150

5 5.8 12.5

36 31 97

Sample 2

400 500 600

55.8 52.14 51.6

57.1 55 53.8

0.199 0.243 0.220

7.1 9.3 8.5

69 483 18

Sample 3

400 500 600

75.7 67.6 21.4

78.8 70.7 22.7

0.243 0.232 0.113

6.4 6.8 10.5

11 77 652

Sample 4

400 500 600

61.8 57.4 29.5

63.7 59.4 30.6

0.235 0.209 0.107

7.6 7.3 7.3

68 88 50

reaction rate constants k (min−1 g−1 ) as calculated from the slopes of the first order plots are listed in Table 4. The intrinsic activity of the catalysts may become clearer if the rate constant of the catalytic reaction is expressed considering the surface area of the catalyst, i.e. k (min−1 m−2 ), column 7 of Table 4. Table 4 reveals that: (i) k (min−1 g−1 ) continuously increased with increasing the reaction temperature from 20 to 50 ◦ C. (ii) The catalytic activities of cobalt ferrite samples decreased with increasing the calcination temperature within 400 and 600 ◦ C. This finding can be attributed to the observed increase in the crystallite size

and the degree of crystallization of cobalt ferrite with subsequent decrease in its specific surface area. (iii) Sample 1 exhibited the biggest catalytic activity as evidenced by the biggest value of rate constants measured at the different reaction temperatures. (iv) Cobalt ferrite catalysts prepared at lower pH showed higher catalytic activity towards H2 O2 decomposition than catalysts prepared at higher pH. (v) The change in the coprecipitation temperature of the mixed hydroxides plays a decisive role in the catalytic activities of the produced ferrites. The increase in coprecipitation temperature from 60 to 70 ◦ C brought about a significant decrease in the

Fig. 4. The pore size distribution curves for cobalt ferrite solids.

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Fig. 5. First-order plots of H2 O2 decomposition conducted at different temperatures over cobalt ferrite catalyst produced from sample 1 and calcined at 400, 500 or 600 ◦ C.

catalytic activities expressed as k and k . The decrease was, however, more pronounced for the k value. Determination of the apparent activation energy (E) for H2 O2 decomposition over cobalt ferrite catalysts shed some light on the

possible change in the mechanism of the catalyzed reaction and hence gives useful information about possible changes in the concentration and nature of the catalytically active constituents. (E) values were calculated by direct application of the Arrhenius equa-

Table 4 Reaction rate constant (k) and rate constant per unit surface area (k ) of liquid phase H2 O2 decomposition over the prepared cobalt ferrite catalysts. Sample

Calcination temperature

k, 20 ◦ C

k, 30 ◦ C

k, 40 ◦ C

k, 50 ◦ C

k , 30 ◦ C

Sample 1

400 500 600

2.04 1.32 0.26

2.96 1.78 0.44

3.43 2.07 0.72

4 2.9 0.86

0.03 0.02 0.018

Sample 2

400 500 600

0.96 0.57 0.21

1.69 1.01 0.37

2.18 1.52 0.77

2.64 1.87 0.98

0.03 0.019 0.007

Sample 3

400 500 600

0.74 0.65 0.48

1.1 0.92 0.64

1.98 1.4 1.05

2.55 2.54 1.15

0.015 0.014 0.03

Sample 4

400 500 600

0.61 0.4 0.42

1.02 0.71 0.54

1.89 1.39 0.8

2.19 1.56 0.98

0.016 0.012 0.018

k, values expressed in (min−1 g−1 ). k , values expressed in (min−1 m−2 ).

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Table 5 Apparent activation energies E, E* and frequency factor A of liquid phase H2 O2 decomposition conducted over the prepared cobalt ferrite catalysts. Sample

Calcination temperature (◦ C)

E (kJ/mol)

ln A

E* (kJ/mol)

Sample 1 Sample 2 Sample 3 Sample 4

400

17 26 34 35

7.84 10.77 13.81 14.06

17 19 20 20

Sample 1 Sample 2 Sample 3 Sample 4

500

20 32 35 38

8.65 12.5 13.96 14.75

20 22 22 23

Sample 1 Sample 2 Sample 3 Sample 4

600

32 42 24 23

11.96 15.74 9.29 8.45

32 33 31 31

tion and are given in Table 5 which also includes the values of the pre-exponential factor A of the Arrhenius equation. The changes in the value of ln A are normally followed by corresponding changes in the magnitude of (E). The catalytic activities of all catalysts investigated ran parallel to the changes in (E) values. In fact, the smaller the (E) values the bigger the catalytic activities and vice versa. So, the computed (E) values express the results obtained. In order to account for the changes in ln A value by changing the calcinations temperature and preparation conditions of various catalysts the (E) values were recalculated (E*) by adopting the ln A value of the catalyst 1 calcined at 400,500 and 600 ◦ C to the other catalysts (2, 3 and 4) being calcined at the same temperatures. The computed (E*) values are given in the last column of Table 5. Examination of Table 5 shows that the change in calcination temperature of various catalysts within 400 and 600 ◦ C did not change the activation energy of the catalyzed reaction. So, one may conclude that increasing the calcination temperature of different catalysts within 400–600 ◦ C did not change the mechanism of the catalyzed reaction but changed the concentration of active sites participated in the catalyzed reaction. These assumption find evidence from the plots of the equation F(Ei ) = a exp(hEi ), where Ei is the energy of site “i” with the substrate. A = a exp(hE) derived on the basis of dissipation function of active sites by their energy as a consequence of surface heterogeneity[26]. Fig. 6 shows the plots of ln A verses E for all catalysts calcined at 400, 500 and 600 ◦ C. The computed values of the constant “h” (slope) of the plots in Fig. 6 are 0.35, 0.34 and 0.37 mol kJ−1 , respectively. The “a” values calculated for the previous solids are 1.82, 1.75 and 0.11 min−1 , respectively. The constant “h” and “a” values indicate that the increase in calcination temperature of different solids within 400–600 ◦ C did not change the dissipation of active sites, i.e. the character of surface heterogeneity. In other words, this treatment did not change the energetic nature of active sites but changed their concentration. Examination of Table 5 shows the following: (i) the prepared cobalt ferrite catalysts showed small values of apparent activation energies as compared to the un-catalyzed H2 O2 decomposition (76 kJ/mol [22]). (ii) The values of (E*) of the prepared catalysts are found to be virtually the same within the experimental error for various investigated catalysts calcined at various temperatures. These results indicate that changing both pH and temperature of coprecipitation did not modify the mechanism of H2 O2 decomposition but brought about an increase in the concentration of the catalytically active sites present in the outermost surface layers of the catalysts [23,24]. This can be discussed in terms of creation of oxygen vacancies on the surface of the catalysts. More oxygen vacancies may facilitate the adsorption of H2 O2 which result in increasing the catalytic activity [25]. It may be suggested that coprecipitation at 60 ◦ C during the preparation of samples

Fig. 6. Variation of ln A verses E for the catalytic reaction carried out over different catalysts calcined at 400, 500 and 600 ◦ C.

1 and 2 led to creation of more oxygen vacancies at the surface. 4. Conclusions The results obtained permitted to draw the following conclusions: 1. Thermal treatment of the cobalt–iron mixed hydroxides coprecipitated at different pH values and various coprecipitation temperatures, led to formation of nanocrystalline cobalt ferrite spinel phase at temperature starting from 400 ◦ C.

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2. The crystallite size of cobalt ferrite increases with increasing the pH and the coprecipitation temperature. 3. The crystallite size and the degree of crystallinity of cobalt ferrite increased by increasing the calcination temperature, while SBET values were decreased. 4. Cobalt ferrite solids were mesoporous materials with specific surface areas ranging between 99.3 and 21.4 m2 /g. 5. The catalytic activity of cobalt ferrite catalysts was decreased by increasing the temperature and pH of the coprecipitation medium. Acknowledgment Financial support from Suez Canal University is gratefully acknowledged. References [1] N. Rezlescu, C. Doroftei, E. Rezlescu, P.D. Popa, Sens. Actuators B 133 (2008) 420. [2] V.G. Harris, A. Geiler, Y. Chen, S.D. Yoon, M. Wu, A. Yang, Z. Chen, P. He, P.V. Parimi, X. Zuo, C.E. Patton, M. Abe, O. Acher, C. Vittoria, J. Magn. Magn. Mater. 321 (2009) 203. [3] H. Yang, J. Yan, Z. Lu, X. Cheng, Y. Tang, J. Alloys Compd. 476 (2009) 715. [4] A. Kraus, K. Jainae, F. Unob, N. Sukpirom, J. Colloids Interface Sci. 338 (2009) 359. [5] K. Praveena, K. Sadhana, S. Bharadwaj, S.R. Murthy, J. Magn. Magn. Mater. 321 (2009) 2433.

[6] T. Tanaka, R. Shimazu, H. Nagai, M. Tada, T. Nakagawa, A. Sandhu, H. Handa, M. Abe, J. Magn. Magn. Mater. 321 (2009) 1417. [7] S.W. Lee, S. Bae, Y. Takemura, I.B. Shim, T.M. Kim, J. Kim, et al., J. Magn. Magn. Mater. 310 (2007) 2868. [8] H. Hirazawa, S. Kusamoto, H. Aono, T. Naohara, K. Mori, Y. Hattori, T. Maehara, Y. Watanabe, J. Alloys Compd. 461 (2008) 467. [9] M. Hatakeyama, Y. Mochizuki, Y. Kita, H. Kishi, K. Nishio, S. Sakamoto, M. Abe, H. Handa, J. Magn. Magn. Mater. 321 (2009) 1364. [10] M. Pita, J.M. Abad, C.V. Dominguez, C. Briones, E.M. Martí, J.A.M. Gago, M.P. Morales, V.M. Fernández, J. Colloids Interface Sci. 321 (2008) 484. [11] P.T.A. Santos, A.C.F.M. Costa, R.H.G.A. Kiminami, H.M.C. Andrade, H.L. Lira, L. Gama, J. Alloys Compd. 483 (2009) 399. [12] L. Menini, M.C. Pereira, L.A. Parreira, J.D. Fabris, E.V. Gusevskaya, J. Catal. 254 (2008) 355. [13] J. Tonga, L. Boc, Z. Li, Z. Lei, C. Xia, J. Mol. Catal. A 307 (2009) 58. [14] J. Akla, T. Ghaddara, A. Ghanemb, H. El-Rassy, J. Mol. Catal. A 312 (2009) 18. [15] P.C. Rajath, R.S. Manna, D. Banerjee, M.R. Varma, K.G. Suresh, A.K. Nigam, J. Alloys Compd. 453 (2008) 298. [16] R.K. Sharma, O. Suwalka, N. Lakshmi, K. Venugopalan, A. Banerjee, P.A. Joy, J. Alloys Compd. 419 (2006) 155. [17] M. Gharagozlou, J. Alloys Compd. 486 (2009) 660. [18] Q. Chen, Z.J. Zhang, Appl. Phys. Lett. 73 (1998) 3156. [19] B.D. Cullity, Elements of X-ray Diffraction, Third ed., Addison-Wesley, Reading, 1967. [20] J. Deren, J. Haber, J. Catal. 4 (1965) 22. [21] F. Rouquerol, J. Rouquerol, K. Sing, Adsorption by Powders and Porous Solids: Principles, Methodology and Applications, Academic Press, San Diego, 1999, p. 19. [22] P. Atkins, J. De Paula, Atkins’ Physical Chemistry, 8th edition, Oxford University Press, 2006, p. 839. [23] A.M. Turky, N.R.E. Radwan, G.A. El-Shobaky, Colloids Surf. A 181 (2001) 57. [24] W.M. Shaheen, A.A. Zahran, G.A. El-Shobaky, Colloids Surf. A 231 (2003) 51. [25] D. Guin, B. Baruwati, S.V. Manorama, J. Mol. Catal. A: Chem. 242 (2005) 26. [26] A. Balandin, A.A. Dok, Akad. Nauk. SSSR 93 (1953) 55.