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Ind. Eng. Chem. Res. 1997, 36, 4628-4637
Selective Sulfur Dioxide Removal Using Organic Solvents M. H. H. van Dam,† A. S. Lamine,† D. Roizard,‡ P. Lochon,‡ and C. Roizard*,† Laboratoire des Sciences du Ge´ nie Chimique and Laboratoire de Chimie Physique Macromoleculaire, CNRS-ENSIC, 1 rue Grandville, BP 451, F-54001 Nancy Cedex, France
It has been found that sulfur dioxide solubility in organic solvents can be well-correlated by Lewis basicity numbers. Mixtures of sulfur dioxide and solvent (N-methylpyrrolidone (NMP), N,N′-(dimethylpropylene)urea (DMPU) and methyldiethanolamine (MDEA)) have been prepared in a 1:1 mole ratio at low temperatures. The melting points of the mixtures are measured, and the mixtures have been analyzed by infrared, ultraviolet-visible, and nuclear magnetic resonance spectroscopy. The solubility of sulfur dioxide has been measured in NMP and DMPU at 25 °C. At sulfur dioxide concentrations of 0.2-0.6 kPa, the solubility is determined using a continuous stirred tank reactor. The absorption of sulfur dioxide proved to be completely reversible. A model where physical and chemical interactions are incorporated showed good correlation with the experimental results. I. Introduction Removal of sulfur dioxide from industrial gases is an increasingly important environmental challenge, on one hand, because of the lowering of the admissible emission limit, and, on the other hand, due to the fact that numerous desulfurization processes, such as limestone scrubbing, produce a large volume of solid waste. There is a growing interest in the use of organic solvents for sulfur dioxide removal. The idea of using organic solvents is not new; a process with dimethylaniline has already been patented in 1940 (Gleason et al., 1940). A process using tetraethyleneglycol dimethyl ether as a solvent has recently been commercialized (Heisel and Bellani, 1991). A major drawback of the last mentioned process is that the solvent is not selective to sulfur dioxide and other gaseous components are absorbed as well. In this article, a systematical approach is presented for selecting appropriate solvents without the abovementioned drawbacks. The solvent is to be used in a classical absorption/regeneration process (Figure 1) in order to recycle the solvent and produce a concentrated sulfur dioxide stream, which can be reused in other processes. Regeneration can be done by pressure reduction, by temperature rise, and by use of a vector gas. The choice of what to do with the concentrated sulfur dioxide stream depends on the situation of the specific industrial production site. The study of new solvents for sulfur dioxide absorption includes the following five steps: (1) preselection of solvents, (2) study of SO2-solvent interactions, (3) measurement of solubility, (4) measurement of absorption kinetics, and (5) study of regeneration. The present article covers the first three steps of this study.
Figure 1. Simplified flow scheme for SO2 absorption and regeneration.
allows us to classify solvents with respect to their capacity and selectivity toward sulfur dioxide. Since carbon dioxide is a major component in waste gases, sulfur dioxide solubility is compared to carbon dioxide solubility in order to examine the solvent selectivity. Luehrs and Godbole (1994) have correlated sulfur dioxide solubilities at atmospheric pressure with solvent parameters. Regarding our field of application (low sulfur dioxide gas concentrations), it is more appropriate to express the solubility in terms of Henry coefficients. Let ng be the initial number of moles of solute and ns the initial number of moles of solvent; the global composition of the liquid phase is given by
xg )
ng ng + ns
The Henry coefficient, H, is defined as
II. Preselection of Solvents The solvent should respond to the following criteria: high absorption capacity, high selectivity with respect to SO2 compared to NOx and CO2, easy to regenerate, low vapor pressure, and low toxicity. First, the absorption capacity and selectivity are dealt with. Our aim is to find an adequate parameter which * To whom correspondence should be addressed. E-mail:
[email protected]. Telephone: 33.(0)3.83.17.52.52. Fax: 33.(0)3.83.32.29.75. † Laboratoire des Sciences du Ge ´ nie Chimique. ‡ Laboratoire de Chimie Physique Macromoleculaire. S0888-5885(97)00111-5 CCC: $14.00
(1)
H ≡ lim xf0
() Pg xg
(2)
where Pg is the solute partial pressure. The Henry coefficients of sulfur dioxide for different solvents are given in Table 1. Attention has to be paid to the significance of these Henry coefficients: they are based on global mole fractions. Further on in this article, we will take a closer look at the meaning of the Henry coefficient. 1. Testing of General Solvent Parameters. First, the utility of general solvent parameters is studied. In © 1997 American Chemical Society
Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 4629 Table 1. Henry Coefficients of Sulfur Dioxide in Different Organic Solvents at 298 K solvent
HSO2, kPa
1,2-dichloroethane 1,2-ethanediol nitrobenzene tetramethylsulfone propylene carbonate diethyl ether tetrahydrofuran diethylene glycol-dimethyl ether triethylene glycol-dimethyl ether tetraethylene glycol-dimethyl ether tributyl phosphate dimethylformamide dimethylacetamide dimethyl sulfoxide dimethylaniline pyridine
428a 345a 166a 101b 122a/151b/148c 236a 39.2b 26.8d 25.7d 21.8d 34.5d 17.2b 11.9d 11.9a/10.9b 9.3e 8.8b
a Lenoir et al. (1971). b Benoit and Milanova (1979). c Xu et al. (1992) d Demyanovich and Lynn (1991) e Balej and Regner (1956).
Figure 2a-c, the sulfur dioxide and carbon dioxide solubilities are shown as a function of the dipole moment, the Hildebrand solubility parameter, and the empirical parameter of solvent polarity, ETN (Reichardt, 1990), respectively. Figure 2 shows clearly that these parameters do not describe sulfur dioxide solubility adequately. Another parameter, showing a better correlation with sulfur dioxide solubility, has to be found. 2. Testing of Parameters Based on the Lewis Basicity of the Solvent. Since sulfur dioxide is a Lewis acid, parameters based on the Lewis basicity of the solvent are tested with respect to the sulfur dioxide and carbon dioxide solubilities. Demyanovich and Lynn (1991) used the Gutmann donor number, DN (Gutmann and Schmid, 1974), for the prediction of infinite dilution activity coefficients of sulfur dioxide in organic solvents. The DN of a solvent is defined as the negative reaction enthalpy (in kcal/mol) of the formation of the 1:1 adduct between the solvent and antimony(V) chloride (SbCl5) in dilute 1,2-dichloroethane. This donor number has been criticized by Maria and Gal (1985), who have used boron trifluoride (BF3) as the acceptor instead of antimony(V) chloride and dichloromethane as the solvent instead of 1,2-dichloroethane. It is beyond the scope of this article to discuss in detail the “correctness” of the different basicity scales, and we simply test several of them with respect to their ability to describe solubility behavior. Next to the abovementioned basicity scales, two other scales are tested: Ds, defined as the difference between the symmetrical Hg-Br stretching frequencies of the neutral mercuric bromide complex in the gaseous phase and in a saturated solution of the studied solvent (Persson et al., 1987), and Cuλ,max, the wavelength in the visible region of the absorption maximum of the copper(II) N,N,N′,N′-tetraethylenediamineacetylacetonate complex dissolved in the pure solvents (Soukop and Schmid, 1985). Basicity parameters of all kinds of solvents have been collected (DN, Demyanovich and Lynn, 1991; Hahn et al., 1985; -∆BF3, Maria and Gal, 1985; Ds and Cuλ,max, Persson et al., 1987). In Figure 3 the Henry coefficients of sulfur dioxide and carbon dioxide are shown as a function of the different basicity parameters. Firstly, sulfur dioxide solubility shows a logarithmic dependency on the different basicity numbers, while carbon dioxide solubility is constant. Secondly, the four different basicity scales give the same qualitative results with
Figure 2. Henry coefficients of SO2 and CO2 as a function of the solvent dipole moment µ (a), the solvent Hildebrand solubility parameter dH (b), and the solvent polarity ETN (c) (Reichardt, 1990).
respect to the solubility behavior of sulfur dioxide and carbon dioxide, and any one of the four can be used. By these means, we have found a simple tool for screening solvents with respect to their capacity and selectivity for sulfur dioxide absorption. It should be noticed that at high basicity numbers, the solvent-sulfur dioxide interactions tend to become bonds of the covalent type, which can turn out to be a major drawback with respect to the regenerability of the solvent. Because of its wide availability in the literature, we use the Gutmann donor number as a screening parameter for choosing the appropriate solvents. 3. Other Criteria for Solvent Selection. Two other important criteria for solvent selection are the solvent vapor pressure and its toxicity. In Table 2, several solvents are presented with their DN, vapor pressure, and TLV (threshold limit for airborne concentration). From this table, three solvents have been selected, namely, N-methylpyrrolidone (NMP), N,N′(dimethylpropylene)urea (DMPU), and N-methyldiethanolamine (MDEA). For NMP, DN ) 27.3, and for DMPU, DN ) 29.6 (Hahn et al., 1985). No basicity number was found in the literature for MDEA.
4630 Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997
Figure 3. Henry coefficient of SO2 and CO2 vs basicity numbers of different organic solvents. (a) Henry coefficient vs Gutmann donor number DN, (b) Henry coefficient vs reaction heat of reaction with BF3 in dichloromethane, (c) Henry coefficient vs Ds, (d) Henry coefficient vs Cuλ,max. Table 2. Donor Numbers (DN), Vapor Pressure (pv), and and Treshhold Limits (TLV) for Different Solvents at 298 K solvent N-methylpyrolidone N,N′-(dimethylpropylene)urea pyridine dimethyl sulfoxide N,N-dimethylacetamide N-methyldiethanolamine
DN, kcal/mol
pv, kPa
TLV, ppmv
27.3 29.6 33.1 29.8 27.8
0.04 0.01a 2.8 0.08 0.27 99.5% GLC) from Carlo Erba, N,N′-(dimethylpropylene)urea (>99% GC) from Fluka, and N-methyldiethanolamine (>99%) from Janssen Chimica). The mixtures were kept at -20 °C. It was observed that the NMP-SO2 mixture was solid at this temperature. The MDEA-SO2 mixture, initially
liquid, crystallized slowly to end up completely solidified after 3 weeks. The DMPU-SO2 mixture stayed in liquid form. 1. Melting Point of Sulfur Dioxide-Solvent Mixtures. It is known that mixtures of sulfur dioxide with a Lewis-type base can give rise to so-called chargetransfer complexes. In these complexes, the solvent is the electron pair donor, while sulfur dioxide is the electron pair acceptor. Melting points of such complexes are reported in the literature. For 1:1 mole ratio sulfur dioxide-solvent mixtures, the following melting points can be found: N,N-dimethylaniline +12.6 °C (Seidell and Linke, 1952), pyridine -7.4 °C (Hoffman and Vanderwerf, 1946), dioxane +2.5 °C (Albertsen and Fernelius, 1943), aniline +65 °C (Foote and Fleischer, 1934). The melting points of some solvent-sulfur dioxide complexes can be found as a function of their composition (Hoffmann and Vanderwerf, 1946; Seidell and Linke, 1952). The melting points reach a maximum at the composition where the most stable complex is formed. Results and Discussion. The melting points of the three SO2-solvent complexes have been determined by heating slowly small solid samples of each 1:1 mixture in closed glass tubes with a thermostated bath at temperature increments of about 1 °C. The experimental melting point lies between -0.7 and +2.5 °C for the NMP-SO2 complex, whereas its value for pure NMP is -24 °C. The melting point of DMPU-SO2 lies between -21.5 and -20.0 °C. This is close to the melting point of DMPU, which was found to be between -18.0 and -17.0 °C. In the case of SO2-MDEA, the complex is much more stable, and after the already mentioned slow crystallization, its melting point is close to 80 °C. 2. Spectrophotometrical Techniques. 2a. Infrared and Ultraviolet/Visible Spectrophotometry. Generally, solutions containing interacting electron pair donor (D) and electron pair acceptor (A) species show
Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 4631
Figure 4. Two stretching bands of SO2: (a) in CCl4, (b) in NMP, (c) in DMPU.
not only the absorptions of D and A, sometimes noticeably modified, but also new bands, which are assigned to intermolecular charge-transfer transitions of the complex as a whole (Foster, 1969). The spectrophotometrical techniques allow us to get more information about the specific interaction site of the sulfur dioxide and solvents molecules. Only NMP-SO2 and DMPU-SO2 mixtures have been investigated by means of UV/vis. A small quantity of each mixture is diluted in cyclohexane in a 1-cm cup and a spectrum is recorded from 800 to 200 nm at ambient temperature. Because of the strong absorption when using 30-mm NaCl cells, the IR spectra have been recorded from a thin film of sulfur dioxide-solvent mixture between two KBr pellets at ambient temperature. This means that only qualitative information could be obtained from the recorded spectra. Results and Discussion. The characteristic bands of the solvents and SO2 were identified with UV/vis, but no information on a complexation reaction could be obtained. It was observed that the DMPU-SO2 mixture
was much less soluble in cyclohexane than in pure DMPU, which might be explained by a complexation reaction between sulfur dioxide and DMPU. The recorded IR spectra are shown in Figures 4-6. Two stretching bands of the sulfur dioxide are observed at 1323 and 1146 cm-1 in the 1:1 SO2-NMP mixture and at 1321 and 1144 cm-1 in the SO2-DMPU mixture, while these bands are found at 1344 and 1145 cm-1 for SO2 in noncomplexating CCl4 (Figure 4a-c). Comparing the SO2-NMP spectrum with the pure NMP spectrum, it is observed that the amide I band is shifted to a lower wavelength, i.e., from 1685 cm-1 in pure NMP to 1668 cm-1 in the mixture, under the influence of sulfur dioxide (Figure 5a). The band due to the stretching vibration of the DMPU carbonyl group is also shifted to a higher frequency under the influence of SO2 (from 1638 to 1606 cm-1, Figure 5b). This indicates that the electron pair donor site is located at the oxygen atom in the NMP and the DMPU molecule. After exposing the samples to air, the amide band shift decreased rapidly. This observation
4632 Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997
Figure 5. Amide I band in (a) NMP and NMP-SO2 and in (b) DMPU and DMPU-SO2.
can be related to the progressive disappearance of the characteristic SO2 absorption bands. The spectrum of the SO2-MDEA 1:1 mixture has been recorded before crystallization took place. The spectrum (Figure 6a) exhibits a new band at 655 cm-1. The characteristic bands of sulfur dioxide are not found, as was the case for the other two solvents. After a first exposure to air, the spectrum changed drastically and new peaks were observed at 1174, 969, and 437 cm-1. Again, this is in contrast with the observations of the NMP-SO2 and DMPU-SO2 mixtures. After another exposure to air, no further change in the spectrum is observed, and it is different from the spectrum of pure MDEA (Figure 6b). It is not clear how to explain the observations for the MDEA-SO2 mixture. It might be
Figure 6. Spectrum of (a) MDEA-SO2 before and after exposure to air, and (b) MDEA-SO2 after exposure to air and pure MDEA.
possible that absorbed water from ambient air plays a role. 2b. Nuclear Magnetic Resonance. When the molecular environment of a nucleus undergoes a rapid reversible change, the position of the magnetic resonance absorption of the nucleus represents a timeaveraged resultant of its behavior in the different environments (Foster, 1969). We consider the simple case of an equilibrium between a 1:1 AD complex (acceptor-donor) and its components A (SO2) and D (D ) solvent):
A + D a AD
(I)
When the forward and reverse reactions are very fast, the measured (mean) chemical shift of a nucleus in the
Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997 4633
donor molecule is the averaged mean of the chemical shift of the pure donor molecule, δD, and of the chemical shift of the pure complex of acceptor and donor, δAD:
δ)
xD xAD δD + δ xD + xAD xD + xAD AD
(3)
where xD is the mole fraction of uncomplexed donor and xAD is the fraction of complexed donor. The mole fractions are related to each other by the equilibrium constant of reaction I, defined as
K)
xAD xAxD
(4)
The global composition of the mixture gives the following two equations (n0A and n0D: the initial number of moles of pure components):
xA + xAD n0A ) xD + xAD n0
(5)
xA + xD + xAD ) 1
(6)
D
The three mole fractions can be expressed as functions of K and the ratio n0A/n0D (see Appendix). Varying the acceptor concentration and thus xD and xAD and measuring the chemical shift of the pure solvent and of the mixtures, both K and δAD can be estimated by minimizing the difference between the measured and calculated values of δ from eqs 3-6. 13C NMR spectra are recorded at different sulfur dioxide mole fractions in the solvent. Starting with the original 1:1 mixtures, the sulfur dioxide concentration has been decreased by bubbling dry argon through the mixture in the NMR tube in order to remove sulfur dioxide. The fraction of remaining sulfur dioxide has been determined by weighing the tube. Varying the sulfur dioxide concentration in the MDEA mixture was not possible; the sulfur dioxide simply stayed in solution. Results and Discussion. Because the SO2 influence on the CdO chemical shift is the largest for NMP and DMPU, the chemical shift of this carbon atom is used in order to determine the equilibrium constant. In Figure 7, the measured chemical shift is given as a function of sulfur dioxide concentration for the SO2NMP and SO2-DMPU mixtures. The agreement between the experimental points and the model is quite satisfactory (average deviation 2%). The obtained values of the fitted parameters are K ) 12 and δAD ) 176.15 ppm for the NMP-SO2 complexation reaction and K ) 18 and δAD ) 157.25 ppm for the DMPUSO2 complexation reaction. IV. Determination of SO2-Solvent Solubilities In this part, we will focus on the third step of the study: the solubility of sulfur dioxide in organic solvents. Although in industry the sulfur dioxide concentration in the gas phase does not exceed a few volume percents, the solubility will be measured up to 100 kPa of sulfur dioxide pressure in order to find a reliable model, which includes all aspects such as physical solubility, Henry’s law, deviation from Raoult’s law, and chemical reaction.
Figure 7. Measured and modeled δ-δD vs SO2 mole fraction for NMP-SO2 (a) and DMPU-SO2 (b) mixtures.
1. Theory. Consider a gaseous component A, which absorbs into a liquid phase:
A(g) a A(l)
(II)
The liquid phase is now a binary system of solute and solvent. If the mixture would show an ideal behavior, the partial pressure of the gaseous component, PA, is the product of the mole fraction of the solute, xA, and the vapor pressure of pure liquid solute, Psat A (Raoult’s law):
PA ) xPsat AA
(7)
At modest pressures, this relation provides a reasonable approximation only when the components of the mixture are physically and chemically similar. Often, serious deviations occur because the pure components differ in molecular size, shape, and intermolecular forces, and thus, this equation should rather be used as a reference and not as a predictive equation. In order to describe deviations of Raoult’s law, three approaches are possible, i.e., deviation caused by (1) physical interaction, (2) chemical reaction, and (3) physical and chemical interactions. 1a. Henry’s Law and the Duhem-Margules Equation. In order to describe the thermodynamic equilibrium at a gas-liquid interface, the equality of the gas and liquid chemical potentials should be used. But, often, the solubilities of the gaseous compounds are simply expressed in terms of Henry coefficients, based on global mole fractions (eq 2). Often, the liquid phase doesn’t show ideal behavior, which results in a deviation from linear behavior when the solute concentration increases (Balej and Regner, 1956; Cooper and Smith, 1974; Fogg and Gerrard, 1991; Harris and Prausnitz, 1969). This deviation of linearity is expressed as an
4634 Ind. Eng. Chem. Res., Vol. 36, No. 11, 1997
interaction parameter, R, in the Duhem-Margules equation in the following form (Fogg and Gerrard, 1991):
ln
()
PA ) ln(H) - R(1 - xD2) xA
(8)
Under the present experimental conditions, it is supposed that the gas phase shows ideal behavior. Note. Since nonideality is described by physical interaction only, xA and xD represent the global mole fractions of respectively sulfur dioxide and solvent. 1b. Chemical Reaction. Another approach is to model the solubility behavior by means of absorption followed by a reversible chemical reaction between solute and solvent molecules. For the absorption of A, eq 8 is used with R ) 0. Note. In this approach, xA represents the mole fraction of uncomplexed sulfur dioxide. The other two species present in solution are uncomplexed solvent D and the complex AD. The mole fractions of these species are related by the equilibrium constant given by eq 4. 1c. Combination of the Two Models. In the literature, the solubility behavior is sometimes modeled by eq 8 (sulfur dioxide in acetone (Fogg and Gerrard, 1991)) and sometimes by absorption followed by chemical reaction (sulfur dioxide in tri-n-butylphosphate (Cooper and Smith, 1974)). However, it is very likely that, when a complexation of the charge transfer type takes place, deviation from Raoult’s law is caused by physical and chemical interactions. Harris and Prausnitz (1969) combined physical and chemical interactions in a model in order to describe the activity coefficients for acetylene in three organic solvents. They used Van Laar’s theory to describe the deviation from Raoult’s law. Here, the Duhem-Margules rule is used to describe the physical part of the interactions between solute and solvent. Thus, eqs 4 and 8 are combined, which allows us to model the solubility behavior by weak (physical) and strong (chemical) interactions. 2. Experimental Part. Two different experimental installations have been used in order to obtain sufficient precision over the whole pressure range. The first installation consists of a stirred tank reactor, in which the solubility is measured at low sulfur dioxide concentrations in the gas phase (up to 1 kPa). The second experimental setup consists of a bubble column reactor and allows measurements from 2 to 100 kPa of sulfur dioxide pressure. The sulfur dioxide solubility in two pure solvents was measured: NMP and DMPU. 2a. Solubility at 0.2-0.6 kPa. Sulfur dioxide solubility is measured using a 10-4-m3-volume (diameter ) height ) 0.05 m) agitated stirred tank reactor, thermostated at 25 °C by a water jacket (Figure 8). Starting with a pure unloaded solvent, a gaseous mixture of sulfur dioxide and nitrogen is bubbled through the reactor at a flow rate of about 2.5 10-5 m3 s-1. The reactor is closed with respect to the liquid phase. The outlet sulfur dioxide concentration is analyzed by UV spectrophotometry and recorded as a function of time by a computer until equilibrium is reached, that is, when the outlet concentration becomes constant. The amount of sulfur dioxide absorbed is calculated from the gas-phase mass balance. 2b. Solubility up to 100 kPa of Sulfur Dioxide Pressure. A cylindrical Pyrex reactor (diameter ) 0.02 m, height ) 0.14 m) has been built, as shown in Figure 9. The procedure proposed by Fogg and Gerrard (1991)
Figure 8. Scheme of the experimental setup for measuring SO2 solubilities at low SO2 pressures (
