1294
Znd. Eng. Chem. Res. 1995,34,1294-1301
Distribution Equilibrium of Citric Acid between Aqueous Solutions and Tri-n-octylamine-Impregnated Macroporous Resins Ruey-Shin Juang* and Hwai-Luh Chang Department of Chemical Engineering, Yuan-Ze Institute of Technology, Nei-Li, Taoyuan, Taiwan 32026,ROC
In this paper, experiments were conducted on equilibrium distribution of citric acid between aqueous solutions and tri-n-octylamine (T0A)-impregnated resins in the 288-318 K temperature range. Analysis of the results showed t h a t the sorption of citric acid could be explained by the formation of acid-TOA complexes in the resin phase with the general composition (H&,(TOA),. Mole fraction of the complexes formed in the resin phase as a function of citric acid concentration in the aqueous solution was obtained. The effect of temperature on equilibrium distribution was also studied, and the apparent thermodynamic data for sorption reactions were determined. Finally, the distribution of citric acid between aqueous solutions and the resins impregnating with a mixture of TOA and bis(2-ethylhexy1)phosphoric acid was examined.
Introduction Due to the difficulty and low yield of acid crystallization, as well as the extra consumption of chemicals and the damage to the environment, amine extraction has been found to be a promising alternative to the conventional precipitation process for the recovery of organic acids from aqueous solutions (Baniel, 1982; Baniel and Gonen, 1991; Hartl and Marr, 1993; Kerest and King, 1986; King, 1992; Wennersten, 1983). For example, many works have been done on the extraction equilibrium of citric acid with tertiary amines like tri-noctylamine (TOA), trilaurylamine, and Alamine 336 (Bizek et al., 1992a,b, 1993; Juang and Huang, 1994; Sat0 et al., 1985; Vanura and Kuca, 1976). However, the solvent extraction process requires violent mixing of the phases to provide sufficient contact area for a satisfactory rate of extraction followed by gravity settling of the mixed phases (Tavlaride et al., 1987). It needs the mixing and settling apparatus and must overcome the problems of reagent loss through entrainment and those of phase separation due to the formation of the third phase and emulsion. In addition, this process is generally considered to be economical for aqueous feed in the metal and carboxylic acid concentration range of 0.01-1.0 and 0.1-2.0 mol/dm3,respectively (Akita and Takeuchi, 1990; Hartl and Marr, 1993). The use of synthetic neutral macroporous resins as polymeric adsorbents offers another choice for the recovery of carboxylic acids from dilute aqueous solutions because the resins have the necessary physical and chemical strength, the high surface aredvolume ratio, and the easy regeneration nature by appropriate organic solvents (Hasanain and Hine, 1981; Kulprathipanaja, 1988; Thurman et al., 1978). Nevertheless, the adsorption capacity of the acids for these resins is decreased as the hydrophility of the acid is increased (Hasanain and Hine, 1981; Thurman et al., 1978). For example, the adsorption of citric acid is measurable only when the acid concentration is relatively high (Kulprathipanaja, 1988). It has been recognized that extractant-impregnated resin (EIR) is a novel separation medium. This is due to the fact that it can bridge the gap between solvent
* To whom correspondence should be addressed.
extraction and resin ion-exchange (or adsorption) processes. It combines not only the advantages of solid ion exchange for processing very dilute liquors with specific properties of the extractants but also the high distribution and selectivity distinctive to the extractants dissolved in a liquid organic phase with the simplicity of equipment and operation distinctive t o solid ionexchange technology (Tavlaride et al., 1987;Warshawsky, 1981). Accordingly, the EIR is very suitable for scavenging a specific solute with high selectivity over other species. A large amount of work on the sorption and separation of metals with EIR has been carried out (Akita and Takeuchi, 1990; Akita et al., 1993; Cortina et al., 1992, 1993a,b, 1994a,b; Juang and Su, 1992a,b). The impregnated extractants mainly include organophosphorus acids (Cortina et al., 1992, 1993a,b, 1994a,b; Juang and Su, 1992a,b)and TOA (Akita and Takeuchi, 1990; Akita et al., 1993). It has been reported that the extractants would exhibit strong affinity to the polymeric matrix and could behave as in the liquid state (Warshawsky, 1981; Cortina et al., 1993a, 1994a). In addition, the possibility of the application of EIR for the recovery and separation of metals has been already justified. To our knowledge, little attention has been paid to examining the sorption of organic solutes (Gao and Su, 1991). In the present study, the sorption of citric acid from aqueous solutions with TOA-EIR was investigated as a function of citric acid concentration in the aqueous phase, TOA concentration in the EIR phase, and temperature. The operating lifetime of TOA-EIR was also examined. The effect of a water-immiscible organic acid, bis(2-ethylhexy1)phosphoricacid (BPEHPA),on the equilibrium distribution of citric acid with TOA-EIR was finally studied.
Experimental Section Reagents and Macroporous Resins. TOA and BBEHPA were the products of Tokyo Chemical Industry Co., Ltd., Japan, and Merck Co., respectively. They both had a purity of about 98.5% and were used without further purification. Citric acid, n-hexane, and the other inorganic chemicals were also supplied by Merck Co. as analytical reagent grade, and all were used as received. Deionized water produced by a Millipore Milli-Q water system was used throughout the work.
0888-5885/95/2634-1294$09.00/00 1995 American Chemical Society
Ind. Eng. Chem. Res., Vol. 34, No. 4,1995 1295 Amberlite XAD-2 and XAD-4 macroporous resins, supplied by Merck Co., were made of styrene-divinylbenzene copolymer with highly aromatic structure. On a dry basis, they had a specific surface area of 330350 and 750-780 mz/g, a porosity of 0.42 and 0.51, and an average pore diameter of 9 and 5 nm, respectively (Komiyama and Smith, 1974). The particle size of both resins was 0.3-0.9 mm (20-50 mesh). These resins were washed by acetone and n-hexane and dried a t 323 K in a vacuum for 2 h before impregnation. Preparation of the TOA-EIR and Solutions. The TOA-impregnatedAmberlite XAD resins were prepared by the following dry procedure (Akita and Takeuchi, 1990; Juang and Su, 1992a). An aliquot of TOA (0.10.5 g) was first dissolved into a precalculated amount of n-hexane (3 cm3). The resulting n-hexane solution was then contacted with fresh resins (1-5 g) until all the organic solution was absorbed by the resins. This step was accomplished within 12 h in a drying oven at 333 K. These resins were finally evaporated to completely remove the solvent at 323 K in a vacuum for 2 h. The content of TOA held in the EIR phase was determined from the amount of HC1 adsorbed by shaking the EIR with 0.1 mol/dm3 HC1 (Akita et al., 1993). In this work, the concentrations of TOA and citric acid sorbed in the impregnated resin were expressed on the basis of the dry EIR. The aqueous solution was prepared by dissolving citric acid in deionized water without pH adjustment, and the initial acid concentration ranged from 0.002 to 0.024 mol/dm3. In the EIR phase, the initial TOA concentration varied from 0.257 to 0.945 mol/kg. Experimental Procedures. In the distribution experiments, the TOA-EIR (1g) and aqueous citric acid solution (50 cm3) were placed in a 125 cm3 glassstoppered flask and shaken at 110 rpm for at least 16 h using a thermostated shaker (Firstek Model B603, Taiwan). Preliminary experiments had shown that the sorption studied was complete after 12 h. After standing for 1 h, the aqueous phase was separated from the EIR and its equilibrium pH was measured with a pH meter (Radiometer Model PHM82). The concentration of citric acid in the aqueous phase was titrated with a known NaOH solution using Radiometer Autotitrator Model RTS82. The contents of citric acid sorbed and unreacted TOA in the EIR phase were calculated from a mass balance. Each experiment was duplicated under identical conditions.
Results and Discussion Preparation of the TOA-EIR. Figure 1 shows effect of the amount of TOA in impregnating solution on TOA content in the prepared XAD-4EIR. Similar results are observed in the case of XAD-mIR (not shown). The content of TOA in the EIR phase increases with the solution concentration. The amount of TOA transferred from organic solution to the resin is found to be more than 96% under the ranges studied. It should be noted that the resulting XAD-mIR and XAD4/EIR become adhesive after drying when the TOA content exceeds about 1.023 and 1.114 mol/kg, respectively. The present result for XAD-2/EIR is similar to that obtained by Akita and Takeuchi (1990) of 1.02 mol/ kg under comparable conditions. Also, the content of TOA analyzed by titration agrees well with that calculated from the changes in the weight of the resins before and after the impregnations.
n
y.c 1.5
0
1
41.5
z
$1 1.0
W
0 d
0.5
u
Amount of TOA (g) Figure 1. Effect of the amount of TOA in the impregnating solution (n-hexane) on TOA content in the XAD-mIR. Fresh resin = 1 g.
It is experimentally found that the equilibrium distribution of citric acid is nearly kept constant in a shaker speed range of 80-140 rpm (not shown). A more intensive contact between the EIR and aqueous solutions ('150 rpm) may cause a serious loss of TOA from the EIR. A shaker speed of 110 rpm is thus selected here, in which the TOA loss is less than 3%. In this work, the extractant loss from EIR under various TOA concentrations is also examined (not shown). The equilibrium distribution of citric acid is found to be abnormally low at lower TOA concentrations (especially in the XAD-4/EIR system, x0.470 molkg). This could be clearly seen in Figure 3. In fact, Cortina et al. (1992) also found that the Cyanex 272 (an organophosphinic acidbimpregnated XAD-2 resin has zero water content a t an extractant concentration of 0.8 molkg. Beyond that concentration the extractant is partially adsorbed into the macroporous resin and partially retained by the capillary forces inside the pores; namely, the EIR is more stable. Therefore, in this work TOA concentration in the EIR phase is kept high enough, greater than 0.47 mol/kg, as also in the work of Akita et al. (1993). In order to be an industrial process, the citric acid must be recoverable from the resins. Although that part is not the scope of the present work, it is worth noting that the desorption of citric acid from the loaded TOAEIR has been verified to be rapidly and completely achieved with either NazC03, HCI, HN03, o r HzS04 solution (0.1 moVdm3) from column operation (Juang and Chang, 1995). Sorption Equilibrium. The influences of citric acid concentration in the aqueous phase and TOA concentration in the EIR phase on equilibrium distribution are shown in Figures 2 and 3. It is found that the distribution of citric acid increases with both concentrations of TOA and citric acid. Moreover, the equilibrium distribution in both EIR systems is comparable under the same conditions. As discussed above, however, in the XAD-4/EIR system the distribution is abnormally low
1296 Ind. Eng. Chem. Res., Vol. 34, No. 4,1995
0.4
rcc
1
0,4
k
w-l
0
0
0.3
M
3
\
0.2
w n
0.0 0.000
0.005
0.01 0
0.01 5
0.020
[H3A] (mol/dm3) Figure 2. Effect of citric acid concentration in the aqueous phase and TOA concentration in the XAD-2/EIR system on the equilib= 0.471 ( O ) , 0.543 ( O ) , 0.652 rium distribution at 298 K. [=lo (O), 0.808 (01, and 0.942 ( A ) moykg.
2 \ 0
E
/
0.0 0.0
0.2
0.4
/
0.8
[mIo(mol/kg
0.8
1 .o
of EIR)
Figure 4. Maximum uptakes of citric acid under various TOA concentrations in the EIR phase a t 298 K.
are not chemically bonded to the polymeric matrix (Akita and Takeuchi, 1990, 1992; Akita et al., 1993; Cortina et al., 199313,199410;Juang and Su, 1992a). In this regard, the sorption of citric acid with TOA-EIR can be expressed by the following stoichiometric relation (Bizek et al., 1992a,b, 1993; Juang and Huang, 1994; Vanura and Kuca, 1976):
0.3
d
v
tt
0.2
n 0.1
I/If i" /
[H3A] (mol/dm3) Figure 3. Effect of citric acid concentration in the aqueous phase and TOA concentration in the XAD-I/EIR system on the equilibrium distribution a t 298 K. [=lo = 0.257 (O), 0.471 (01, 0.652 ( O ) , and 0.808 ( 0 )molkg.
-
where the overbar refers to the EIR phase. The sorption equilibrium constant, Kpq,by eq 1 is given by
at lower [TOAIo shown in Figure 3. The maximum uptakes of citric EIR phase are plotted in Figure 4, which also clearly -indicates that a reasonable result is obtained for [TOAIO> 0.47molkg. Although it has been found that the extracted species in the EIR system are less solvated than in organic solvent (Cortina et al., 1993b, 1994b), the mechanism of metal extraction with EIR is generally similar to that of the solvent extraction, provided that the extractants
The total equilibrium concentration of citric acid in the EIR phase is
P P
P 4
In eq 3, the contribution of the "physical" sorption of citric acid with TOA-EIR is ignored (Juang and Huang, 1994). In practice, the adsorption of citric acid with unimpregnated resins was measured in this study and are found to be, on an average, 3.38 x and 8.07 x molkg only for XAD-2 and XAD-4, respectively, at 298 K. They are less than those obtained for propionic, butyric, hexanoic, and heptanoic acids under comparable conditions (Hasanain and Hine, 1981). The term [TOAI in eq 3 is calculated by the following mass balance equation, assuming that the solubility of
Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 1297 0.8
o'6
~
r
~
7X 0.5
0.4
0.3
1
r
r
~
l
l
l
l
l
"
r
r
~
"
"
I
"
l
l
l
l
r
l
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[wA]o=0.543 mol/kp of EIR
\ 3
F 5 0.0 0.008
f I 1 1 1 I
0.013
0.01 8
0.023
(mol/dm3) Figure 5. Dependence of the loading of TOA in the XAD9BIR system on the concentration of citric acid in the aqueous phase. The meaning of each symbol is the same as that in Figure 2.
TOA in the aqueous solution is negligibly small (Akita and Takeuchi, 1990; Akita et al., 1993; Juang and Huang, 1994).
P
Q)
0.2
Q
5
(4) P
Q
A computer program was developed to determine parameters in the law of mass action (eqs 1 and 2) so as to minimize the error between the measured data and the results calculated from eqs 3 and 4. The object function, F, is defined by
[H3AIieXPt - [H3A]?lC
?[
[H3A]?lC
Various sets of stoichiometries of the complexes formed are postulated, and the minimization routine is performed for each set of stoichiometries to find the corresponding best-fit equilibrium constants and the error. Preference is given to complexes whose existence is inferred from the literature. In this work, the values o f p = 1-3 and q = 1-3 are tried (Bizek et al., 1992a,b, 1993; Juang and Huang, 1994; Sat0 et al., 1985;Vanura and Kuca, 1976). If the error is not significantly reduced by inclusion of a particular complex, that complex is not considered to be part of the best model. An additional consideration is that the model should represent physically reasonable complex formulation. The best-fit formulations for the acid-TOA complex (p,q) are obtained to be 1,1, 1,2, 1,3, 2,1, and 2,3. The value of F obtained for the best-fit is about 0.08.
[&AI o (m0l/dm3) Figure 6. Mole fraction of the complex formed in the XAD-2BIR system as a function of citric acid concentration in the aqueous phase - at 298 K. Curve 1, [(H&TOA)HE]t; 2,[(H&TOA)d [Hdlt; 3, t(H3A)(TOA)3HElt; 4, [(H3A)dTOA)Y[=lt; and 5, [(H3A)z(TOA)3l/[Elt.
It is found that the calculated equilibrium constants are equivalent in both EIR systems, which may be attributed to both comparable equilibrium distribution. They are calculated to be K11= 1.21 x lo2 dm3/mol,KIZ = 7.06 x 10 (dm3/moll2,K13 = 1.02 x lo3 (dm3/moU3, K21 = 2.81 x lo3 (kg/mol)(dm3/mol),and K23 = 4.52 x lo3 (kg/mol)(dm3/moU3,respectively, a t 298 K. In the case of solvent extraction of citric acid with xylene solutions of TOA, the equilibrium constants were found to be K11 = 0.68 dm3/mol, K12 = 1.68 (dm3/moU2,and K23 = 4.15 x lo2 (dm3/moU4,respectively, a t 298 K (Juang and Huang, 1994). Evidently, they are significantly smaller than the corresponding results obtained in EIR sorption. The present findings are supported by the loading curve shown in Figure 5. The loading curve is a plot of the loading of TOA vs log [HA], where the loading of TOA is defined as the total concentration of the acid sorbed a t equilibrium divided by the initial TOA concentration in the EIR phase. As indicated by Tamada et al. (1990), there is no effect of initial TOA concentration on the loading for systems with only one amine per
1298 Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995
0.5 Oe6
3
0.35
j
0.4
0.3
0.20
0.2
0.15
0.1
0.10
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1
0.00 0.000
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0.5
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0.05
0.005
0.010
0.01 5
0.020
[H3A] (mol/dm3) Figure 8. Effect of citric acid concentration in the aqueous phase and temperature on the equilibrium distribution. XAD-B/EIR, [=lo = 0.571 molflrg. '7 = 288 (O), 298 (O),308 (01,and 318 ( A )
K.
0.3
0.2
0.1 0.0 -l' 0.010
'
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I I
I
I
'
5 I
0.015
I I
I I
I I I I I
0.020
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I 6
1
'
1
0.025
Figure 7. Mole fraction of the complex formed in XAD-2/EIR (lower) and XAD-4/EIR (upper) systems as a function of citric acid concentration in the aqueous phase at 298 K. The meaning of each curve is the same as that in Figure 6 .
complex. On the other hand, if there is more than one amine per complex, the loading increases with TOA concentration at low acid concentrations. Systems that exhibit aggregation, formation of complexes with large numbers of acid and amine molecules, exhibit an abrupt increase in loading at a critical acid concentration, as encountered at [HsAI = 2 x mol/dm3in this study. The mole fraction of the species formed in the EIR phase as a function of initial concentrations of citric acid and TOA is shown in Figures 6 and 7. Under the ranges studied, the mole fractions of 1,2 and 2,3 complexes are always less than 0.1. With increasing [HAIo, the mole fraction of 1,l and 2,l complexes increases and that of the 1,3 complex decreases, but all finally reach a plateau. As clearly shown in Figure 6 , under the whole [HA10 ranges examined, the 1,l and 1,3 complexes dominate and the sum of the mole fraction of these two species is greater than about 0.75. With an increase in TOA concentration, nevertheless, the mole fraction of the 1,lcomplex decreases at an expense of the formation of the 2,l complex. It also follows from Figure 7 that the distribution of each complex has no detectable difference in XAD-B/EIR and XAD-4EIR systems.
In solvent extraction of citric acid with TOA in xylene (Juang and Huang, 19941, however, the mole fraction of the 1,2 complex is always less than 0.2 under -the ranges studied ([HA10 = 0.01-0.5 mol/dm3, [TOAIo = 0.05-0.5 mol/dm3). With increasing [Halo, the mole fraction of the 1,lcomplex decreases but that of the 2,3 complex increases. For the whole [H& ranges, the 1,l complex dominates for [TOAIo < 0.1 mol/dm3 -and the 2,3 complex becomes dominant at higher [TOAIo, especially for [TOAIo = 0.5 mol/dm3. In contrast to solvent extraction, the extremely small amount of the 2,3 complex present in EIR systems is believed to be due t o the steric hindrance for complex formation on the EIR surface or within the polymeric network (Cortina et al., 1994b; Tamada and King, 1990a). Effect of Temperature on Citric Acid Sorption. Figures 8 and 9 show the effect of temperature on equilibrium distribution in XAD-2EIR and XAD-4EIR systems, respectively. It is found that the amount of citric acid sorbed decreases as the temperature is increased. In general, the acid-TOA complexation reactions involve proton transfer or hydrogen-bond formation and are thus expected to be exothermic. Also, the formation of a complex makes the system more ordered and therefore decreases the entropy (Bizek et al., 199213; Juang and Huang, 1994; Tamada and King, 1990b). The apparent Gibbs free energy change, AG, for the sorption reaction at a given temperature is defined as
AG = -2.303RT log Kpq
(6)
where R is the universal gas constant and T is the absolute temperature. The enthalpy change, AH, is usually given by the van't Hoff relation: (7)
Ind. Eng. Chem. Res,, Vol. 34,No. 4,1995 1299 Table 1. Apparent Enthalpy and Entropy Changes for the Formation of the Complex (H&JTOA), in EIR Sorption and Solvent Extraction Systems in the 288-318 K Temperature Range e
complex
P=
G
@,q)
1,1 1,2 1,3 2,1 2,s
y.c
0
2? \ d
0
E
v
U
[H3A] (mol/dm3) Figure 9. Effect of citric acid concentration in the aqueous phase temperature on the equilibrium distribution. XAD-4/EIR, and [TOAIo = 0.671 molkg. The meaning of each symbol is the same as that in Figure 8.
10
10
10
( i / ~ > ~ i(K-') o~ Figure 10. Dependence of the sorption equilibrium constants, Kps, on temperature.
The entropy change, AS, is thus calculated from eq 8:
A S = (AH- AG)/T
(8)
Similar computer calculations were performed to treat the distribution data obtained at different temperatures. Figure 10 shows the temperature dependence of sorption equilibrium constants. The apparent enthalpy change and entropy change calculated in this study, together with those obtained in solvent extraction (Juang and Huang, 1994), are listed in Table 1 for
EIR sorption
solvent extraction
AH (kJ/mol) A S (J/mol K) AZI (kJ/mol) A S (J/mol K) -168.9 -95.6 -59.6 -88.4 -157.6
-525.6 -284.5 -142.1 -230.0 -457.2
-69.2 -35.5 -17.4
-237.5 -115.6 -8.40
comparison. For the present systems, 1:1 and 2:3 complexations are much more exothermic and involve a much greater loss of entropy than the formation of 1,2, 1,3, and 2,l complexes. The result for the case of the 1,l complex agrees with the findings of solvent extraction, which is attributed to the fact that 1:l complexation involves the formation of an ion pair but the higher complexes such as 1,2, 1,3, and 2,l involve hydrogen-bond formation (Tamada and King, 1990b). In contrast to the solvent extraction system, the exceptionally larger enthalpy change for 2:3 complexation obtained in this work could be a result of the steric effect on the EIR surface or within the polymeric network; futhermore, the more negative entropy changes for all complexation reactions in EIR systems could be due to the more restricted reaction environment. Sorption of Citric Acid with TOA and B2EHPA. BBEHPA (HR) is an organophosphoric acid. The extractability of citric acid with BBEHPA is negligibly low, as expected from the cation-exchange nature of BBEHPA. On the basis of our previous work on solvent extraction of citric acid with xylene solutions of TOA and BBEHPA (Juang and Huang, 1994), the synergistic - - and antagonistic effects are observed for [HRldlTOAlo e 1 and - [HRld[TOA10 > 1, respectively. This can be explained by the interaction between TOA and BSEHPA in the organic phase. Figures 11and 12 show the equilibrium distribution of citric acid between aqueous solutions and the resins impregnated with a mixture of TOA and BBEHPA in XAD-B/EIR and XAD-mIR systems, respectively. It is found that the presence of BBEHPA in the EIR phase has an antagonistic effect on the equilibrium distribu- tion even for [HRI~TOAIO e 1. In addition, it follows -Figure 11 that the smaller the value of [mid from [TOAIo the more pronounced that effect. The detailed mechanism leading to the difference between EIR sorption and solvent extraction still remains unknown at this stage. Presumably, the steric hindrance for the formation of some complex on the EIR surface may play an important role. This is supported by the fact that the antagonistic effect is comparatively remarkable for the resins with smaller pore size (i.e., XAD-4) a t high [H& (r0.015moVdm3), as shown in Figures 11 and 12. Another reason may be the interaction between BBEHPA and TOA in the EIR phase, which reduces the effective concentration of TOA for citric acid sorption and thus decreases the equilibrium distribution (Juang and Huang, 1994).
Conclusions The equilibrium of the sorption of citric acid between aqueous solutions and Amberlite XAD resins impregnated with TOA has been studied in the temperature range of 288-318 K. The following results are obtained.
1300 Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 0'25
3
n
E
fractions of both 1,2 and 2,3 complexes are less than 0.1 and the 1 , l and 1,3 complexes dominate. (3) The apparent enthalpy change and entropy change for the sorption of citric acid with TOA-EIR are given in Table 1. All sorption reactions are favored by entropy change and unfavored by enthalpy change. The 1:l and 2:3 complexation reactions are much more exothermic and involve a much greater loss of entropy than the formation of 1,2, 1,3, and 2,l complexes. (4) When B2EHPA is additionally impregnated onto the resins, an antagonistic effect on the distribution of citric acid is observed, in comparison with the impregnation of TOA only. This may be explained by the steric hindrance on the resin surface or within the pores.
v
Acknowledgment
n
This work was supported by the ROC National Science Council under Grant No. NSC83-0402-El55006, which is greatly appreciated.
Nomenclature
Figure 11. Effect of citric acid concentration in the aqueous phase on the equilibrium distribution with XAD-2l'EXin the presence = 0.543 molkg; [HRIo = 0 ( O ) , 0.272 of B2EHPA a t 298 K. [=lo (A), and 0.543 ( 0 )mobkg.
O"
0.5
B2EHPA = bis(2-ethylhexy1)phosphoric acid EIR = extractant-impregnated resin F = error of object function defined in eq 5 AG = apparent Gibbs free energy change for sorption reaction, kJ/mol AH = apparent enthalpy change for sorption reaction, kJ/ mol H a = citric acid HR = BPEHPA monomer Kpq= sorption equilibrium constant defined in eq 2, (dm3/ mol)D+q-l p = number of citric acid molecules involved in the complex q = number of TOA molecules involved in the complex A S = apparent entropy change for sorption reaction, J/(mol K) T = temperature, K TOA = tri-n-octylamine [ 3 = molar concentration of species in the brackets, mol/ dm3
L
0.2
Subscripts
t = total value at equilibrium 0 = initial n
Superscript
0.1
- (overbar) = EIR
phase
Literature Cited 0.0
0.00
0.01
0.02
0.03
Figure 12. Effect of citric acid concentration in the aqueous phase on the equilibrium distribution with XAD-4EIRin the presence of B2EHPA a t 298 K. [%%I0 = 0.771 mobkg; [HRlo = 0 ( 0 ) and 0.771 ( 0 )mol/kg.
(1)The impregnation of TOA on Amberlite XAD-2 and XAD-4 resins can be achieved by dry method to about 96% and to a content below 1.203 and 1.114 moVkg, respectively. (2) The equilibrium distribution of citric acid is comparable in both EIR systems. The distribution data can be best described by assuming the simultaneous formation of the acid-TOA complexes (HsA),(TOA), in the EIR phase as p , q = 1,1, 1,2, 1,3, 2,1, and 2,3. Under the conditions studied, the mole
Akita, S.; Takeuchi, H. Sorption and Separation of Metals from Aqueous Solution by a Macromolecular Resin Containing Trin-octylamine. J. Chem. Eng. Jpn. 1990,23(4), 439-443. Akita, S.; Hirano, K.; Ohashi, Y.; Takeuchi, H. Equilibrium Distribution of Palladium Between Hydrochloric Acid Solution and a Macromolecular Resin Containing Tri-n-octylamine. Solvent Extr. Ion Exch. 1993,11 ( 5 ) , 797-810. Baniel, A. M. Process for the Extraction of Organic Acids from Aqueous Solution. Eur. Patent Appl. EP 49,429, 1982. Baniel, A. M.; Gonen, D. Production of Citric Acid. US Patent US 4,994,609, 1991. Bizek, V.; Horacek, J.; Kousova, M.; Heyberger, A.; Prochazka, J. Mathematical Model of Extraction of Citric Acid with Amine. Chem. Eng. Sci. 1992a,47 (6), 1433-1440. Bizek, V.; Horacek, J.; Rericha, R.; Kousova, M. Amine Extraction of Hydroxy- carboxylic Acids. 1. Extraction of Citric Acid with 1-Octanolln-Heptane Solutions of Trialkylamine. Ind. Eng. Chem. Res. 1992b,31(61, 1554-1562. Bizek, V.; Horacek, J.; Kousova, M. Amine Extraction of Citric Acid: Effect of Diluent. Chem. Eng. Sci. 1993,48 (8), 14471457.
Ind. Eng. Chem. Res., Vol. 34, No. 4, 1995 1301 Cortina, J. L.; Miralles, N.; Sastre, A.; Aguilar, M.; Profumo, A.; Pesavento, M. Solvent Impregnated Resin Containing Cyanex 272. Preparation and Application to the Extraction and Separation of Divalent Metals. React. Polym. 1992, 18, 67-75. Cortina, J. L.; Miralles, N.; Sastre, A.; Aguilar, M.; Profumo, A.; Pesavento, M. Solvent Impregnated Resins Containing Di(2,4,4trimethylpenty1)phosphinic Acid. I. Comparative Study of Di(2,4,44rimethylpentyl)phosphinicAcid Adsorbed into Amberlite XAD-2 and Dissolved in Toluene. React. Polym. l993a, 21,89101. Cortina, J. L.; Miralles, N.; Sastre, A.; Aguilar, M.; Profumo, A,; Pesavento, M. Solvent Impregnated Resins Containing Di(2,4,4trimethylpenty1)phosphinic Acid. 11. Study of the Distribution Equilibria of Zn, Cu and Cd. React. Polym. 199313, 21, 103116. Cortina, J. L.; Miralles, N.; Aguilar, M.; Sastre, A. M. Solvent Impregnated Resins Containing Di(2-ethylhexy1)phosphoric Acid. I. Preparation and Study of the Retention and Distribution of the Extractant on the Resin. Solvent Extr. Zon Exch. 1994a, 12 (2), 349-369. Cortina, J. L.; Miralles, N.; Aguilar, M.; Sastre, A. M. Solvent Impregnated Resins Containing Di(2-ethylhexy1)phosphoric Acid. 11. Study of the Distribution Equilibria of Zn, Cu and Cd. Solvent Extr. Zon Exch. 199413, 12 (2), 371-391. Gao, H. Q.; Su, Y. F. Solvent Extraction and Adsorption by Extracting Resins of Phenol. Huadong Huagong Xueyuan Xuebao 1991, 17 (3), 249-254. Hartl, J.; Marr, R. Extraction Processes for Bioproduct Separation. Sep. Sci. Technol. 1993,28 (1-31, 805-819. Hasanain, M. A.; Hines, A. L. Application of the Adsorption Potential Theory to Adsorption of Carboxylic Acids from Aqueous Solutions onto a Macroreticular Resin. Ind. Eng. Chem. Process Des. Dev. 1981,20 (4), 621-625. Juang, R. S.; Su, J. Y. Sorption of Copper and Zinc from Aqueous Sulfate Solutions with Bis(2-ethylhexy1)phosphoric Acid-Impregnated Macroporous Resin. Znd. Eng. Chem. Res. 1992a, 31 (121, 2774-2779. Juang, R. S.; Su, J. Y. Separation of Zinc and Copper from Aqueous Sulfate Solutions with Bis(2-ethylhexy1)phosphoric Acid-Impregnated Macroporous Resin. Znd. Eng. Chem. Res. 1992b,31 (12), 2779-2783. Juang, R. S.; Huang, W. T. Equilibrium Studies on the Extraction of Citric Acid from Aqueous Solutions with Tri-n-octylamine. J. Chem. Eng. Jpn. 1994,27 (4),498-504. Juang, R. S.; Chang, H. L. Column Sorption of Citric Acid from Aqueous Solutions Using Tri-n-octylamine-Impregnated Macroporous Resins. Sep. Sci. Technol. 1995,30 (61,917-931. Kertest, A. S.; King, C. J. Extraction Chemistry of Fermentation Product Carboxylic Acids. Biotechnol. Bioeng. 1986, 28, 269282.
King, C. J. Amine-Based Systems for Carboxylic Acid Recovery. Chemtech 1992,22 (51, 285-291. Komiyama, H.; Smith, J. M. Surface Diffusion in Liquid-Filled Pores. AZChE J . 1974,ZO (6), 1110-1117. Kulprathipanja, S. Separation of Citric Acid from Fermenation Broth with a Neutral Polymeric Adsorbent. US Patent US 4,720,579, 1988. Sato, T.; Watanabe, H.; Nakamura, H. Extraction of Lactic, Tartaric, Succinic, and Citric Acids by Trioctylamine. Bunseki Kagaku 1985,34, 559-563. Tamada, J. A.; King, C. J. Extraction of Carboxylic Acids with Amine Extractants. 2. Chemical Interactions and Interpretation of Data. Znd. Eng. Chem. Res. 1990a, 29 (7), 1327-1333. Tamada, J. A.; King, C. J. Extraction of Carboxylic Acids with Amine Extractants. 3. Effect of Temperature, Water Coextraction, and Process Considerations. Znd. Eng. Chem. Res. 199013, 29 (7), 1333-1338. Tamada, J. A.; Kertes, A. S.; King, C. J. Extraction of Carboxylic Acids with Amine Extractants. 1. Equilibria and Law of Mass Action Modeling. Znd. Eng. Chem. Res. 1990,29 (71,1319-1326. Tavlaride, L. L.; Bae, J. H.; Lee, C.. K. Solvent Extraction, Membranes, and Ion Exchange in Hydrometallurgical Dilute Metals Separation. Sep. Sci. Technol. 1987,22 (2 and 3), 581617. Thurman, E. M.; Malcolm, R. L.; Aiken, G. R. Prediction of Capacity Factors for Aqueous Organic Solutes Adsorbed on a Porous Acrylic Resin. Anal. Chem. 1978,50 (6), 775-779. Vanura, P.; Kuca, L. Extraction of Citric Acid by the Toluene Solutions of Trilaurylamine. Collect. Czech. Chem. Commun. 1976,41,2857-2877. Warshawsky, A. Extraction with Solvent-Impregnated Resins. In Zon Exchange and Solvent Extraction; Marinsky, J. A., Marcus, Y., Eds.; Marcel-Dekker: New York, 1981; Vol. 8, pp 229-310. Wennersten, R. The Extraction of Citric Acid from Fermentation Broth Using a Solution of Tertiary Amine. J . Chem. Tech. Biotechnol. 1983, 33B, 85-94.
Received for review August 10, 1994 Revised manuscript received November 30, 1994 Accepted December 6 , 1994* IE940482L *Abstract published in Advance A C S Abstracts, March 1, 1995.
