Investigations on the Solubility, Density, and Viscosity in the NaVO3+

Article pubs.acs.org/jced

Investigations on the Solubility, Density, and Viscosity in the NaVO3 + Na2SO4 + H2O System from 293.15 K to 323.15 K Krzysztof Mazurek* and Sebastian Druzẏ ński Faculty of Chemistry, Nicolaus Copernicus University, ul. Gagarina 7, 87-100 Toruń, Poland ABSTRACT: The solubility of sodium vanadate(V) and sodium sulfate(VI) in the system NaVO3 + Na2SO4 + H2O within a temperature range from (293.15 to 323.15) K was measured with the method of isothermal saturation of solutions. The concentration of ions in solutions was determined using the EDXRF method. The part of the polythermal solubility surface for the investigated system was plotted based on the obtained data. Additionally, the solution density and viscosity was investigated. Knowledge of these data is necessary to evaluate the possibility of the use of vanadium compounds from a spent vanadium catalyst for the production of Na2CO3, and to assess the possibility of selective precipitation of vanadium compounds from a solution after leaching the waste.

1. INTRODUCTION Spent vanadium catalyst is among the most hazardous wastes of the sulfuric(VI) acid industry. As a result of change in the structure and texture of catalysts caused by operation in an industrial environment, they are subject to deactivation, and therefore must be withdrawn from industrial use. The coefficient of relative deactivation, which is a measure of decline in operating activity of catalysts, is a function of time and operating conditions of catalysts.1−6 Regardless of the type of plant in which the catalysts are used, vanadium, potassium, or sodium compounds, sulfates(VI), and free sulfur oxide(VI) are the basic components of waste vanadium. Owing to the presence of free SO3, sulfates(VI), and moisture, the pH of the effluent is acidic and can be about pH 1 to pH 2. Therefore, stored catalysts may be a source of environmental hazard. The risk is mainly that harmful chemicals found in catalysts may seep into soil and water. The acidity of the effluent may systematically rise, which would increase the mobility of pollutants contained in the waste. Previous works indicated two possibilities of spent vanadium catalyst management. The first option is to use it to obtain Na2CO3 by applying the vanadate method which involves carbonization of water−ammonia solutions of sodium vanadate(V).7−14 In this method, the spent vanadium catalyst would be a valuable and cheap source of vanadium(V) and could be used for the synthesis of sodium vanadate(V) from NaCl and with the participation of water vapor or oxygen in accordance with the following equations:12,13

then to separate vanadium compounds from the leach solution in the form of a marketable product (e.g., NaVO3), or to use the leach solution to produce fresh catalyst.15−19 After prior stabilization and cementation, the postextraction solid waste, consisting mainly of SiO2, can be used as a building material in road construction or mining. These methods of spent vanadium catalyst management require economic development and environmentally friendly ways of separating synthesized sodium vanadate(V) from the postreaction mixture and separating sodium vanadate(V) from the leach solution. Therefore, it is essential to determine the mutual water solubility of the vanadium(V) salt and sulfates(VI),20−22 including the three-component system of NaVO3 + Na2SO4 + H2O.

2. EXPERIMENTAL DETAILS Reagents and Apparatus. Reagents used in the tests (Na2SO4, ≥ 99 wt %, POCh Gliwice, Poland; NaVO3, ≥ 98 wt %, Aldrich Chemical Co., Inc.) were of analytical grade. Deionized water was used for preparing the solutions and chemical analysis. The experiments were carried out in a thermostatic bath in which samples were magnetically stirred until the equilibrium between the solution and precipitate was guaranteed at each respective temperature. The temperature was kept constant with the Polystat CC1 thermorelay with an accuracy of ± 0.02 K. PANalytical’s MiniPal 4 compact energy dispersive X-ray spectrometer was employed to determine the concentrations of V and S in solution. Philips X’Pert PRO X-ray powder diffractometer was used to identify the crystal structures of solid phases. The automatic viscometer Lauda iVisc was used to measure the kinematic viscosity of the test equilibrium solutions.

2NaCl + V2O5(spent catalyst) + 0.5O2 → 2NaVO3 + Cl 2↑ (1)

2NaCl + V2O5(spent catalyst) + H 2O → 2NaVO3 + 2HCl↑ (2)

Received: November 16, 2013 Accepted: February 27, 2014 Published: March 18, 2014

Another possibility is to recover vanadium compounds by leaching with the use of various media, including NaOH, and © 2014 American Chemical Society

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dx.doi.org/10.1021/je4009969 | J. Chem. Eng. Data 2014, 59, 1468−1475

Journal of Chemical & Engineering Data

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Table 1. Solubility, Density, and Viscosity Data of the System NaVO3 + Na2SO4 + H2O at a Temperature Range of (293.15 − 323.15) K and Pressure of 0.1 MPaa ρ g·cm

c/mol·kg−1

ν −3

2 −1

mm ·s

x

NaVO3

Na2SO4

1.120 1.077 1.084 1.105 1.127 1.152 1.153 1.152 1.151 1.150 1.149 1.146

1.274 1.199 1.221 1.326 1.434 1.590 1.597 1.591 1.589 1.583 1.581 1.576

1.455 0.864 0.438 0.262 0.191 0.138 0.131 0.118 0.107 0.073 0.040 0.000

0.00 0.24 0.58 0.86 1.05 1.32 1.36 1.37 1.38 1.38 1.37 1.37

1.139 1.127 1.130 1.136 1.145 1.158 1.171 1.193 1.212 1.235 1.258 1.283 1.284 1.283 1.282 1.281 1.281 1.279

1.110 1.048 1.066 1.121 1.177 1.251 1.325 1.455 1.581 1.767 1.969 2.240 2.252 2.251 2.242 2.229 2.216 2.211

1.719 1.445 1.235 0.962 0.717 0.510 0.398 0.261 0.207 0.142 0.098 0.083 0.077 0.069 0.054 0.046 0.034 0.000

0.00 0.30 0.51 0.79 1.06 1.37 1.54 1.91 2.08 2.39 2.66 2.71 2.78 2.79 2.80 2.77 2.81 2.81

1.155 1.151 1.147 1.159 1.179 1.208 1.244 1.307 1.322 1.321 1.320 1.319 1.318 1.317

1.006 0.911 0.951 1.021 1.113 1.261 1.482 2.007 2.153 2.150 2.148 2.146 2.144 2.138

1.951 1.504 1.155 0.760 0.483 0.287 0.182 0.124 0.074 0.071 0.049 0.039 0.035 0.000

0.00 0.41 0.75 1.16 1.55 2.05 2.48 2.85 3.35 3.40 3.39 3.39 3.43 3.42

1.173 1.158 1.157 1.166 1.174 1.19 1.202 1.208 1.234 1.259

0.956 0.877 0.850 0.878 0.911 0.980 1.013 1.053 1.178 1.360

2.286 1.894 1.419 0.952 0.725 0.490 0.422 0.366 0.294 0.203

0.00 0.34 0.78 1.13 1.35 1.72 1.85 2.00 2.31 2.71

NaVO3 T = 293.15 K 1.000 0.782 0.428 0.234 0.153 0.095 0.088 0.079 0.072 0.050 0.028 0.000 T = 303.15 K 1.000 0.829 0.708 0.549 0.404 0.272 0.205 0.120 0.090 0.056 0.036 0.030 0.027 0.024 0.019 0.016 0.012 0.000 T = 313.15 K 1.000 0.786 0.605 0.397 0.238 0.122 0.068 0.042 0.022 0.021 0.014 0.011 0.010 0.000 T = 323.15 K 1.000 0.847 0.646 0.458 0.349 0.222 0.186 0.155 0.113 0.070

1469

Na2SO4

solid phase

0.000 0.218 0.572 0.766 0.847 0.905 0.912 0.921 0.928 0.950 0.972 1.000

NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O, Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O

0.000 0.171 0.292 0.451 0.596 0.728 0.795 0.880 0.910 0.944 0.964 0.970 0.973 0.976 0.981 0.984 0.988 1.000

NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O, Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O Na2SO4·10H2O

0.000 0.214 0.395 0.603 0.762 0.878 0.932 0.958 0.978 0.979 0.986 0.989 0.990 1.000

NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O, Na2SO4 Na2SO4 Na2SO4 Na2SO4 Na2SO4 Na2SO4

0.000 0.153 0.354 0.542 0.651 0.778 0.814 0.845 0.887 0.930

NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O NaVO3·2H2O



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Table 1. continued ρ g·cm

c/mol·kg−1

ν −3

1.305 1.308 1.306 1.304 1.303 1.303 1.303

2 −1

mm ·s

1.566 1.644 1.639 1.638 1.633 1.632 1.626

NaVO3 0.102 0.089 0.070 0.065 0.057 0.039 0.000

x Na2SO4

NaVO3 T = 323.15 K 0.031 0.027 0.021 0.019 0.017 0.012 0.000

3.21 3.24 3.27 3.29 3.28 3.28 3.31

Na2SO4

solid phase

0.969 0.973 0.979 0.981 0.983 0.988 1.000

NaVO3·2H2O NaVO3·2H2O, Na2SO4 Na2SO4 Na2SO4 Na2SO4 Na2SO4 Na2SO4

Notation: x, mole fraction of salts in solution on a water free basis; c, mol·kg−1H2O. The uncertainty of T is ± 0.02 K, ρ is ± 0.002 g·cm−3, ν is ± 0.8 %, c is ± 2.5 %. a

Table 2. Comparison of Pure NaVO3 Solubility in Water Trypuć and Kiełkowska8

this work

Trypuć and Białowicz10

T/K

ρ/g·cm−3

c/g·100 g of soln−1

ρ/g·cm−3


ρ/g·cm−3


293.15 303.15 313.15 323.15

1.120 1.139 1.155 1.173

15.07 17.33 19.21 21.80

1.124 1.138 1.155 1.179

15.05 17.25 19.25 21.64

1.124 1.138 1.151 1.179

15.08 17.25 19.28 21.61

Table 3. Comparison of Pure Na2SO4 Solubility in Water Stephen and Stephen26

this work T/K

ρ/g·cm−3


293.15 303.15 313.15 323.15

1.146 1.279 1.317 1.303

16.27 28.54 32.69 32.02

ρ/g·cm−3

Okorafor27


ρ/g·cm−3


16.3 29.0 32.8 31.8

1.15 1.28 1.35 1.29

15.93 28.89 32.48 31.65

Figure 1. Course of branch I (corresponding to the saturated solutions of Na2SO4) of the solubility polytherm in the NaVO3 + Na2SO4 + H2O system. 1470



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The capillary Lauda Micro-Ubbelohde, calibrated and certified with constant k = 0.03, was used in the study. Experimental Procedure. Erlenmeyer flasks, each of 100 cm3 volume, containing the adjusted amount of components and 70 cm3 of water were placed in a thermostatted bath until equilibrium between the solid phase and solution was gained. The samples were stirred for at least 72 h. After a fixed time interval, stirring was discontinued and solids were allowed to settle (next 24 h). A clear equilibrium solution was transferred into a calibrated at given temperature Ostwald’s pycnometer. The overall pycnometer content was used for the solution density determination. The accuracy of that determination was ± 0.002 g·cm−3. Next, the solution was transferred quantitatively to the graduated flask (500 cm3) and diluted with deionized water, and the vanadium and sulfur concentration were examined. The slight under-pressure conditions were applied during the transfer of the pycnometer contents to the flasks. A clear solution was drawn into the Ubbelohde capillary to determine the kinematic viscosity of the equilibrium solution. Prior to sampling and measurement, the capillary was thermostatted for an appropriate period of time at a given temperature to prevent crystallization from the solution. Liquid flow time was measured at least three times, and the viscosity measurement uncertainty was ± 0.8 %. Analytical Methods. The total concentration of vanadium and sulfur in solutions was determined using the EDXRF method.23 All measurements were performed in triplicate and the uncertainty of the analysis was estimated to be ± 2.5 %. The qualitative analysis of the solid phase for selected points was carried out by XRD. Diffraction patterns obtained were then compared with data available in the published databases.24,25

Figure 2. Course of branch II (corresponding to the saturated solutions of NaVO3) of the solubility polytherm in the NaVO3 + Na2SO4 + H2O system.

3. RESULTS AND DISCUSSION The results of a quantitative analysis of solutions relating to the mutual solubility of NaVO3 and Na2SO4 are summarized in Table 1. The columns of the table present the solution density (g· cm−3), kinematic viscosity of the solution (mm2·s−1), concentration of the system components expressed in (mol·kg−1) and mole fractions, and the composition of the solid phase. The mole fractions have been calculated as x=

[A] [A] + [B]

(3)

Figure 3. Density relationship of Na2SO4 saturated equilibrium solutions versus the salts amount expressed in mole fractions.

where A and B are NaVO3 and Na2SO4, respectively. The levels of solubility of pure NaVO3 and Na2SO4 presented in Table 1 are summarized and compared with results of previously published studies (Tables 2 and 3). The maximum differences in solubility do not exceed 1% for NaVO3 and 2.1% for Na2SO4. The results summarized in Table 1 have been used to present an isothermal section for NaVO3 + Na2SO4 + H2O, shown in Figures 1 and 2, and the correlation between the density of the solutions and the concentration of sodium vanadate(V) expressed in mole fractions are in Figures 3 and 4. Each isotherm consists of two branches. Figure 1 shows the location of points on branch I which correspond to solutions saturated with sodium sulfate(VI). These branches begin at points A corresponding to solutions saturated with Na2SO4 at a given temperature, and end at points E. Branches II, presented in Figure 2, describe changes in the solubility of sodium vanadate(V) with increasing concentration

of Na2SO4 in the equilibrium solution. They begin at points B which correspond to solutions saturated with NaVO3 at a given temperature and end at points E. Points E are points of intersection of branches I and II and correspond to solutions saturated with respect to both NaVO3 and Na2SO4. The data presented in Table 1 indicate that the solubility of sodium vanadate(V) grows linearly with increasing temperature. The analysis of water solubility of Na2SO4 depending on temperature shows that the solubility of the compound grows with increasing temperature (293.15 to 313.15) K. However, at 323.15 K, there is a minimal decrease in the solubility of the compound relative to the value reached at 313.15 K. As is apparent from Figure 1, the presence of NaVO3 does not affect the solubility of sodium sulfate(VI). The concentration of 1471



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strong salting-out effect of sodium sulfate(VI). Even a small addition of this salt to a solution saturated with sodium vanadate(V) results in a significant reduction of NaVO3 concentration in the equilibrium solution. After crossing points E, the densities and viscosities of the equilibrium solutions slightly decrease reaching density and viscosity values of the solutions saturated with Na2SO4. Sodium vanadate(V) in the form of β-NaVO3 may additionally be in the solid phase coexisting with the equilibrium solution for points on branch II. The hydrated form is hardly stable and even at room temperature is readily converted into the isomeric form of β-NaVO3. Trypuć and Białowicz10 showed that solid phase samples analyzed immediately after collection contained only the dihydrate form. However, the isomer β-NaVO3 was found in samples analyzed after 24 h. The so-called “melting point” of sodium sulfate(VI) occurs at a temperature of 305.55 K. Above this temperature, the anhydrate is the stable phase and below this temperature it is the decahydrate salt. Literature data, including von Plessen27,28 and Rossenblatt et al.,29 clearly show that at 303.15 K only the decahydrate is the stable phase. However, there are literature reports that confirm the presence of the anhydrous salt at 303.15 K. Conducting research on the NaCl + Na2SO4 + H2O system at 303.15 K,30 W. F. Linde noted the existence of two points E, transition from the dehydrated to the anhydrous form and the clear salting-out effect of sodium chloride on the solubility of sodium sulfate(VI). It should also be noted that the concentration of NaCl at the eutonic point in the system studied by Linde is 6 mol·kg−1. In the ternary system presented herein, the concentration of NaVO3 at point E, at a temperature of 303.15 K, is less than 0.08 mol·kg−1, while the isothermal curve at this temperature corresponding to other temperatures, and the lack of effect of sodium vanadate(V) on the solubility of sodium sulfate(VI) indicate that there is no new solid phase in the test solution. Mathematical equations have been developed based on experimental results obtained to allow calculation of the concentration of the ternary system components at temperatures other than those presented herein, yet falling within the range of (293.15 to 323.15) K. To describe the course of branch I, the following notations have been used:

Figure 4. Density relationship of NaVO3 saturated equilibrium solutions versus the salt amounts expressed in mole fractions.

Na2SO4 practically does not change for increasing concentrations of NaVO3 over the range of temperatures. The curve of points in Figure 2 indicates that as the concentration of sodium sulfate(VI) increases in the solution, the concentration of NaVO3 systematically undergoes a large reduction in the direction of points E, which shows a strong salting-out effect of Na2SO4 on the water solubility of sodium vanadate(V). The effect of the decreasing solubility of NaVO3 caused by the presence of Na2SO4 in the solution increases with temperature and it is 2.197 mol·kg−1 at 323.15 K. Figure 5 compares the levels of water solubility of sodium vanadate(V) and at points E of the system studied.

concentration of Na2SO4 in mol·kg−1: [Na 2SO4 ] = f ([NaVO3], T )

concentration of Na2SO4 in water in mol·kg−1: [Na 2SO4 ]A = f (T )

concentration of Na2SO4 at point E in mol·kg−1: [Na 2SO4 ]E = f ([NaVO3], T )

concentration of NaVO3 at point E in mol·kg−1: Figure 5. Comparison of the levels of water solubility of sodium vanadate(V) and at points E of the system studied.

[NaVO3]E = f (T )

concentration of NaVO3 in mol·kg−1: The analysis of the experimental data provided in Table 1 and the course of points in Figures 3 and 4 show that the density and viscosity of the equilibrium solutions initially decrease and then increase significantly with increasing concentration of sodium sulfate(VI) (for the solutions saturated with NaVO3), reaching a maximum at points E. The initial decrease in the density and kinematic viscosity of the equilibrium solutions is due to a very

[NaVO3] temperature in K

T Temperature-dependent changes in the concentration of [Na2SO4]A are best described in the polynomial equation 1472


Journal of Chemical & Engineering Data [Na 2SO4 ]A = a0 + b0T + c0T 2

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linear, regardless of temperature and the concentration of NaVO3 in the solution equilibrium. The temperature-dependent change in the solubility of sodium sulfate(VI) is the only determinant of differences in the position of isotherm branches I.

(4)

Dependence of [Na2SO4]E on both temperature and [NaVO3]E can be shown by the equation [Na 2SO4 ]E = (a1 + b1T ) + (c1 + d1T )[NaVO3]E

(5)

[Na 2SO4 ] = [Na 2SO4 ]A = a0 + b0T + c0T 2

Dependence of [NaVO3]E on temperature can be shown as the following polynomial equation: [NaVO3]E = a 2 + b2T + c 2T 2 + d 2T 3

(9)

The relative error of the results calculated from eq 9 and the experimental results in 96 % is less than ± 2.0 %. The summary of experimental results and those calculated from eqs 4, 7, and 9 describing branches I and the test system are shown in Figure 6.

(6)

By substituting eq 6 to eq5, we obtain an equation describing changes in the concentration of sodium sulfate(VI) at points E, depending on temperature [Na 2SO4 ]E = (a1 + b1T ) + (c1 + d1T )(a 2 + b2T + c 2T 2 + d 2T 3)

(7)

All of the above equation coefficients have been calculated on the basis of the results of experimental studies and summarized in Table 4. Table 4. The Multiple Correlation Coefficients ao b0 c0 a1 b1 c1 d1 a2 b2 c2 d2 a3 b3 a4 b4 c4 d4 a5 b5 c5 a6 b6 c6 e6 f6 g6 i6 j6 k6

−380.48 2.4262 3.8328·10−3 −28.72 0.1103 219.36 −0.8073 177.436 1.67326 5.257·10−3 5.5·10−6 −6.544 2.725·10−2 1.057 −2.96·10−3 −4.287·10−1 1.316·10−3 −379.28 2.42 −3.8259·10−3 −329.8697 2.1037 3.365·10−3 281.893 1.79265 2.8506·10−3 −56.9929 3.62151·10−1 5.7501·10−4

Figure 6. The studied system with the experimental points and the approximating lines calculated from eqs 4, 7, and 9: branches I.

In the approximation of branches II the following notations have been used: concentration of NaVO3 in mol·kg−1: [NaVO3] = f ([Na 2SO4 ], T )

concentration of pure NaVO3 in water in mol·kg−1: [NaVO3]B = f (T )

concentration of NaVO3 at the E point in mol·kg−1: [NaVO3]E = f ([Na 2SO4 ], T )

concentration of Na2SO4 at point E in mol·kg−1: [Na 2SO4 ]E = f (T )

concentration of Na2SO4 in mol·kg−1:

[Na 2SO4 ]

The relative deviation of the approximation has been calculated from eq 8: calculated value − experimental value δ (%) = 100 experimental value

temperature in K:

T Dependence of the solubility of sodium vanadate (V) on temperature is best described in the linear equation

(8)

The relative deviations calculated from eq 8 give for eq 4: (−0.8 to 0.3) %; for eq 5: (−0.9 to −0.1) %; for eq 6: (0.8 to 1.6) % and for eq 7: (−2.1 to −1.3) %. On the basis of the position of points on branches I of the solubility isotherms (Figure 1), it has been assumed that they are

[NaVO3]B = a3 + b3T

(10)

Dependence of [NaVO3]E on both temperature and [Na2SO4]E can be shown in the linear equation 1473



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[NaVO3]E = (a4 + b4T ) + (c4 + d4T )[Na 2SO4 ]E

expressed in mole fraction was plotted. Additionally the approximation of the experimental data on the mutual solubility with the use of mathematical functions was presented. The results have shown, that sodium sulfate(VI) has a strong saltingout effect on NaVO3 in its saturated solutions. The presence of sodium vanadate(V) does not affect the sodium sulfate(VI) solubility. The density and viscosity of the equilibrium solutions depends on the concentration of sodium sulfate(VI). No new double salts are formed in the investigated system.

(11)

Dependence of [Na2SO4]E on temperature can be shown as the polynomial [Na 2SO4 ]E = a5 + b5T + c5T 2

(12)

By substituting eq 12 to eq 11, we obtain the following dependence which allows calculation of the concentration of sodium vanadate(V) at points E

■

[NaVO3]E = (a4 + b4T ) + (c4 + d4T )(a5 + b5T + c5T 2) (13)

The relative deviations calculated from eq 8 gave for eq 10: (−1.1 to 2.0) %, for eq 11: the error is δ = 0, for eq 12 (−1.3 to 0.4) % and for eq 13: (0.0 to 1.4) %. On the basis of the course of points located on isotherm branches II (Figure 2), it has been assumed that the solubility of NaVO3 depending on the concentration of sodium sulfate(VI) and temperature is best described in the polynomial equation:

Corresponding Author

*E-mail: [email protected]. Tel.: (0048)566114309. Fax: (0048)566542477. Funding

The project was funded by the National Science Centre, Poland. Notes

The authors declare no competing financial interest.

2

■

[NaVO3] = [NaVO3]B + (a6 + b6T + c6T )[Na 2SO4 ] + (e6 + f6 T + g6T 2)[Na 2SO4 ]2 2

3

+ (i6 + j6 T + k6T )[Na 2SO4 ]

AUTHOR INFORMATION

REFERENCES

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(14)

The approximation for ca. 75 % of the experimental data gave the relative error not exceeding a few percent. Only single points had fallen out of that range. The lines approximating branches II for the Na2SO4− NaVO3−H2O system and the experimental points are presented in Figure 7.

Figure 7. The studied system with the experimental points and the approximating lines calculated from eqs 10, 13, and 14: branches II.

4. CONCLUSION The solubility of sodium vanadate(V) and sodium sulfate(VI) in the system NaVO3 + Na2SO4 + H2O within a temperature range from (293.15 to 323.15) K was presented in this paper. On the basis of the obtained data the part of the polythermal solubility surface for the investigated system and the solution density dependence versus the sodium vanadate(V) concentration 1474



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(17) Mazurek, K. Studies on the optimum conditions for leaching the spent vanadium catalyst from metallurgical plants with sodium hydroxide solutions. Przem. Chem. 2012, 91 (2), 234−238. (18) Anioł, S.; Karolewicz, T.; Kubała, J. Recovery of V2O5 from the spent vanadium catalysts. Pol. J. Appl. Chem. 1997, 41, 25−31. (19) Mazurek, K. Recovery of vanadium, potassium and iron from a spent vanadium catalyst by oxalic acid solution leaching, precipitation and ion exchange processes. Hydrometallurgy 2013, 134−135, 26−31. (20) Trypuć, M.; Chałat, M.; Mazurek, K. Solubility investigations in the K2SO4 + V2O5 + H2O system from 293 to 323 K. Pol. J. Appl. Chem. 2003, 4, 287−293. (21) Trypuć, M.; Chałat, M.; Mazurek, K. Solubility investigations in the Na2SO4 + V2O5 + H2O system from 293 to 323 K. J. Chem. Eng. Data 2006, 51 (2), 322−325. (22) Mazurek, K. Investigations on the solubility in the KVO3 + K2SO4 + H2O system from 293.15 to 323.15 K. J. Chem. Eng. Data 2013, 58 (4), 980−985. (23) Brouwer, P. Theory of XRF; PANalytical B.V.: Almelo, Netherlands, 2006. (24) Powder Difraction File; International Center for Diffraction Data: USA, 2001. (25) Przybora, E. Rentgenostrukturalne Metody Identyfikacji Minerałów i Skał (Spectroscopy Methods of Identifying Minerals and Rocks); WG Electronics: Warszawa, 1957. (26) Stephen, H.; Stephen, T. Solubilities of Inorganic and Organic Compounds, Vol. 1. Binary Systems. Macmillian: New York, 1963. (27) Okorafor, O. C. Solubility and density isotherms for the sodium sulfate−water−methanol system. J. Chem. Eng. Data 1999, 44, 488−490. (28) von Plessen, H. Sodium Sulfates. Ullmann’s encyclopedia of industrial chemistry; Verlag Chemie: Weinheim, Germany, 1993. (29) Rosenblatt, D.; Marks, S. B.; Pigford, R. L. Kinetics of phase transitions in the system sodium sulfate−water. Ind. Eng. Chem. Fundam. 1984, 23, 143−147. (30) Linke, W. F. Solubilities of Inorganic and Metal Organic Compounds; American Chemical Society: USA, 1965.

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Investigations on the Solubility, Density, and Viscosity in the NaVO3+

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