Ind. Eng. Chem. Res. 1997, 36, 2353-2358
2353
Enhancement of the Absorption of CO2 in Alkaline Buffers by Organic Solutes: Relation with Degree of Dissociation and Molecular OH Density Gonzalo Va´ zquez,* Francisco Chenlo, and Gerardo Pereira Department of Chemical Engineering, University of Santiago de Compostela, Avda. das Ciencias, s/n, E-15706 Santiago (La Corun˜ a), Spain
Absorption of CO2 in a wetted-wall column by 0.5 M Na2CO3/0.5 M NaHCO3 buffer with and without various concentrations of saccharose, fructose, glucose, formaldehyde, glycerin, methanol, or ethanol was measured under conditions in which the reaction of CO2 was of pseudo-firstorder. For the purposes of comparison, experiments were also carried out with arsenite in both buffer and pure water. For all of these solutes, the absorption enhancement factor increased with solute concentration. Rate constants kc for the overall reaction were obtained by fitting Danckwerts’ expression for absorption to the experimental data, and correlation of kc with solute concentration then afforded rate constants kcat for the catalyzed reactions. For each solute, an empirical correlation was found between the absorption enhancement factor and Hatta numbers calculated from the rate constant for the uncatalyzed reaction and from k ) kcat[cat]. The notion that the enhancement of absorption by acid-base reaction is facilitated by dipole-dipole interaction between the solutes and CO2 is supported by an empirical correlation relating kcat to the degree of dissociation of the solute and the density of OH groups in the solute molecule. Introduction Absorption of CO2 by liquid phases is a common industrial process, improvement in the efficiency of which is always of interest (Giavarini and Moresi, 1981). This process has generally been studied in water (Lee, 1991) or alkaline solutions (Yonemoto et al., 1986; Hagewiesche et al., 1995); the most usual absorbent is probably carbonate/bicarbonate buffer (Danckwerts, 1970), which is moderately alkaline and undergoes no change in pH during absorption of CO2. In this buffer solution dissolved CO2 reacts both with water and the hydroxyl ion
CO2 + H2O a HCO3- + H+
(a)
CO2 + OH a HCO3
(b)
-
-
and chemical absorption can, however, be enhanced by means of homogeneous catalysts by the addition of solutes such as arsenite (Sharma and Danckwerts, 1963), hypochlorite (Pohorecki, 1968; Benadda et al., 1994), or carbonic anhydrase (Alper and Deckwer, 1980). Since the works of Astarita et al. (1981, 1982) and Danckwerts (1981), it is generally accepted that in aqueous solution these species, or their hydration or dissociation products, act as catalysts whose reaction with CO2 is followed by a process in which HCO3- is released as the catalyst is regenerated. All of these species feature O- or OH groups, all can act as Lewis bases for CO2 as Lewis acid (some through atoms with electron lone pairs), and all have a pyramidal or tetrahedral structure (or tetrahedral carbon units) facilitating the approach of the CO2 molecule to the basic site. Figure 1 shows the catalytic reaction for an active species XO-. Another class of additives employed in absorption studies comprises the organic solutes used to control the viscosity of the liquid phase. In studies of the effect of * E-mail:
[email protected];
[email protected]. S0888-5885(96)00635-5 CCC: $14.00
Figure 1. Scheme of the catalytic reaction process for a general active species XO-.
viscosity on interfacial area as measured by Danckwerts’ method (Rizzuti and Brucato, 1989; Va´zquez et al., 1989), which requires the use of buffered absorbents because the absorption reaction must be of pseudo-firstorder with respect to the absorbed gas, it has been observed that certain solutes of this kind increase absorption of CO2 dose-dependently (Va´zquez et al., 1992; Alvarez, 1993). This effect was assumed to be related to the structural similarities between these organic solutes and the inorganic and enzymatic solutes mentioned above, namely the presence of OH groups and the tetrahedral nature of their carbon atoms. However, unlike inorganic enhancers such as arsenious acid, the organic additives are so weakly acidic that they fail to generate the effective catalysts (the corresponding anions) in pure water, so that CO2 absorption in their aqueous solutions remains a purely physical process (Limin˜ana et al., 1987; Va´zquez et al., 1991); they only exhibit catalytic action in alkaline buffers. To investigate this process further, we have carried out a systematic study of the absorption of CO2 in a wetted-wall column at 25 °C and atmospheric pressure by 0.5 M Na2CO3/0.5 M NaHCO3 buffer with and without various concentrations of sodium arsenite, saccharose, fructose, glucose, commercial formaldehyde (formalin), glycerin, methanol, and ethanol. The wetted-wall column is a suitable contact device for absorp© 1997 American Chemical Society
2354 Ind. Eng. Chem. Res., Vol. 36, No. 6, 1997
tion studies because its laminar flow allows easy calculation of contact areas and exposure times. Some of the organic compounds chosen are cyclic and bear several OH groups; some of those that are acyclic likewise bear several OH groups whereas others bear only one, and all have tetrahedral carbon atoms (formaldehyde exists in aqueous solution as H2C(OH)2). The observed absorption enhancement factors were correlated with Hatta numbers calculated from the rate constants for the uncatalyzed reaction and the catalytic process, and the rate constants for the catalytic process were correlated with the degree of dissociation of the organic absorption promoters and a measure of the density of OH groups in the solute molecule.
Table 1. Physical Properties of a Solution of Organic CO2 Absorption Enhancers in 0.5 M Sodium Carbonate/ 0.5 M Sodium Bicarbonate Buffer, with the Corresponding Values of kc, the Overall Rate Constant for the Reaction of CO2a solute
r ) -d[CO2]/dt ) kH2O[CO2] + kOH[OH-][CO2] (1) If the reaction takes place at constant [OH-], as in carbonate/bicarbonate buffers, which maintain a practically constant pH ) 9.6, it is therefore pseudo-first-order with respect to CO2, and eq 1 may be written
r ) (kH2O + kOH[OH-])[CO2] ) ko[CO2]
(2)
Under pseudo-first-order conditions in the presence of a catalyst we write
r ) (ko + kcat[cat])[CO2] ) kc[CO2]
0.09 0.18 0.27 0.36 0.45
1079.0 1079.8 1080.3 1080.6 1080.9
1.299 1.315 1.331 1.347 1.363
19.38 19.38 19.38 19.38 19.38
1.56 1.55 1.53 1.52 1.50
3.02 6.16 8.49 11.16 13.88
saccharose
0.05 0.09 0.15 0.21 0.25 0.31
1078.7 1084.5 1091.8 1100.1 1106.1 1113.6
1.358 1.439 1.533 1.640 1.731 1.868
18.94 18.63 18.21 17.82 17.61 17.27
1.51 1.44 1.36 1.29 1.23 1.16
2.74 4.81 6.81 9.20 11.10 13.60
fructose
0.08 0.17 0.25 0.35 0.43
1081.2 1086.1 1090.9 1097.7 1103.0
1.376 1.443 1.511 1.607 1.680
18.96 18.53 18.22 17.94 17.71
1.49 1.43 1.38 1.31 1.26
2.60 4.10 6.02 7.90 9.59
glucosea
0.08 0.17 0.25 0.35 0.43
1081.6 1085.2 1089.2 1097.3 1102.8
1.384 1.447 1.509 1.597 1.669
18.90 18.47 18.20 18.00 17.84
1.48 1.43 1.38 1.32 1.27
2.40 4.22 5.87 7.80 9.20
glycerin
0.11 0.26 0.41 0.60
1076.1 1079.4 1082.4 1086.2
1.342 1.384 1.439 1.527
18.94 18.63 18.38 18.14
1.52 1.48 1.44 1.37
1.31 1.93 2.52 3.29
methanol
0.31 0.46 0.62
1073.9 1073.9 1073.9
1.302 1.310 1.318
19.38 19.38 19.38
1.56 1.55 1.54
1.17 1.29 1.47
ethanol
0.20 0.30 0.40 0.50
1073.9 1073.9 1073.9 1073.9
1.352 1.381 1.410 1.439
19.30 19.26 19.23 19.19
1.51 1.49 1.46 1.44
1.25 1.42 1.70 1.95
(3)
Values of ko and kc were obtained by optimizing the fit between the experimental absorption data and the expression for the quantity of gas absorbed per unit contact area in exposure time te
Q)
Nte ) kLCeteE A
(4)
with the enhancement factor function obtained by applying the surface renewal theory when gas absorption is due to a pseudo-first-order reaction
(
E ) Ha 1 +
) (x )
π erf 8Ha2
(
)
4Ha2 4Ha2 1 + exp π 2 π
(5)
and Hatta number
Ha ) xπkte/4
(6)
For a wetted-wall column with laminar flow, the exposure time and total contact area depend only on column geometry and the physical properties and flow rate of the liquid phase: 3
2 te ) h 3
x
3π2d2µ q2Fg
(7)
3
3µq xπdFg
A ) πhd + 2πh
(8)
k c, 1/s
formaldehyde
Definition and Determination of Kinetic Constants Under conditions in which reactions a and b are more or less equally significant (8 < pH < 10), the overall reaction rate is given by
F, µ × 103, Ce × 103, D × 109, [ ], m2/s kmol/m3 kg/m3 kg/(m‚s) kmol/m3
a
Values taken from Alvarez (1993).
For each catalyst, an optimum value of kc was sought by the Nelder-Mead method to minimize the mean square deviations between the calculated values of Q from eqs 4-8 and the experimental values corresponding to each value of liquid flow rate, and the values of kc were used to calculate values of ko and kcat, using eq 3, to correlate kc with solute concentration. Experimental Section The absorption of CO2 by 0.5 M Na2CO3/0.5 M NaHCO3 buffer with and without various concentrations of saccharose, fructose, glucose, commercial formaldehyde (formalin), glycerin, methanol, and ethanol, Table 1, was measured in a wetted-wall column at 298 K and atmospheric pressure. For the purposes of comparison, similar measurements were carried out for sodium arsenite in both buffer and pure water, Table 2. Note that the formalin solutions contained both the concentrations of formaldehyde listed in Table 1 and 17% by weight of methanol as stabilizer; in what follows, it will always be stated whether or not the methanol content of the formalin solutions has been taken into account in deriving or displaying parameters obtained from the experimental data for formalin. The wetted-wall column used, with a diameter of 8 mm and a height of 150 mm, has been described elsewhere (Va´zquez et al., 1991, 1992). The gas phase was CO2 of purity >99.95% that had been saturated with water vapor prior to being fed into the column at
Ind. Eng. Chem. Res., Vol. 36, No. 6, 1997 2355 Table 2. Physical Properties and pH of Solutions of Sodium Arsenite in Water and in 0.5 M Sodium Carbonate/0.5 M Sodium Bicarbonate Buffera F, µ × 103, Ce × 103, D × 109, [NaAsO2], m2/s liquid kmol/m3 kg/m3 kg/(m‚s) kmol/m3 pHo water
0.0150 0.0300 0.0600
997.3 997.4 997.5
0.896 0.897 0.898
33.60 33.60 33.60
1.92 1.92 1.92
10.64 10.80 10.93
buffer
0.0025 0.0050 0.0075 0.0100 0.0125 0.0150
1073.9 1073.9 1073.9 1073.9 1073.9 1073.9
1.284 1.284 1.284 1.284 1.284 1.284
19.38 19.38 19.38 19.38 19.38 19.38
1.58 1.58 1.58 1.58 1.58 1.58
9.60 9.60 9.60 9.60 9.59 9.58
kc , 1/s
1.60 2.43 3.23 4.10 4.68 5.50
a For the buffer solutions, the corresponding values of k , the c overall rate constant for the reaction of CO2, is also shown.
a flow rate of about 1.5 × 10-6 m3/s; absorption rate was determined as the difference between the measured gas flow rates at the entry to and the outflow from the column. The liquid phase was supplemented with 5 × 10-3% by weight of the surfactant sodium lauryl sulfate so as to prevent the formation of surface wavelets (Jyotirmay et al., 1987; Llorens et al., 1988). The densities of the liquid phases, Tables 1 and 2, were measured with a Bosch S2000/30 densimetric balance and their viscosities with a Schott AVS350 capillary viscosimeter. The results were in keeping with published values (Va´zquez et al., 1994a). The low concentrations of arsenite employed in this work caused no detectable change in the density or viscosity of the liquid phases, and the changes in buffer density caused by addition of ethanol and methanol were likewise negligible. The solubility of CO2 in solutions of saccharose, fructose, glucose, glycerin, and ethanol, Table 1, was calculated from its solubility in unsupplemented buffer (Hikita et al., 1974, 1976), assuming the same functional dependence on solute concentration as has been reported for unbuffered aqueous solutions of these solutes (Stephen and Stephen, 1964; Va´zquez et al., 1994b). It was assumed that methanol, formaldehyde, and arsenite have no effect on CO2 solubility. The diffusivity of CO2 in all the absorbent solutions, Tables 1 and 2, was calculated from the expression (Kamal and Canjar, 1966; Joosten and Danckwerts, 1972)
DµR ) cte
(9)
with R ) 0.82 (Ratcliff and Holdcroft, 1963; Alper, 1981). Experimental Results and Discussion Sodium Arsenite. When arsenite is added to aqueous medium, successive hydration equilibria produce the active catalyst H2AsO3- and its conjugate acid (Piechura and Bub, 1980):
AsO2- + H2O a H2AsO3-
(f)
H2AsO3- + H2O a H3AsO3 + OH-
(g)
In this work, addition of arsenite to unbuffered water increased CO2 absorption dose-dependently, Figure 2. Unbuffered arsenite solutions are alkaline due to reaction g, but the enhancement of CO2 was not due simply to OH- and reaction b, because NaOH solutions with
Figure 2. Rates of CO2 absorption by water, NaOH solution, and various solutions of sodium arsenite, for various liquid flow rates.
Figure 3. Rates of CO2 absorption by buffer solutions with and without various concentrations of commercial formaldehide, for various liquid flow rates.
similar OH- concentrations failed to enhance CO2 absorption, Figure 2. Furthermore, the consumption of hydroxyl ions in reaction b led to a fall in pH down the column of only about 1 unit in the case of arsenite solutions, as against about 5 units in the case of NaOH solutions. Hence the enhancement of CO2 absorption by arsenite solutions may be attributed to reaction c between CO2 and H2AsO3-, according to Figure 1. However, the overall reaction cannot be considered as catalytic sensu stricto because, due to the production of protons by reaction d or consumption of OH- by reaction e, reaction g proceeds from left to right and [H2AsO3-] falls. Arsenite also enhanced CO2 absorption by carbonate/ bicarbonate buffers. In this case the overall reaction is properly catalytic because the constant pH of these solutions ensures the maintenance of equilibrium g. Fitting equations 4-8 to the data and correlating the resulting kc values, Table 2, with NaAsO2 concentration afforded the expression
kc ) 0.87 + 310[NaAsO2]
(10)
the constants of which are in keeping with previously published values (Alper et al., 1980).
2356 Ind. Eng. Chem. Res., Vol. 36, No. 6, 1997
Figure 4. Rates of CO2 absorption by buffer solutions with and without various concentrations of glycerin, for various liquid flow rates.
Figure 5. Rates of CO2 absorption by buffer solutions with and without various concentrations of ethanol, for various liquid flow rates. Table 3. Rate Constant ko and kcat Obtained for Each Organic Solute by Correlating kc against [cat] solute
ko, 1/s
kcat, m3/(kmol‚s)
saccharose fructose glucose glycerin
0.83 0.87 0.88 0.86
40.84 20.21 19.59 4.07
solute
k o, 1/s
ethanol 0.83 methanol 0.86 formaldehyde 0.86
kcat, m3/(kmol‚s) 2.16 0.97 28.20
Organic Additives. CO2 absorption by carbonate/ bicarbonate buffers was increased dose-dependently by saccharose, fructose, glucose, and formalin (Figure 3 shows the effect of formalin), was reduced by glycerin (Figure 4), and was not significantly affected by methanol or ethanol (Figure 5 shows the effect of ethanol). Table 1 lists the kc values obtained by fitting eqs 4-8 to the data, and Table 3 the values of ko and kcat obtained by correlating kc with solute concentration (the values for formaldehyde in Tables 1 and 3 have taken into account the methanol content of the formalin). Note that although overall CO2 absorption is reduced by glycerin, kcat for glycerin is positive, showing that
Figure 6. CO2 absorption enhancement factors E plotted against the Hatta number. Table 4. Optimized Parameters of the Expression of Eq 11 for Each Solute solute
a
b
c
f
arsenite saccharose fructose glucose glycerin ethanol methanol formaldehyde
0.414 0.414 0.414 0.414 0.414 0.414 0.414 0.414
2 2 2 2 2 2 2 2
0.365 0.362 0.363 0.363 0.369 0.383 0.380 0.364
1.85 1.74 1.77 1.78 1.90 1.99 1.96 1.71
the overall reduction is due to the effect of the additive on the physical properties of the system. Hatta Numbers and Enhancement Factors. The Hatta numbers calculated from the kc values of the various solutes and their corresponding values of enhancement factors, Figure 6, indicate the occurrence of moderately fast reactions at the interface (Charpentier, 1981) and hence zero concentration of dissolved CO2 in the bulk liquid. For each solute the experimental enhancement factor increased with solute concentration and with exposure time. If the total absorption of CO2 is the effect of water, buffer, and promotor, the joint influence of these three factors was empirically modeled by fitting the data for each solute with an expression of the form
E ) 1 + aHaBb + cHacatf
(11)
where HaB and Hacat are the Hatta numbers calculated from eq 6 with k ) ko and k ) kcat[cat], respectively. The experimental values of E differ by less than 1% from those afforded by the fitted expressions, and the latter differ chiefly in the value of f, Table 4. Relationship between Solute Reactivity and Structure. The enhancement of CO2 absorption by the organic solutes in carbonate/bicarbonate buffers is attributed to the general mechanism of Figure 1, following deprotonation of the solute by OH-. As was pointed out in the Introduction, it is the need for this prior deprotonation step that makes an alkaline medium essential for the efficacy of the organic solutes, which have no absorption-enhancing effect in pure water. The need for the prior deprotonation step might also suggest that the efficacy of the organic solutes depends on their acidity, but there is no correlation between the kcat
Ind. Eng. Chem. Res., Vol. 36, No. 6, 1997 2357 Table 5. pKa, Molecular Mass, and Number of OH Groups of Each Solute solute
pKa
M, kg/kmol
NOH
arsenite saccharose formaldehyde fructose glucose glycerin ethanol methanol
9.23 12.35 9.75 12.30 12.50 13.70 15.98 15.55
126 342 48 180 180 92 46 32
3 8 2 5 5 3 1 1
Figure 7. Values of kcat given by the correlation equation (12) plotted against the experimental values, for each solute.
values of Table 3 and the corresponding pKa values of Table 5, which were taken from Weast and Astle (1987) and Budavari et al. (1989), and Dennard and Williams (1966) similarly found no correlation between the efficacy of a series of enhancers and a number of molecular properties (acidity, geometry, number of oxygens bonded by the central atom, etc.). Under the assumption that step c of Figure 1 is the rate-controlling step, it is possible that the catalytic reaction may be facilitated not only by the acidity of the enhancer but also by the ability of the active species to capture CO2 molecules via dipole-dipole interactions, which would increase the probability of encounter. To explore this possibility, we sought a correlation of the form
kcat ) xδyβ-z
(12)
where δ ) Ka/(Ka + [H+]) is the degree of dissociation of the solute and, as a rough measure of the capacity of the solute to capture CO2 molecules via polar interactions, β is defined as the OH density of the solute molecule (on a mass basis, for simplicity)
β)
NOHm M
(13)
where NOH is the number of OH groups in the solute molecule, M is its molecular mass, and m is the molecular mass of the OH group (which has been introduced merely to make β dimensionless). Using the measured pH of the buffer solutions and the Ka values of Table 5 (after prior verification by measurement of the pH of unbuffered aqueous solutions of the solutes), fitting eq 12 to the experimental data afforded the
values x ) 9.83, y ) 0.35, and z ) 3.95, with excellent fit even for arsenious acid, Figure 7. In conclusion, the enhancement of CO2 absorption in carbonate/bicarbonate buffer solutions by a series of organic solutes has been satisfactorily correlated with their degree of dissociation and their OH density, the latter parameter as a measurement of their capacity for polar interaction with CO2. Nomenclature A ) gas-liquid transfer area, m2 a ) optimized parameter of eq 11 b ) optimized parameter of eq 11 Ce ) solubility of CO2 in the liquid phase, kmol/m3 c ) optimized parameter of eq 11 D ) diffusivity of CO2 in the liquid phase, m2/s d ) diameter of the wetted-wall column, m E ) CO2 absorption enhancement factor f ) optimized parameter of eq 11 g ) acceleration due to gravity, m/s2 Ha ) Hatta number h ) height of the wetted-wall column, m Ka ) acidity constant, kmol/m3 kc ) overall pseudo-first-order rate constant for the chemical reaction of CO2 with the liquid phase, 1/s kcat ) rate constant of the catalytic reaction, m3/(kmol‚s) kH2O ) rate constant of reaction (a), 1/s kL ) mass transfer coefficient, m/s ko ) rate constant for the uncatalyzed reaction in buffer solution, 1/s kOH ) rate constant of reaction b, m3/(kmol‚s) M ) molecular mass of the solute, kg/kmol m ) molecular mass of the OH group, kg/kmol N ) CO2 absorption rate, kmol/s NOH ) number of OH groups per solute molecule pHo ) initial pH of the liquid phase Q ) quantity of CO2 absorbed per unit contact area in exposure time, kmol/m2 q ) flow rate of the liquid phase, m3/s r ) reaction rate, kmol/(m3‚s) te ) gas-liquid exposure time, s x ) optimized parameter in eq 12, m3/(kmol‚s) y ) optimized parameter in eq 12 z ) optimized parameter in eq 12 [ ] ) concentration in the liquid phase, kmol/m3 R ) parameter in eq 9 β ) molecular OH density, on a molecular mass basis δ ) degree of dissociation µ ) viscosity of the liquid phase, kg/(m‚s) F ) density of the liquid phase, kg/m3
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Received for review October 9, 1996 Revised manuscript received February 12, 1997 Accepted February 27, 1997X IE9606350
X Abstract published in Advance ACS Abstracts, April 15, 1997.
