Znd. Eng. Chem. Res. 1992, 31, 1222-1227
1222
e = fractional liquid holdup in mixer E = engulfment rate coefficient (8-l) ED = direct energy dissipation rate (W-kg-') k = rate constant (rn3.mol-'.s-') L = length of mixer (m) LL = length of loop reactor (m) LR = length of reaction zone (m) Ne = Newton number nZ = number of recirculations of A-solution during feeding of B-solution o-R, o = ortho monoazo dyestuff (2-[ (4-~ulfophenyl)azo]-lnaphthol) p - R , p = para monoazo dyestuff (4-[(4-sulfophenyl)azo]-lnaphthol) Ap = pressure drop over static mixer (bar) Q = 2-naphthol monoazo dyestuff (1-[(4-sulfophenyl)azo]-2naphthol) Q* = volumetric flow rate ( m 3 d ) S = bisazo dyestuff (2,4-bis[(4-sulfophenyl)azo]-l-naphthol) T = dimensionless time for mixing tF= feed time (8) u = velocity (ms-') VA = volume of naphthol solution (m3) V , = volume of diazotized sulfanilic acid solution (m3) VM = volume of liquid in static mixer (m3) XS' = product yield (new test reaction) XS = product yield (old test reaction) XQ = product yield (new test reaction)
Greek Letters a = volume ratio v A / V B yA1= stoichiometric ratio NA,,/Nb
= turbulent energy dissipation rate (W-kg-') 0 = total energy dissipation rate (W-kg-')
t
Y
= kinematic viscosity ( m 2 d )
Subscripts 0 = initial
s,s=s
Literature Cited Baldyga, J.; Bourne, J. R. Simplification of micromixing calculations, Part 1. Chem. Eng. J. 1989, 42, 83. Baldyga, J.; Bourne, J. R. The effect of micromixing on parallel reactions. Chem. Eng. Sci. 1990, 45, 907. Bourne, J. R.; Tovstiga, G. Micromixing and fast chemical reactions in a turbulent tubular reactor. Chem. Eng. Res. Des. 1988,66,26. Bourne, J . R.; Maire, H. Micromixing and fast chemical reactions in static mixers. Chem. Eng. Process. 1991a, 30, 23. Bourne, J. R.; Maire, H. Influence of the Kinetic Model on Simulating the Micromixing of l-Naphthol and Diazotized Sulfanilic Acid. Ind. Eng. Chem. Res. 1991b, 30, 1285. Bourne, J. R.; Kut, Oe. M.; Lenzner, J. An Improved Reaction System To Investigate Micromixing in High-Intensity Mixers. Znd. Eng. Chem. Res. 1992, 31, 949. Davies, J. T. A physical interpretation of drop sizes in homogenizers and agitated tanks,including the dispersion of viscous oil. Chem. Eng. Sci. 1987, 42, 1671. Godfrey, J. C. Static mixer. In Miring in the process industries; Harnby, N., Edwards, M. F., Nienow, A. W., Eds.; Butterworth: London, 1985. Lenzner, J. Der Einsatz rascher, kompetitiver Reaktionen zur Untersuchung von Mischeinrichtungen. Ph.D. Thesis No. 9469, ETH Ziirich, 1991.
( = stoichimetric ratio NA2,,/NAl0 p
--
lo = reaction A 1 o-R l p = reaction A 1 p-R 20 = reaction p - R S 2p = reaction o-R S 3 = reaction A2 Q A1 = l-naphthol A2 = 2-naphthol B = diazotized sulfanilic acid o = O-R p = p-R R = sum of o-R and p - R
Received for review October 1, 1991 Accepted January 13, 1992
= fluid density (kgm-3)
Calculation of the Thermodynamic Data for Zinc Extraction from Chloride Solutions with Di-n -pentyl Pentanephosphonate Ruey-Shin Juang* and Jiann-Der Jiang Department of Chemical Engineering, Yuan-Ze Institute of Technology, Nei-Li, Taoyuan, 32026,
Taiwan, ROC
The thermodynamic data for the extraction of zinc from chloride solutions with di-n-pentyl pentanephosphonate (DPPP) dissolved in kerosene have been calculated on the basis of the temperature dependence of extraction equilibrium constanta over the temperature range of 20-55 "C.The method, employing either Bromley or the simplified Pitzer equations to estimate the stoichiometric activity coefficient of various species in the aqueous phase, is found to be effective for the evaluation of the thermodynamic data. The extraction reaction is favored by the enthalpy change and unfavored by the entropy change. Introduction In recent years numerous papers concerning the various factors affecting the extraction of a metal chelate have been found. However, only a few thermodynamic studies have been done. It was reported that the thermodynamic data could support the mechanism for synergism of metal extraction with 2-thenoyltrifluoroacetone and neutral phosphorus-based compounds (Nash and Choppin, 1977; *To whom all correspondence should be addressed.
0888-5885/92/2631-1222$03.00/0
Kandil and Ramadan. 1980) or with diDhenvlcarbazone and pyridine (Yamada'et al.,'1982). It isthoight that the thermodynamic data can provide valuable information regarding the extraction mechanism. For the thermodynamic approach of extraction reactions, the enthalpy change (AH) was generally obtained on the basis of the van't Hoff relation from the temperature dependence either of the distribution ratio (Patil et al., 1973; Otu and Westland, 1990) or of the extraction constant (Nash and Choppin, 1977; Kandil and Ramadan, 1980; Kalina et al., 1981; Yamada et al., 1982). However, 0 1992 American Chemical Society
Ind. Eng. Chem. Res., Vol. 31, No. 4, 1992 1223 most of the expressions of distribution ratio or extraction constant contain only the terms of concentration rather than the activity of species in both phases. Recently, Otu and Westland (1990) have employed the extended Debye-Huckel equation to estimate the stoichiometric activity coefficients of various species in the aqueous phase and assumed that the activity coefficients remain unchanged over the temperature range involved. The applicability was, however, questionable, especially in higher ionic strength media. They only gave the relative and even incorrect AH values; thus a great accuracy in the determination of the extraction equilibrium constants would be required. Since 1970 calculations of the activity coefficients of the aqueous electrolytes in high ionic strength media involving the multicomponent system have been improved due to the work done by either Bromley (1973) or Pitzer and his co-workers (Pitzer and Mayorga, 1973; Pitzer and Kim, 1974; Silvester and Pitzer, 1977, 1978). The application of thermodynamics to the extraction equilibria in practical hydrometallurical processes is now possible (Hughes and Hou, 1986; Tanaka, 1990). In this paper, the thermodynamic data for the extraction of zinc from chloride solutions with kerosene solutions of di-n-pentyl pentanephosphonate (DPPP) were calculated by considering the activity corrections of various species in the aqueous phase based on Bromley and simplified Pitzer equations of electrolyte solutions and compared with those obtained by the conventional methods.
Source of Experimental Data The experimental distribution data determined by Nogueira and Cosmen (1983) for the extraction of zinc chloride with DPPP are considered. DPPP was dissolved in kerosene to the required concentration. The aqueous solutions were adjusted using sulfuric acid to a value of pH 1.50-2.00 to avoid hydrolysis of zinc chloride. Nogueira and Cosmen (1983) have reported that the aqueous pH has no effect on the extraction of zinc chloride. In this calculation, the contributions of sulfuric acid to ionic strength and activity coefficienb are ignored for simplicity since it is present at a relatively trace level. For the aqueous phase, the initial concentration of sodium chloride ranged from 0.70 to 2.00 mol dm-3,while of to 1.10 mol dm-3. For zinc chloride was from 7.57 X the organic phase, the initial concentrations of DPPP were from 0.647 to 1.618 mol dm-3. The extraction experiments were performed under various temperatures ranging from 20 to 55 "C. Thermodynamics of the Extraction Reaotion Extraction Equilibrium Relationships. Di-n-pentyl pentanephosphonate (DPPP) is an extraordinarily selective solvating extractant for zinc in chloride solutions. It was reported that DPPP diluted in kerosene has a good phase disengagement rate and its aqueous solubility was determined to be below 4 mg L-l. All these features, together with ita high selectivity for zinc in chloride solutions, make DPPP a promising commercial extractant for zinc chloride (Nogueira et al., 1980). The extraction of zinc from chloride solutions by solvating extractanta such as DPPP was suggested to proceed according to the following solvating reaction (Sato and Nakamura, 1980; Nogueira and Cosmen, 1983): Zn2++ 2C1- + 2DPPP + ZnC12.2DPPP (1) and K,, = [ZnClZ-2DPPP]/ [Zn2+][C1-l2[DPPPlZ (2)
where K,,is the stoichiometricextraction constant and the overbar indicates the species in the organic phase. The distribution ratio of zinc is defined as D = [Kle/[znle (3) where the subscript "ewdenotes the total value at equilibrium. The use of the activity concept only for the aqueous phase has been introduced in order to adapt the model to different aqueous compositions (Tanaka, 1990). The use of activities in the organic phase would require polynomial fitting of the change of activity or molarity ratio of organic species in order to obtain a model for extraction equilibrium (Hughes and Hou, 1986; Tanaka, 1990). Such treatment does not necessarily make the model convenient to use. In addition, Sat0 and Nakamura (1980) have calculated the stability constants of aqueous zinc-chloro complexes by the extraction of zinc chloride with tri-noctylphosphine oxide (TOPO). They presumed that the molar concentration of TOPO may be used since the stability constants obtained with TOPO extraction are essentially similar to those obtained by other methods. It is thought to be justifiable to use the concentration concept for the organic phase formulation. Thus, the activity coefficients of the species in the organic phase are assumed to be kept constant, and eq 2 can be written as K,, = [ZnC12.2DPPP]/ U ~ , ~ + U ~ ~ - ~ [ D P (4) PP]~ where Ke, is the thermodynamic extraction constant. By taking the logarithms of both sides in eq 4, we obtain log [ZnC1,.2DPPP] = log (aZnz+aCl-2)+ 2 log [DPPP] + log K,, (5) It is evident that an adequate method will give a straight line with a slope of 1for the log-log plot of [ZnCl2.2DPPP] [DPPPI2. The value on the left-hand versus (uZnZ+uCl-2) side in eq 5 is given by the experimental data. The second term on the right can be determined by the mass balance: [ m ]= - 2[ZnC12.2DPPP] (6) The value of uznz+ucl-2is calculated by the following four methods. Method names taken are only for convenience. (1) Stoichiometric Method. Most simply, uZnz+and ucl- are replaced by [Zn], and [Cl],, respectively. This treatment is not theoretically correct. The extraction constant obtained here is a conditional constant and will change with the aqueous compositions. (2) Complex Method. Since zinc is strongly complexed by chloride ions, uznz+and ucl- are replaced by their free ion concentrations [Zn2+]and [Cl-I, respectively,while the following complexation reactions are considered: Zn2+ iC1- + ZnCli(2-i), Bi = [ZnCli(2-i)] /[Zn2+][C1-]i (7) where s : 8 (i = 1-4) are the overall stability constants and are taken to be 5.40, 0.80, 0.30, and 0.06 dm3 mol-', respectively, at an ionic strength of unity (Sato and Nakamura, 1980). It is assumed that these constants remain unchanged with temperature (Patil et al., 1973). Thus, [Zn2+]and [Cl-] can be obtained by simultaneously solving the following mass balance equations: [Zn], = [Znz+](l+ Cpi[Cl-li) (8) [Cl], = [Cl], - 2[ZnCl2.2DPPP] = [Cl-] + CiBi[Zn2+][Cl-Ii (9) (3) Bromley Method. The value of uznz+ucl-2is estimated by Bromley equations (Bromley, 1973). The
[m],
+
1224 Ind. Eng. Chem. Res., Vol. 31, No. 4, 1992
equations for electrolyte MX in the mixed electrolyte solution is given by log yMx = -{A,Iz+XIP/'/(~ + P/'))+ (uM/V)CBMAzMA'mA + (YX/V)CBCXzCX'mC (10) Throughout eq 10, the subscript C denotes cation (Na+, Zn2+, or ZnCl+) and A denotes anion (Cl-, ZnCl,-, or ZnClt-). Sums over C and A cover all cations and anions, respectively. Here, BMX,zMX,u, and I are given in eqs 11-14, respectively: Bm= ((0.06 + O.6B&)l%~%xl/(l + ~.~I/I%M%x~)') + Bf& (11) %MX =
v = VM
+ IzXl)/2
(12)
+ VX
(13)
I = 0.5C(mizi2) (14) at 25 "C. BRx is a constant and A, = 0.511 (kg characteristic to the single electrolyte MX and is taken from Bromley (1973). In the actual calculation, we neglect the difference between the molality and molarity scales for simplicity since the average density of aqueous solutions involved here is measured to be 1.032 g/mL. The contributions involving ZnCli(2-i)(i = 1, 3, and 4) to the ionic strength I are considered, but those to the activity coefficients as indicated in eqs 10 and 11 are ignored because of the lack of their Bromley parameters, BRx. It was noted that the experiments were performed ranging from 20 to 55 "C, but all Bromley parameters given above were applicable only at 25 "C. As suggested by Bromley, the Meissner equation (Meissner and Peppas, 1973) could be well used to correlate the temperature dependence of the activity coefficient: log rt = (1.125 - o.oo5t) log rZ5 - (0.125 - o.oo5t) log rref(15) where the temperature, t , is in "C, and log rref= -(0.4iP/~)/(i+ PI2) + 0.03910.~~ (16) log r = (1/1zMzx1)log YMX
(17)
Equation 15 calculates the reduced activity coefficient, r, at temperature t for a given ionic strength from that obtained at 25 "C. Thus, the value of U ~ , . , Z + U C ~can - ~ be calculated by solving eqs 8 and 9 together with the appropriate correlations for the activity coefficient of eqs 10-17 at any temperature. (4) Pitzer Method. The value of U ~ Z + U C is~ -estimated ~ by Pitzer equations (Pitzer and Mayorga, 1973; Pitzer and Kim, 1974; Silvester and Pitzer, 1977, 1978). Here, the mixing coefficients of two ions of the same sign and that of three ions including an ion of the opposite sign are omitted because of the lack of literature data (Tanaka, 1990). The equations for cation M and anion X in the mixed electrolyte solution are expressed as In yM = ZM'F+ CmA(28MA + ZCMA) + IZMlZ)3CmcmACC!A (18) In yx = ZX'F+ Cmc(2Pcx + &x) + I ~ x I C C ~ C ~ A(19) CCA where F = -A6{PI2/(l + bP/') + (2/b) In (1 + bP1')) + CZmcmAPcA (20) z = Cmilzil = 2Cmczc
(21)
CMX
(22)
C&X/(2l~M~X1"')
and A, = 0.391, b = 1.2 at 25 "C. For 1-1, 1-2, and 2-1 electrolytes, Pm and P ' M ~are given by PMX
= && + &!d(x)
(23)
/I
(24)
PMX
= P!$Xf'(x)
where
+ x ) exp(-x))/x2 = -2(1 - (1 + x + x2/2) exp(-x))/x2
f ( x ) = 211 - (1 f'(x)
x =
aI1/2
(25) (26) (27)
and a = 2.0 at 25 "C. In eqs 22-24, figx,/3#x, and ChXare parameters characteristic to the single electrolyte MX and are taken from Pitzer and Mayorga (1973) except those for ZnClz. For the strongly complexing electrolyte ZnClz, they are adapted from Rard and Miller (1989),which are valid up to 1.5 mol kg-*. Similarly, the effect of temperature on the activity Coefficientswas considered. As indicated by Silvester and Pitzer (1977,1978), there was very little effect on the pa&&,and C& for a 10-20-deg change rameters a, b, @gX, in temperature, whereas the DebyeHuckel parameter, A,, was affected to a much greater degree. The correlation between A, and t, ranging from 10 to 70 "C, is given by (Silvester and Pitzer, 1977) A, = 3.770 X lo-' + 4.548 X 10-3t + 4.302 X 10-5t25.977 x 10-8t3 (28) The correlation coefficient for this regression is found to be 0.998. Thermodynamic Relationships. The free energy change of the extraction reaction at any temperature is defined as AG = -2.303RT log K,, (29) where R denotes the universal gas constant (8.314 J mol-' K-l) and T the absolute temperature. The enthalpy change, AH, is usually given by the van't Hoff relation: d(l0g KeJ/d(l/T) = -AH/(2.303R) (30) The entropy change, AS, is thus calculated from the equation A S = (AH - AG)/T (31)
Calculation Results and Discussion As shown in the original source (Nogueira and Cosmen, 1983), the distribution ratio of zinc increases with an increase in the initial chloride concentration for the fixed initial concentrations of DPPP and ZnClz. This confirms that the extraction reaction of eq 1 is reasonable. Also, it is found that the distribution ratio of zinc decreases with increasing temperature. The typical calculation results are shown in Figures 1-4 at 30 "C by the above four methods, respectively. As shown in Figures 1-4, the plotted points by the stoichiometric and complex methods are more scattered, but those by the Bromley and the simplified Pitzer methods lie nearby a straight line of slope 1. Actually, the smallest difference in log K,,is found in the simplified Pitzer method. It is evident that in Pitzer equations the mixing coefficients can almost be negligible as far as this work is concerned. The extraction constants and thermodynamic data calculated at 30 OC are listed in Table I. It is found that
Ind. Eng. Chem. Res., Vol. 31,No. 4,1992 1225 I
,
7
1
I
I
7
, , ,
I
e
/
5
-!
-
-
3
k
lo-',: I-
i:
T-4
4-
2N
J-
a-
Y
I
o
10 -1.k
I
/ I
I
10
/
-=
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a
I I a711
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I
J
4 II?I#
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I
a
1
-2,
I 7 2
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J
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'
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-?
(azn2+ aC1-2)[DPPP12, (mol dm-3)4 Figure 1. Test of the stoichiometricmethod in the extraction of zinc from chloride solutions with DPPP at 30 OC. The slope of the solid line is 1, as indicated in eq 5.
chloride solutions with DPPP at 30 O C . The slope of the solid line is 1, as indicated in eq 5. a
1
I
7-
Figure 3. Test of the Bromley method in the extraction of zinc from
a-
/
I-
-
,
, , , ,,,I
I
7
F6 a d
:I 4-
'1-
8
W
k a a
R
# Y
lo-',: I7 I-
I4-
J-
I-
0
(azn2+ aC1-2-)[DPPP12, (mol d ~ n - ~ ) ~ Figure 2. Test of the complex method in the extraction of zinc'from
chloride solutions with DPPP at 30 "C. The slope of the solid line is 1, aa indicated in eq 5.
Table I. Extraction Constants and Thermodynamic Data Obtained by Various Methods for the Extraction of Zinc from Chloride Solutions with DPPP at 30 OC method 1% K e t stoichiometric 4 6 8 2 f 0.229 complex 0.588 f 0.056 Bromley 1.447 f 0.042 Pitzer 1.619 f 0.030 a
AG,kJ
AH,kJ
mol-'
mol-'
AS,J mol-' K-'
3.96 -3.41 -8.40 -9.40
-18.78 -15.50 -14.38 -14.93
-75.00 -39.88 -19.75 -18.26
Obtained from 110 seta of experimental data.
the scatter in the values of extraction constant, as measured by the standard deviation of the average, is substantially reduced when activity corrections are used. Figure 5 shows the van't Hoff relation by various methods. The obtained equations and the changes in enthalpy (AW and in entropy (AS) over the temperature range of 20-55
Figure 4. Test of the simplified Pitzer method in the extraction of zinc from chloride solutions with DPPP at 30 O C . The slope of the solid line is 1, as indicated in eq 5. Table 11. Equations and the Changes in Enthalpy and in Entropy for the Extraction of Zinc from Chloride Solutions with DPPP over the Temperature Range Studied log K,, =
P + q(l/T) method stoichiometric complex Bromley Pitzer a
AH,kJ
p
u
p
mol-'
-3.93 -2.08 -1.30 -0.95
980.7 809.7 751.4 779.9
0.968 0.973 0.988 0.992
-18.78 -15.50 -14.38 -14.92
AS,J mol-' K-' -75.22 -39.87 -19.73 -18.21
f 0.56 f 0.47
k 0.29 k 0.22
Correlation coefficient.
OC are listed in Table 11. It is evident that AH values obtained by the four methods are not apparently different, but significant differences in the values of AG and AS are observed due to the effect of activity corrections on the extraction constant. Obviously, an accurate determination of the extraction constant plays an important role in
1226 Ind. Eng. Chem. Res., Vol. 31, No. 4, 1992
Council under Grant No. NSC80-0402-E155-04,which is greatly acknowledged.
2
M 0 3
Nomenclature a = activity A , = constant defined in eq 10 (0.511 at 25 "C) t i A, = Debye-Huckel parameter defined in eq 20 (0.392 at 25 0.8 c i "C) E?& = parameter defined in eq 11 for electrolyte MX 0.4 ChX = parameter defined in eq 22 for electrolyte MX D = distribution ratio of zinc AG = free energy change of the extraction reaction, kJ mol-' 0.0 L 1 AH = enthalpy change of the extraction reaction, k J mol-' I = ionic strength, mol kg-' -0.4 K,, = stoichiometric extraction constant, (dm3 K,, = thermodynamic extraction constant log = logarithm of base 10 m = stoichiometric molality, mol kg-I A S = entropy change of the extraction reaction, J mol-' K-' 1 -''s.O 3.1 3.2 3.3 3.4 3.5 t , T = temperature in "C and K, respectively zi = charge of the ion i UT) x lo3, [ ] = molar concentration, mol dm-3
I/ I-
3
1
(~-9
Figure 5. van't Hoff relations for the extraction of zinc from chloride solutions with DPPP by various methods: stoichiometric (O), complex ( O ) , Bromley (A),and Pitzer (W).
calculating the thermodynamic data of a solvent extraction reaction. Actually, the data calculated by Bromley and the simplified Pitzer methods are essentially similar, and the average AH and A S values are found to be -14.65 kJ mol-' and -18.97 J mol-' K-', respectively. This extraction reaction is therefore favored by the enthalpy change and unfavored by the entropy change. The enthalpy change obtained here is a typical value for a solvating reaction (Kalina et al., 1981). The medium, negative entropy term reflects the ordering that must occur as the five molecules and ions on the left-hand side of eq 1condense into a single complex species and also the reordering of the aqueous phase that occurs when the structure-breaking anions are removed. It should also be noted that Nogueira and Cosmen (1983) have calculated the extraction constant and AH value for the extraction of zinc chloride with DPPP to be log K,, = 0.627 and AH = -18.1 kJ mol-', respectively, by the complex method. The small difference compared with our determinations (log K,, = 0.588 and AH = -15.50 kJ mol-') results from directly replacing the free chloride concentration, [Cl-1, in eq 8 by the total chloride concentration at equilibrium, [Cl],, in their study.
Conclusions An accurate determination of the extraction constant plays an important role in calculating the thermodynamic data of a solvent extraction reaction. The AH values calculated from the temperature dependence of the extraction constant by various methods are not apparently different, but significant differences in the values of AG and A S are observed due to the effect of activity corrections when either the Bromley or the simplified Pitzer equations for electrolyte solutions are employed. The average changes in enthalpy and in entropy are found to be AH = -14.65 kJ mol-' and A S = -18.97 J mol-' K-l, respectively, over the temperature range of 20-55 "C. The extraction reaction is thus favored by the enthalpy change and unfavored by the entropy change. Acknowledgment This work was supported by the ROC National Science
Greek Letters a = constant defined in eq 27 &&, /3& = parameters defined in eqs 23 and 24 for electrolyte MX pi = overall stability constants defined in eq 7, dm3 mol-' r = the reduced activity coefficient defined in eq 17 y = stoichiometric activity coefficient vi = number of ion i in one electrolyte molecule Subscripts e = total value at equilibrium 0 = initial value Superscript _ -- species in the organic phase Registry
No. DPPP, 2528-38-3; Zn,7440-66-6.
Literature Cited Bromley, L. A. Thermodynamic Properties of Strong Electrolytes in Aqueous Solutions. AZChE J. 1973, 19 (2), 313-320. Hughes, M. A.; Hou, S. S. Equilibria in the System Cobalt/Di(2ethylhexy1)phosphoricAcid/Water. J. Chem. Eng. Data 1986,31, 4-11.
Kalina, D. G.; Mason, G. W.; Horwitz, E. P. The Thermodynamics of Extraction of Americium(II1) and Europium(II1) from Nitrate Solution by Neutral Phosphorus-Based Organic Compounds. J. Znorg. Nucl. Chem. 1981,43 (3), 579-582. Kandil, A. T.; Ramadan, A. The Synergistic Solvent Extraction of Cobalt(I1) Ion. Radiochim. Acta 1980,27 (4), 229-232. Meissner, H. P.; Peppas, N. A. Activity Coefficients-AqueousSolutions of Polybasic Acids and Their Salts. AZChE J. 1973,19 (4), 806-809.
Nash, K. L.; Choppin, G. R. Thermodynamics of SynergisticSolvent Extraction of Zinc. J. Znorg. Nucl. Chem. 1977, 39, 131-135. Nogueira, E. D.; Coamen, P. The Solvent Extraction of Zinc Chloride with Di-n-pentyl Pentaphosphonate. Hydrometallurgy 1983,9, 333-347.
Nogueira, E. D.; Regife, J. M.; Blythe, P. M. Zincex-The Development of a Secondary Zinc Process. Chem. Znd. 1980 (Jan 191, 63-67.
Otu, E. 0.; Westland, A. D. The Thermodynamics of Extraction of Some Lanthanide and Other Ions by Dinonylnaphthalenesulfonic Acid. Solvent Extr. Zon Exch. 1990,8 (6), 827-842. Patil, S.K.; Ramakrishna, V. V.; Avadhany, G. V. N.; Ramaniah, M. V. Some Studies on the TBP Extraction of Actinides. J. Znorg. Nucl. Chem. 1973, 35 (7), 2537-2545. Pitzer, K. S.; Mayorga, G. Thermodynamics of Electrolytes. 11. Activity and Osmotic Coefficients for Strong Electrolytes with One or Both Ions Univalent. J. Phys. Chem. 1973,77,2300-2308. Pitzer, K. S.; Kim, J. J. Thermodynamics of Electrolytes. IV. Activity and Osmotic Coefficients for Mixed Electrolytes. J. Am. Chem. SOC.1974, 96, 5701-5707.
1227
Ind. Eng. Chem. Res. 1992,31,1227-1231 Rard, J. A.; Miller, D. G. Isopiestic Determination of the Osmotic and Activity Coefficients of ZnClz at 298.15 K. J. Chem. Thermodyn. 1989,21,463-482. Sato, T.; Nakamura, T. The Stability Constants of the Aqueous Chloro-Complexesof Divalent Zinc, Cadmium and Mercury Determined by Solvent Extraction with Tri-n-octylphosphineOxide. Hydrometallurgy 1980,6, 3-12. Silvester, L.F.; Pitzer, K. S. Thermodynamics of Electrolytes. VIII. High Temperature Properties, Including Enthalpy and Heat Capacity, with Application to Sodium Chloride. J. Phys. Chen. 1977,81, 1822-1828. Silvester, L. F.; Pitzer, K. S. Thermodynamics of Electrolytes. X.
Enthalpy and the Effect of Temperature on the Activity Coefficients. J. Solution Chem. 1978, 7 (5), 327-337. Tanaka, M. Modelling of Solvent Extraction Equilibria of Copper(I1) from Nitric and Hydrochloric Acid Solutions with &Hydroxyoxime. Hydrometallurgy 1990,24, 317-331. Yamada, E.; Nakayama, E.; Kuwamoto, T.; Fujinaga, T. A. Thermodynamic Study of the Synergistic Solvent Extraction of a Series of Zinc(I1) and Cadmium(I1) Complexes. Bull. Chem. SOC. Jpn. 1982,55 (lo), 3155-3159. Received for review October 10, 1991 Accepted December 19, 1991
Solvent Effects on the Hydration of Cyclohexene Catalyzed by a Strong Acid Ion Exchange Resin. 1. Solubility of Cyclohexene in Aqueous Sulfolane Mixtures Henk-Jan Panneman and Antonie A. C. M. Beenackers* Department of Chemical Engineering, State University of Groningen, Nijenborgh 4, 9747 AG Groningen, The Netherlands
The solubility of cyclohexene in different water-sulfolane mixtures was measured between 313 and 413 K. The results demonstrate a sharp increase of the solubility of cyclohexene with increasing percentages of sulfolane in the solvent mixture. Without sulfolane the increase of the solubility with temperature is higher than in mixtures with sulfolane. From thermodynamic calculations, using UNIFAC to predict the activity coefficients of all components, the solubility of cyclohexene in various mixtures and a t different temperatures could be computed. The experimental and computed solubilities are of comparable magnitude and show the same temperature and mixture-composition dependence. The solubility of cyclohexene in a mixture of 90 mol 9% sulfolane and 10 mol 9% water is, depending on the temperature, between 100 and 200 times higher than in pure water.
Introduction The direct hydration of alkenes catalyzed by a macroporous strong acid ion exchange resin is a simple and cheap process for producing alcohols. A disadvantage is the low mutual solubility of alkenes and water. A possible alternative is the use of an organic cosolvent, which gives a single liquid phase and a greatly improved solubility of alkenes in the solvent mixture. For the conditions that have to be satisfied for a proper cosolvent see part 3 (Panneman and Beenackers, 1992). Addition of a cosolvent will not only increase the solubility of the alkene in the reaction mixture but also change the equilibrium conversion, the reaction rate constant, and the activity of the ion exchange resin. We have studied the hydration of cyclohexene and used sulfolane as a cosolvent. In this contribution, we present both experimentally measured and theoretically predicted solubilities of cyclohexene in solvent mixtures of water and sulfolane at various temperatures. Such data are necessary for designing a process for the synthesis of cyclohexanol from cyclohexene with sulfolane as a cosolvent. The solubility of cyclohexene in water at room temperature is 0.003 km01.m-~(McAuliffe, 1966). As far as the authors know, no literature data are available for the solubility of cyclohexene in water-sulfolane mixtures. Mixtures of cyclohexene with water or water-sulfolane are far from ideal, and as a consequence predicted solubilities, which can be obtained from thermodynamic models, are highly uncertain. In this paper experimentally obtained solubilities of cyclohexene are compared with predicted solubilities, calculated with UNIFAC. Experimental Section A dynamic method was applied to measure the solubility of cyclohexene as a function of solvent composition. A
schematic representation of the equipment used is given in Figure 1. The feed liquids, cyclohexene and a watersulfolane mixture, were stored in two displacement pumps (Model 314 ISCO pumps) with a maximum flow rate of 5.56 X m 3 d (200 mL-h-'). The flow rate could be adjusted a t any level below the maximum value. During an experiment, the two liquid feed streams were continuously fed to a static mixer (i.d. = 3 X m; L = 4 m; filled with glass beads, d = 0.25 X m) immersed in a thermostatic oil batk. As shown in the Appendix, conditions in the mixer were such that at the mixer outlet phase equilibrium was established. The two-phase stream leaving that mixer was fed to a separator, also immersed in the oil bath. This separator could be considered as a pair of communicating vessels, interconnected both at the top and the bottom. The two-phase stream entered the first vessel in the middle. The light and heavy phases passed through the top and bottom connections, respectively. Here, continuous sample streams could be withdrawn from either the light or the heavy phase and directed to a liquid injection valve of a gas-liquid chromatograph. The sample tubes were electrically heated to prevent phase separation. A needle valve behind the injection valve kept the sample stream at pressure. The analysis of the samples is described elsewhere (Panneman and Beenackers, 1992; Marsman et al., 1988). The two-phase stream from the second vessel of the separator left the system via a back-pressure valve, which kept the total pressure in the system at 2 MPa.
Results The solubility of cyclohexene has been measured in various mixtures of sulfolane and water at temperatures between 313 and 413 K and a pressure of 2 MPa. The water-sulfolane ratio is influenced by the amount of sul-
088S-5885/92/2631-1227$03.QO/O0 1992 American Chemical Society
