Solubility of Na2CO3 and NaHCO3 in Aqueous Sodium Sulfate


Apr 27, 2015 - To separate a salt mixture (mainly consisting of Na2CO3 and Na2SO4) which is formed from the 4,4′-diaminostilbene-2,2′-disulfonic a...
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Solubility of Na2CO3 and NaHCO3 in Aqueous Sodium Sulfate Solutions and Its Application to Separating Na2CO3 and Na2SO4 Salt Mixtures Wen-Yu Zhu, Yu Gu, Lin Zhang, Hui Jing, Bo Liu, Feng-Bao Zhang, Guo-Liang Zhang, and Qing Xia* School of Chemical Engineering and Technology, Tianjin University, Tianjin, 300072, PR China S Supporting Information *

ABSTRACT: To separate a salt mixture (mainly consisting of Na2CO3 and Na2SO4) which is formed from the 4,4′diaminostilbene-2,2′-disulfonic acid wastewater treatment, the solubility data of Na2CO3 and NaHCO3 in Na2SO4−water solvent mixtures were measured using a dynamic method over the 275−335 K temperature range. The solute-free mass fraction of Na2SO4 in the solvent mixtures extended from 0 to 0.25. The solubility of Na2CO3 and NaHCO3 in Na2SO4−water solvent mixtures decreased with addition of Na2SO4. With rising temperature, the solubility of NaHCO3 increased while the solubility of Na2CO3 increased at first and then decreased after transition points. The experimental data were correlated with the electrolyte nonrandom two-liquid model. The root-mean-square deviations of solubility temperature varied from 0.35 to 1.50 K. Based on the solubility data of this work, a new strategy for separating the Na2CO3 and Na2SO4 mixture was proposed theoretically.

1. INTRODUCTION As an important intermediate that is widely used in the synthesis of direct dyes, mothproof agents, and fluorescent brighteners,1,2 4,4′-diaminostilbene-2,2′-disulfonic acid (DSD acid) is synthesized from 4-nitrotoluene via three main steps, namely, sulfonation, oxidation, and reduction.3 As a consequence of the low reaction efficiency of the oxidation process,4 the wastewater produced in the oxidation process contains high concentrations of various substituted derivatives of aromatic compounds as well as inorganic substances (mainly Na2SO4)5 and cannot be discharged directly. Over the past decades, several methods have been investigated for treating this wastewater and recovering the valuable resources, including oxidation with Fenton’s reagent, complexometric extraction, and resin absorption.2,6−9 However, some deficiencies restrict their wide application in industry. Fenton’s reagent is an efficient oxidizing agent to degrade the organic substances, but it has a high cost and is better used in dilute solution. The extraction method results in a secondary contamination as well as an incrase in operating cost because the extracting agent may run off to the aqueous phase. Resin absorption produces mixtures which require further processing to recycle the adsorbate. Evaporation−incineration,3 which includes the multipleeffect evaporation process combined with incineration, has been applied in industry to treat DSD acid wastewater. The multiple-effect evaporation process is used mainly to recycle the volatile organic compounds. Nevertheless, some organic compounds remain in the treated water after the multipleeffect evaporation. After spray drying, the residual organic compounds are converted to Na2CO3 by incineration. With all the procedures, a salt mixture (mainly consisting of Na2SO4 and Na2CO3) is formed. According to the data from industrial practice, every 100 g of salt mixture contains about 82 g of Na2SO4 and 18 g of Na2CO3. Currently, the salt mixture is sold at a very low price; © XXXX American Chemical Society

however, if the mixture is able to be separated efficiently, both Na2SO4 and Na2CO3 can be reused in the DSD acid process to reduce operating cost. As reported by Lupeiko et al.,10 the solubility of NaHCO3 in water is less than that of both Na2SO4 and Na2CO3. Therefore, a speculated method to separate Na2SO4 and Na2CO3 is formed. First, Na2CO3 is converted to NaHCO3, which is later separated from Na2SO4 by solubility difference. Finally, NaHCO3 is calcined and Na2CO3 is obtained as the final product. To design and optimize the above process, the solubility data of Na2CO3 and NaHCO3 in the Na2SO4−water mixed solutions were measured by the dynamic method over the temperature range from 275 to 335 K. The experimental data were regressed with the E-NRTL model. According to the determined data, a new strategy for separating the salt mixture (Na2SO4 and Na2CO3) is presented and discussed.

2. EXPERIMENTAL SECTION 2.1. Material Preparation. Prior to the solubility measurements, Na2SO4 and Na2CO3 were dried in an oven at 343 K for 2 days until constant weight, while NaHCO3 was dried in a vacuum oven at 318 K. (Its decomposition temperature is 323 K.11) The mass fraction purities of Na2SO4, Na2CO3, and NaHCO3 (purchased from Guangfu Chemical Reagents Co., Tianjin, China) were 0.990, 0.998, and 0.995, respectively. Deionized water was obtained from the Nankai Chemical Reagents Co., Tianjin, China. These reagents were used without further purification. 2.2. Apparatus and Procedure. The solubility was determined using the dynamic method, which was described Received: January 28, 2015 Revised: April 23, 2015 Accepted: April 27, 2015

A

DOI: 10.1021/acs.iecr.5b00381 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

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Industrial & Engineering Chemistry Research in detail previously.12 The laser monitoring technique was used to determine the solid−liquid equilibrium (SLE) temperature of a certain composition mixture by observing the dissolving processes. Predetermined solvents and solute weighed by analytical balance (Gibertini, Crystal 200, Italy; accuracy of 0.0001 g) were put into a jacketed glass vessel. The mixture was heated slowly with continuous stirring. The heating rate was decreased when the solid in the solution decreased. In particular, when there is very little solid left in the solution, the temperature is kept for a few hours for every 0.1 K increase. Heating rate must be very slow in the final stage in order to establish SLE. The temperature of the mixture, which was measured by a platinum resistance thermometer Pt-100 (calibrated with an accuracy of 0.01 K), was controlled by a refrigerated−heating circulator (Julabo FP45-HE, Germany; temperature stability ±0.01 K). The temperature at which the last crystal of the salt disappeared and the intensity of the laser beam reached a maximum value was taken as the SLE temperature. The uncertainty of the measured temperature was ±0.3 K.

decline. To confirm the forms of Na2CO3 salts which precipitated at the temperature lower and higher than the transition point, saturated Na2CO3 aqueous solutions were prepared at 300 and 310 K, respectively. The solution (310 K) is heated to 330 K to obtain crystal sample 1, and the solution (300 K) is cooled to 283 K to get crystal sample 2. The samples are analyzed by thermogravimetric analysis (TGA) and X-ray diffraction (XRD). Figure S1 in Supporting Information shows the XRD pattern of sample 1. All peaks can be readily ascribed to Na2CO3·H2O, and no impurity is detected. TGA results of sample 1 gives 14.9% weight loss of hydrate, which is nearly the same as the theoretical value (the weight of H2O in Na2CO3· H2O molecule is 14.5%). It is proved that the SLE is between Na2CO3·H2O and water at a temperature higher than the transition point. Na2CO3·10H2O is unstable and easily losses hydrate when it is exposed to the air. We tried to analyze sample 2 by XRD, but we were not successful. The result in TGA also supports the instability of Na 2CO3 ·10H2O; inconsistent results were obtained in the twice repeated tests. Wells and McAdam13 reported the phase relations of Na2CO3 and water. Three hydrates of Na2CO3 exist in the establishment of SLE between Na2CO3 and water. They are decahydrate, heptahydrate, and monohydrate, respectively. Considering the results of Wells and McAdam, the SLE is established between sodium carbonate decahydrate (Na2CO3·10H2O) and water at a temperature lower than the transition point, while it becomes sodium carbonate monohydrate (Na2CO3·H2O) at a temperature higher than the transition point. The transition point is known to be 306.1 K in pure water in the work of Wells and McAdam, and in this case it is measured as 308.8 K. 3.1.2. Solubility Data of NaHCO3 in Na2SO4−Water Solvent Mixtures. The solubility data investigated of NaHCO3 in Na2SO4−water solution are given in Table S2 in Supporting Information, where Texp is the measured SLE absolute temperature and x2 is the solubility of NaHCO3 in unit of mole fraction; w03 is the solute-free mass fraction of Na2SO4 in solvent mixtures. Figure 2 shows the solubility of NaHCO3 as a function of temperature in the solvent mixtures under different concentration. Each curve indicates that the solubilities of NaHCO3 increase with temperature and decrease

3. RESULTS AND DISCUSSION 3.1. Experimentally Determined Solubility Data. 3.1.1. Solubility Data of Na2CO3 in Na2SO4−Water Solvent Mixtures. The experimentally determined solubility data of Na2CO3 in Na2SO4−water solutions are listed in Table S1 in Supporting Information, where Texp is the measured SLE absolute temperature and x1 expresses the mole fraction solubility of anhydrous sodium carbonate; w03 is the solutefree mass fraction of Na2SO4 in solvent mixtures. Figure 1 is the

Figure 1. Mole fraction solubility data of Na2CO3 (x1) at different compositions of (Na2SO4−water) solvent mixtures. Solute-free mass fraction of Na2SO4 w03: ■, 0 (pure water); ●, 0.05; ▲, 0.10; ▼, 0.15; ★, 0.20; ◆, 0.25; △, □, and ○ indicate literature data in pure water;13−15 the black traces represent solubility curves calculated from the E-NRTL model.

curves of x1 versus temperature at different w03. As shown in Figure 1, the solubility of Na2CO3 first increases with temperature and decreases afterward. For each concentration of the solvent mixtures, the increase is sharper than the decrease. The solubility of Na2CO3 decreases with addition of Na2SO4. There is nearly no evident solubility difference for these five curves at low temperature, whereas the solubility difference becomes obvious after the transition point occurs. With additional Na2SO4, the solubility demonstrates a larger

Figure 2. Mole fraction solubility data of NaHCO3 (x2) at different compositions of (Na2SO4−water) solvent mixtures. Solute-free mass fraction of Na2SO4 w03: ■, 0 (pure water); ●, 0.05; ▲, 0.10; ▼, 0.15; ★, 0.20; ◆, 0.25; □, ○, and △ are literature data in pure water;10,14,15 the black traces indicate solubility curves calculated from the E-NRTL model. B

DOI: 10.1021/acs.iecr.5b00381 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research

Figure 3. Simple process for separation of Na2CO3 and Na2SO4.

range ion−ion interaction (γi*PDH), and the other from the short-range interaction (γi*lc). They are represented by the Pitzer−Debye−Hückel model and the nonrandom two-liquid theory (NRTL), respectively. The E-NRTL model can be written as the following expression:

with addition of Na2SO4. Therefore, Figure 2 indicates that the lowest dissolving amount conditions in Na2SO4−water solvent mixtures is w03 = 0.25 and Texp = 301 K. Saturated NaHCO3 aqueous solution which is prepared at 303 K is cooled to 283 K to obtain a crystalline solid. The XRD pattern of the solid is shown in Figure S2 in Supporting Information. All peaks in the XRD pattern can be readily ascribed to NaHCO3, and no impurity was detected, indicating the form of precipitate is NaHCO3. The SLE is established between NaHCO3 and water. 3.2. Comparison with Literature Data. The solubilities of Na2CO3 and NaHCO3 in water are compared with the published data. Figures 1 and 2 show the comparison results. It can be seen clearly that the experimentally determined solubility data is basically the same as the literature data.10,13−15 3.3. Correlation of Experimental Data with the E-NRTL Model. The solubility product represents the phase equilibrium between the salt’s crystal phase and the solvent phase, and it can be used to describe the SLE in electrolyte systems. The temperature dependence of the solubility product constant (Ks) defined by the activity coefficient and molarity of aqueous ions is expressed as follows: K s = a+ν +a−ν − = (γ+x+)ν + (γ−x−)ν −

ln γi* = ln γi*PDH + ln γi*lc

The E-NRTL model parameters include the energy interaction parameter τij, which is expressed as a function of temperature, and the nonrandomness factor α. The format of τij is written as follows:17

τij = aij + bij /T

(4)

where aij and bij are constants representing the temperature dependence of τij. The results are listed in Tables S4 and S5 in Supporting Information. In addition, in regressing data of a single-solvent electrolyte system, a value of 0.2 is satisfied for the solvent−salt nonrandomness factor α in practice. In this work, the water−salt nonrandomness factor is fixed as 0.2 in the optimization process. The Nelder−Mead simplex method18 along with Matlab is used to determine the E-NRTL model parameters. The object function is the root-mean-square deviation, σ, which is defined as

(1)

where Ks, a, ν, γ, and x denote the solubility product constant, ion activity, electrolyte stoichiometric coefficient, ion asymmetric activity coefficient, and dissociated ion in units of mole fraction, respectively. The subscripts (+) and (−) refer to cation and anion, respectively. Equation 1 can be used for aqueous solutions, and it can also be applied not only to singlecomponent systems but also to multicomponent systems.16 Furthermore, in pure solvent m, the solubility product constant (Ksm) is a function of temperature and can be written as follows: ln K sm = A m + Bm /T

(3)

N

σ = [∑ (T exp − T )2 /(N − 1)]0.5 i=1

(5)

where σ denotes the root-mean-square deviation between Texp and T; Texp is the experimental equilibrium temperature, T the calculated equilibrium temperature, and N the number of experimental data points. The root-mean-square deviation results are listed in Table S6 in Supporting Information. The calculation results are also shown in Figure 1 and Figure 2; good agreement is achieved between the experimental equilibrium temperature acquired in this work and the equilibrium temperature calculated by the E-NRTL model for the Na2CO3/NaHCO3−Na2SO4−water systems. In Table S6 in Supporting Information, the root-mean-square deviations vary from 0.35 to 1.50 K, which indicates that the solubility curves can be well expressed by the E-NRTL model. 3.4. Theoretical Application of the Determined Solubility Data. As seen in Figure 2, the lowest dissolving capacity conditions of NaHCO3 is w03 = 0.25 and Texp = 301 K. It is known that Na2SO4 has a maximum saturation concentration of 0.32 mass fraction in water.19 When the concentration of aqueous Na2SO4 solution (w03) exceeds 0.25, Na2SO4 may precipitate out with NaHCO3. As a result, we choose w03 = 0.25 and Texp = 301 K as the conditions of

(2)

where T is absolute temperature in kelvin; Am and Bm are constants given in Table S3 in Supporting Information, which are obtained by regressing solubility data; and m represent the solvent, water. To calculate the SLE temperature of the dissolution of a solute in a solvent, a thermodynamic model which allows estimating the activity coefficients of all the components considered in solution is needed. The E-NRTL model is used in this work. Applied in the electrolyte system, the E-NRTL model contains two basic assumptions about the lattice structure: (1) local electroneutrality assumption and (2) like-ion repulsion assumption.17 The activity coefficient of ion species i (γi*) is a combination of two contributions, one resulting from the longC

DOI: 10.1021/acs.iecr.5b00381 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Industrial & Engineering Chemistry Research



separating NaHCO3 from Na2SO4−water solvent mixtures. According to the solubility data, every 100 g of 0.25 mass fraction aqueous Na2SO4 solution can dissolve 2.07 g of NaHCO3 or 12.85 g of Na2CO3 under these conditions. The proposed procedure is shown in Figure 3. About 246 g of water is added to 100 g of salt mixture to form w03 = 0.25 aqueous Na2SO4 solution under the temperature of 301 K. Meanwhile, all 18 g of Na2CO3 in the mixture is dissolved according to the determined data. About 7.47 g of carbon dioxide is introduced into the solution to convert 18 g of Na2CO3 to 28.53 g of NaHCO3, theoretically. After the reaction, 21.74 g of NaHCO3 will precipitate out according to the determined solubility data. The cake and the filtrate are separated by filtration at 301 K. The cake, which contains NaHCO3, is calcined, and 13.72 g of Na2CO3 can be obtained. About 3.96 g of sulfuric acid is added to the filtrate, which contains Na2SO4, residual NaHCO3, and water. The residual NaHCO3 is converted to Na2SO4. As a result, 87.74 g of Na2SO4 can be obtained from the filtrate after evaporation and drying processes. Throughout the entire treating process, based on each 100 g of salt mixture, 249 g of water and 3.96 g of sulfuric acid are consumed; 7.47 g of CO2 is consumed only in the initial conditions, and 9.26 g of CO2 which is produced in the subsequent procedures can be recycled. In total, 87.74 g of pure Na2SO4 and 13.72 g of pure Na2CO3 can be obtained, which can be reused in the DSD acid production process.

AUTHOR INFORMATION

Corresponding Author

*Tel.: +86-22-27400292. Fax: +86-22-27408778. E-mail: [email protected] Notes

The authors declare no competing financial interest.



ACKNOWLEDGMENTS The authors acknowledge the support provided by the Program of Introducing Talents of Discipline to Universities, China, B06006.



REFERENCES

(1) Liu, Y.; Zhang, F. B.; Zhang, G. L. Kinetic study on the preparation of 4,4′-dinitrostilbene-2,2′-disulfonic acid (I)Kinetic study on the oxidation of 4,4′-dinitrobibenzyl-2,2′-disulfonic acid to prepare 4,4′-dinitrostilbene-2,2′-disulfonic acid. Dyes Pigm. 2003, 56, 181. (2) Rajagopal, R.; Seshadri, S. Light-stabilized azo dyes containing a built-in ultraviolet absorber residue. Dyes Pigm. 1988, 9, 233. (3) Li, G. Z.; Zhang, F. B.; Zhang, G. L.; Han, J. Recovery of Na2SO4 from remediation of wastewater and reuse for preparation of sodium 4nitrotoluene-2-sulfonate (NTSNa). Desalination 2006, 194, 176. (4) Peng, W. C.; Chen, Y.; Fan, S. D.; Zhang, F. B.; Zhang, G. L.; Fan, X. B. Use of 4,4′-dinitrostilbene-2,2′-disulfonic acid wastewater as a raw material for paramycin production. Environ. Sci. Technol. 2010, 44, 9157. (5) Yu, G.; Zhu, W. P.; Yang, Z. H.; Jiang, Z. P. Physicochemical decolorization technology of dye wastewater. Environ. Sci. (Beijing, China) 1994, 15, 75. (6) Yu, G.; Zhu, W. P.; Yang, Z. H. Pretreatment and biodegradability enhancement of DSD acid manufacturing wastewater. Chemosphere 1998, 37, 487. (7) Chai, L. M.; Zhang, F. B.; Zhang, G. L. Treatment of DSD acid wastewater using a weak basic resin. Desalination 2005, 180, 157. (8) Rachel, A.; Lavedrine, B.; Aguer, J. P.; Boule, P. Photochemical study of 4,4′-dinitrostilbene-2,2′-disulfonate (DSD) degradation in water. J. Photochem. Photobiol., A 2002, 151, 137. (9) Zhu, W. P.; Lang, Z. H.; Wang, L. Application of ferrous hydrogen peroxide for treatment of DSD-acid manufacturing process wastewater. Water. Res. 2001, 35, 2087. (10) Lupeiko, T. G.; Ivleva, T. I.; Ivanova, O. K. System chloridesodium bicarbonate water-triethylamine. Russ. J. Inorg. Chem. 1989, 34, 1047. (11) Pasquali, I.; Bettini, R.; Giordano, F. Thermal behaviour of diclofenac, diclofenac sodium and sodium bicarbonate compositions. J. Therm. Anal. Calorim. 2007, 90, 903. (12) Xia, Q.; Zhang, F. B.; Zhang, G. L.; Ma, J. C.; Zhao, L. Solubility of sebacic acid in binary water + ethanol solvent mixtures. J. Chem. Eng. Data 2008, 53, 838. (13) Wells, R. C.; McAdam, D. J. Phase relations of the system: sodium carbonate and water. J. Am. Chem. Soc. 1907, 29, 721. (14) Perry, R. H.; Green, D. W. Perry’s Chemical Engineers’ Handbook, 7th ed.; McGraw-Hill: New York, 1997. (15) Speight, J. G. Lange’s Handbook of Chemistry; McGraw-Hill: New York, 2005. (16) Li, M. L.; Constantinescu, D.; Wang, L. S.; Mohs, A.; Gmehling, J. Solubilities of NaCl, KCl, LiCl, and LiBr in methanol, ethanol, acetone, and mixed solvents and correlation using the LIQUAC model. Ind. Eng. Chem. Res. 2010, 49, 4981. (17) Aspen Technology. Aspen Physical Property System 11.1, Cambridge. (18) Nelder, J. A.; Mead, R. A. A simplex method for function minimization. Comput. J. 1965, 7, 308. (19) Zhilova, S. B.; Karov, Z. G.; El’mesova, R. M. Solubility isotherm for the Na2MoO4-Na2SO4-H2O. Russ. J. Inorg. Chem. 2008, 53, 628.

4. CONCLUSIONS At a temperature range of 275 to 335 K, the solubility data of Na2CO3 and NaHCO3 in the binary solvent mixtures Na2SO4− water have been investigated by the dynamic method. When the temperature increases, the solubilities of NaHCO3 increase while the solubilities of Na2CO3 increase first then decrease afterward. The solubilities of both solutes decrease rapidly with addition of Na2SO4, except the solubility of Na2CO3 at low temperature. The highest solubilities of both are observed in pure water, and the lowest solubility occurs in 0.25 mass fraction Na2SO4 aqueous solution. The solubilities of NaHCO3 in aqueous Na2SO4 solutions are much smaller than those of Na2CO3 in the same solvent composition. The E-NRTL model is used in correlations of the experimental solubility. The results are acceptable in the systems studied because the root-mean-square deviations for each system range from 0.35 to 1.50 K. The presented strategy leads to a good separation result. Theoretically, 87.74 g of pure Na2SO4 and 13.72 g of pure Na2CO3 can be obtained by separating 100 g of salt mixture.



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ASSOCIATED CONTENT

S Supporting Information *

Tables of the experimental solubilities of Na2CO3 and NaHCO3 at Texp and solute-free mass fraction of Na2SO4 in binary Na2SO4−water solvent mixtures; the parameters of Ksm; E-NRTL model parameters and root-mean-square deviations σ of Na2CO3−Na2SO4−H2O system and NaHCO3−Na2SO4− H2O system. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/ acs.iecr.5b00381. D

DOI: 10.1021/acs.iecr.5b00381 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX