Article pubs.acs.org/jced
Solubility Properties and Spectral Characterization of Dilute SO2 in Binary Mixtures of Urea + Ethylene Glycol Shaoyang Sun,† Yanxia Niu,†,‡ Fei Gao,§ Jun Shen,‡ and Xionghui Wei*,† †
Department of Applied Chemistry, College of Chemistry and Molecular Engineering, Peking University, Beijing 100871, China College of Chemistry and Chemical Engineering, Taiyuan University of Technology, Taiyuan 030024, China § Key Laboratory for Green Chemical Technology of the Ministry of Education, School of Chemical Engineering & Technology, Tianjin University, Tianjin 300072, China ‡
ABSTRACT: Solubilities of SO2 in binary mixtures of urea and ethylene glycol were determined by gas−liquid equilibrium experiments at the temperatures from (298.15 to 318.15 K) and 122.7 kPa, with SO2 partial pressure in gas phase in the range of (0 to 130) Pa. Henry’s law constants were calculated based on the solubility data. It indicated that the mixture with urea molality of 6.65 mol·kg−1 has the most desirable capacity. More than 90 % SO2 can be regenerated by bubbling N2 at 338.15 K. Density and viscosity data for the binary mixtures were also measured. In addition, UV, IR, and NMR experiments were conducted to investigate the interaction between ethylene glycol, urea, and SO 2 . The result demonstrated that SO2 can interact with the mixture by both the charge-transfer interaction between S (SO2)···N (urea) and the hydrogen bond between O (SO2)···H (ethylene glycol).
1. INTRODUCTION Sulfur dioxide, emitted from the combustion of fossil fuels, has brought about tremendous damage to human health and ecological environment for many years. Therefore, the disposal of SO2 emission has become a key issue of global concern. At present, flue-gas desulfurization (FGD) techniques using limestone as absorbent are widely applied to control SO2 emission in industry. However, there are some drawbacks in traditional FGD processes, such as high capital investment and operating cost, useless byproducts (CaSO4 and CaSO3), and wastewater.1−3 Thus, some environmental friendly and energyefficient methods are proposed to be alternatives to current FGD techniques. Recently, room temperature ionic liquids (ILs)4−6 have attracted much interest due to their advantages, such as low vapor pressure, high thermal and chemical stability, good solubility, and high selectivity to SO2. However, the high viscosity and expense of ionic liquids prohibit their application in industry. Because of the favorable absorption and desorption properties of acid gases, organic solvents have become the subject of increasing interest for many years.7−9 Ethylene glycol (EG)/ polyethylene glycol (PEG) and their aqueous solutions are considered to be a kind of promising absorbents for FGD for their advantages of reversible absorption, low price, and nontoxicity. Research measuring SO2 solubility data and the interaction mechanism between SO2 and solvent molecules has been reported in the literature.10−13 The results indicated that the absorption of SO2 is based on physical absorption, so © XXXX American Chemical Society
absorbents can be regenerated by either reducing pressure or increasing temperature. However, because of the physical absorption, the absorption capacity of dilute SO2 in pure EG and its water solution is very low. Considering that the outlet temperature of flue gas is high and the partial pressure of SO2 is low, an improvement of the absorption capacity of EG solution is still needed. Urea is a kind of basic industrial material widely used in chemical industry and agriculture. Because of the weak basicity, its aqueous solution has an effective absorption property of acid gas. In consideration of the excellent absorption properties of urea and EG, their binary mixtures may be a promising alternative in SO2 removal. Here, we report urea and EG binary mixtures to absorb SO2 in order to increase the absorption capacity. Gas−liquid equilibrium (GLE) was used to determine the solubilities of SO2 in binary mixtures of various compositions at different temperatures and a constant pressure of 122.7 kPa. In general, the partial pressure of SO2 ranges from several hundreds to thousands parts per million in flue gas, so the GLE data was measured with SO2 partial pressure of gas phase in the range of (0 to 130) Pa (about 1000 ppm). Henry’s law constants were calculated based on the solubility data. Regeneration was conducted by bubbling N2 at a high temperature. Density and viscosity data for the binary mixtures Received: September 26, 2014 Accepted: December 17, 2014
A
DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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the experiment process, the system temperatures are controlled using a thermostatic bath and determined using a standard thermometer with an uncertainty of 0.01 K. System pressures are inspected by a meter with accuracy of 0.133 kPa. 2.2.2. Data Treatment and Sources of Error. The FPD detector of GC was calibrated with the various standard gas mixtures (SO2 + N2) before being applied to determine the partial pressure of SO2 in the gas phase. The calibration results showed that GC method presented a high stability with 1.0 % uncertainty, and the calibration curve was linear by double logarithm method. The calibration relation was expressed as y = ax + b, where x represented the logarithm of the SO2 peak area, y represented the logarithm of the volume fraction of SO2 in the gas phase, and a and b were the equation coefficients. In this work, the maximum deviation value was defined as
at the temperatures ranging from (298.15 to 318.15) K were also measured. Meanwhile, the mechanism of interactions between SO2 and mixtures were also discussed by spectral characterization. As far as we know, there is no report on SO2 absorption by the binary mixtures of EG and urea, and we believe these renewable and cheap solutions with high SO2 absorption capacity can be potential absorbents in industry.
2. EXPERIMENTAL SECTION 2.1. Materials. A certified standard gas SO2 in N2 (8000 ppmv) which was supplied by Beijing Gas Center, Peking University (China), was employed to determine the GLE data of dilute SO2 in urea + EG mixtures. The ethylene glycol (EG) and urea were purchased from Beijing Yili Fine Chemical Co., Ltd., and Beijing Chemical Works, respectively, and directly used without further purification. EG and urea were both AR grade, and the purities were more than 99.9 %. The purity of EG was also checked by measuring its densities and viscosities at the temperature ranging from (298.15 to 318.15) K. The values are listed in Table 1 and are in good agreement with the literature values.14−24 Distilled water and chromatographic grade ethanol were used in the present work.
y′ = max((ymax − y ̅ ), (y ̅ − ymin ))
where y max and y min were the maximum and minimum measurements of SO2 partial pressure in the gas phase, respectively, and y ̅ was the average of eight measurement values. They were calculated by the below formula:
Table 1. Comparison of Experimental Densities (ρ) and Viscosities (η) of Ethylene Glycol with Literature Values at Various Temperaturesa ρ/(g·cm−3) T/K 298.15
expt. 1.1102
303.15
1.1066
308.15
1.1030
313.15
1.0993
318.15
1.0958
lit.
ymax = max(y1,exp , y2,exp , ..., y8,exp )
(2)
ymin = min(y1,exp , y2,exp , ..., y8,exp )
(3)
8
η/(mPa·s)
Ethylene Glycol 1.1100914 1.110015 1.1099116 1.106617 1.1066218 1.1066514 1.103019 1.1029614 1.1028920 1.0993522 1.0997516 1.0957221 1.0960222
(1)
y̅ =
expt.
lit.
16.2
16.123 16.79519
∑ yi ,exp /8
(4)
i=1
The maximum relative error was obtained by δ = (y′ − y ̅ )/y ̅
24
14.1
13.8678
11.6
11.695624 11.33319
9.75
9.534824
7.91
7.453222
(5)
2.3. Desorption Experiment. To illustrate the desorption effect of the binary mixture, the desorption experiment was carried out using the gas stripping way at the temperature of 338.15 K and N2 flow rate of 2 L·min−1. 2.4. Density and Viscosity Measurements. Densities of pure liquid and their mixtures were determined using a 25 cm3 bicapillary pycnometer. The volume of pycnometer was calibrated using distilled, deionized, and degassed water at each temperature. The pycnometer filled with liquid was kept in a thermostatic controlled and well-stirred water bath for 20 min to attain thermal equilibrium. The density measurements were carried out at T from (298.15 to 318.15) K. Each experimental density value was an average of at least three measurements. The uncertainty of the density measurement was estimated to be 0.01 %. The kinematic viscosities in both the pure component and their mixtures were measured using an Ubbelohde capillary viscometer, calibrated with high pure water and chromatographic grade ethanol at the experimental temperature whose viscosity and density were well-known, as has been described in literatures.26−31 The flow time of the solutions was recorded by an electronic timer with an accuracy of 0.01 s. The same circulating water bath applied in the density measurement was used for the viscosity measurement. Measurements were repeated at least 10 times at each temperature for all solutions, and the final results are the average values of ten measurements. The kinematic viscosity (ν) was calculated from the following equation
a Standard uncertainties for data in this experiment are u(T) = 0.01 K, u(ρ) = 0.0001 g·cm−3 and u(η) = 0.02 mPa·s.
Binary mixtures were prepared by mass using an analytical balance (Sartorius BS 224S) with 0.0001 g accuracy. The possible error in molality of urea for each binary mixture is less than 0.1 %. 2.2. Solubility Measurements. 2.2.1. Apparatus and Experimental Procedures. The solubility data of SO2 in the binary mixtures of urea + EG at the temperature ranging from (298.15 to 318.15) K and the system pressure of 122.7 kPa were measured, and the experimental instrument and process were identical to the literature.25 The gas phase (SO2/N2) was recycled in the absorption bottle by the gas circulatory pump to achieve the gas−liquid equilibrium. The SO2 concentration in gas phase was determined by using an Agilent GC, while the SO2 concentration in the liquid phase was determined with indirect iodometric titration. The relative uncertainty of SO2 molality in the liquid phase was estimated to be 0.6 %. During
ν = At − B /t B
(6) DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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where ν was the kinematic viscosity; t was the flow time of liquids in the viscometer; A and B were instrument constants calculated from measurements with the calibration fluids of water and ethanol. The absolute viscosity (η) was calculated by multiplying the kinematic viscosity by the corresponding density (η = νρ). It was estimated that the uncertainty of the viscosity measurement to be lower than 0.3 %. 2.5. Spectral Measurements. The mechanism of interactions between molecules in the system of EG + urea + SO2 were studied by spectrometric methods. FTIR spectra were recorded on a Bruker VECTOR22 spectrometer with 4 cm−1 resolution, 16 times background scans, and 32 times sample scans from 4000 to 400 cm−1. UV−vis spectra were recorded on a UNIC 4802 UV spectrophotometer. 1H NMR spectra were obtained using an AVANCE III Bruker-500 MHz nuclear magnetic resonance spectrometer, and d6-DMSO was used as the NMR solvent. All spectral experiments were performed at room temperature and atmospheric pressure.
Figure 2. Solubility curves on dilute sulfur dioxide in urea + EG systems at 308.15 K and 122.7 kPa: □, m1 = 0.000 mol·kg−1; ○, m1 = 0.834 mol·kg−1; △, m1 = 1.644 mol·kg−1; ▽, m1 = 2.799 mol·kg−1; ☆, m1 = 4.096 mol·kg−1; ◇, m1 = 5.833 mol·kg−1; +, m1 = 6.648 mol·kg−1; −, the linear smoothed line.
respectively. Their corresponding data are listed in Tables 3 and 4. In Figure 3, the SO2 solubility decreases as the absorption temperature rises, which means that the absorbed SO2 can be desorbed from the liquid when heated. Moreover, the solubility value is almost 1.53·10−2 mol·kg−1 at 318.15 K when the partial pressure of SO2 in gas phase is close to 1100 Pa. This value is still more than the solubility of SO2 in pure EG even at the temperature of 298.15 K, which reveals that the solution with m1 = 6.648 mol·kg−1 will be superior to pure EG for SO2 absorption. In addition, it is also noticed that the molality of SO2 in liquid phase drops linearly with the SO2 partial pressure decline within all measured temperatures. Figure 4 presents the relationship between SO2 solubility and the system total pressure. It is clear that there are few differences for SO2 solubility under different system pressure. It means that the effect of system pressure on SO2 absorption in the studied solution is less than that of the absorption temperature. In other words, the absorbed SO2 is desorbed from liquid more easily by heating method. 3.2. Desorption Result. The absorbed SO2 can be desorbed from the solution with m1 = 6.648 mol·kg−1 by heating and N2 bubbling, and the result of desorption experiment is given in Figure 5. It is clear in Figure 5 that the about 90 % SO2 molecules escape from the investigated solution in the experimental condition. This means that not only the absorbed SO2 can be recovered but also the absorbent solution can be used again. This information is very useful for selecting the candidate absorbent. 3.3. Density and Viscosity Data. For a solution used for capture SO2, the properties of density and viscosity are the same significance as the absorption capability. Figure 6 shows the results of densities on those solutions with m1 ranging from (0.000 to 7.137) mol·kg−1 at the temperatures from (298.15 to 318.15) K, and the corresponding measurements are given in Table 5 and 6. As shown in Figure 6, the density of the mixture increases with the urea molality and declines as the measured temperature arises. Figure 7 describes the results of viscosity on these solutions at the same temperature range. The change trend of viscosity with urea molality and temperature is similar to that of the density. Here, it is remarkable that the viscosity value of the solution with m1 = 6.648 mol·kg−1 is about 16.8 mPa·s at 308.15 K, which is slightly higher than that of EG at 298.15 K (16.2 mPa·s). Furthermore, the viscosity of the solution with m1 = 6.648 mol·kg−1 decreases to 10.8 mPa·s at T = 318.15 K which is even lower than the viscosity of EG at
3. RESULTS AND DISCUSSION 3.1. Solubility Data. A series of solubility data of dilute SO2 in the binary mixtures of urea + EG were measured at temperature range from (298.15 to 318.15) K and 122.7 kPa. A comparison on the solubility of SO2 in pure EG between this work and previous work32 is carried out and shown in Figure 1. It is clear that two sets of data are well in line.
Figure 1. Comparison of two sets of data between this work and previous work on the solubility of SO2 in pure EG at 308.15 K and 122.7 kPa: △, this work.○, previous work.
Figure 2 gives the relationship between SO2 partial pressure in the liquid phase and its partial pressure in the gas phase, and the corresponding data are shown in Table 2. As in Figure 2, it can be seen that the solubility of SO2 increases with the molality of urea rising under the same SO2 partial pressure. When the molality of urea in the mixture (m1) is equal to 6.648 mol·kg−1, the solution is almost at saturation at 293.15 K. For this solution, the solubility value of SO2 in liquid phase approaches to 2.04·10−2 mol·kg−1 with the SO2 partial pressure equal to 110 Pa, which is significantly larger than that of SO2 in pure EG (3.28·10−3 mol·kg−1)32 under the same conditions. It is also found from Figure 2 that the solubility curves fit well with the linear smoothed lines within the investigated range of SO2 partial pressure here. Because the solution with m1 = 6.648 mol·kg−1 has a desirable absorption capability for SO2, solubility experiments of SO2 in this solution were also performed at the temperature from (298.15 to 318.15) K and the system pressure from (122.7 to 138.7) kPa. These results are given in Figures 3 and 4, C
DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 2. Solubility Data of SO2 in Urea (1) + EG (2) Systems at 308.15 K and 122.7 kPaa m1b mol·kg
mSO2 −1
0.000 0.000 0.000 0.000 0.000 0.000 0.834 0.834 0.834 0.834 0.834 0.834 0.834 1.644 1.644 1.644 1.644 1.644 1.644 2.799 2.799 2.799
10 ·ΦSO2 6
381 516 657 716 805 888 257 354 502 632 706 825 895 206 460 683 763 856 924 292 410 504
10
−3
mol·kg
1.26 1.85 2.37 2.99 3.54 4.37 2.80 4.14 5.35 5.84 6.55 7.70 9.36 3.48 6.01 8.33 9.30 10.64 11.45 5.58 6.57 7.85
m1b
pSO2 −1
Pa
mol·kg
46.8 63.3 80.6 87.8 98.8 109 31.5 43.4 61.6 77.5 86.6 101 110 25.3 56.4 83.8 93.6 105 113 35.8 50.3 61.9
mSO2 −1
10 ·ΦSO2
2.799 2.799 2.799 4.096 4.096 4.096 4.096 4.096 4.096 5.833 5.833 5.833 5.833 5.833 5.833 5.833 6.648 6.648 6.648 6.648 6.648 6.648
6
577 690 824 242 449 690 768 870 968 167 334 477 644 785 854 933 221 397 534 656 849 906
10
−3
mol·kg
8.95 9.92 11.43 6.27 10.10 14.05 15.28 16.51 18.00 6.42 9.39 11.83 15.04 16.96 18.61 19.48 8.26 11.69 14.55 16.45 18.70 20.95
pSO2 −1
Pa 70.7 84.7 101 29.7 55.1 84.6 94.2 107 119 20.5 41.0 58.6 79.0 96.3 105 114 27.2 48.6 65.5 80.5 104 111
Standard uncertainties are u(T) = 0.01 K, u(p) = 0.1 Pa, u(m1) = 0.001 mol·kg−1, and u(mSO2) = 0.2·10−3 mol·kg−1. bm1 is the molality of urea in binary mixtures of urea + EG. a
308.15 K (11.6 mPa·s). These results, along with the above solubility results, illustrate that the solution with m1 = 6.648 mol·kg−1 has good absorption performance for SO2 as well as a low viscosity at 308.15 K or higher temperature. 3.4. Thermodynamical Model. Thermodynamical parameters including Henry’s law constant, standard Gibbs free energy (ΔG), phase change enthalpy (ΔHH), and phase change entropy (ΔSH) were calculated for the absorption processes in these studied mixtures at the temperatures between (298.15 and 318.15) K and the system total pressure of 122.7 kPa. These thermodynamical models are deduced as follows: In general, the Henry’s law constant (HLC) is important to evaluate the absorption capability of a solvent for a certain gas. It is defined as
Figure 3. Solubility curves on dilute SO2 in the solution with m1 = 6.648 mol·kg−1 at the temperatures from (298.15 to 318.15) K and 122.7 kPa: □, 298.15 K; ○, 303.15 K; △, 308.15 K; ▽, 313.15 K; ◇, 318.15 K; −, the linear smoothed line.
HLC = p /mSO2
(7)
where HLC is the Henry’s law constant, with the units of Pa·kg· mol−1, mSO2 is the molality of the absorbed gas in liquid phase, and p is the partial pressure of the studied gas in gas phase. Equation 7 implies that the solubility of gas which behaves nearly ideally is linearly related to the pressure. Just as shown in Figures 2 to 4, the absorption processes of SO2 in these solutions follow Henry’s law under the investigated conditions, and these Henry’s law constants can be acquired by calculating the slope of the linear solubility curves. According to ref 33, the dimensionless Henry’s law constant (H′) is defined as H′ = Cg /C l Figure 4. Solubility curves on dilute SO2 in the solution with m1 = 6.648 mol·kg−1 at 308.15 K and the system total pressure from (122.7 to 138.7) kPa: □, 122.7 kPa; △, 128.0 kPa; ☆, 133.3 kPa; +, 138.7 kPa.
(8)
where Cg is the gas concentration in gas phase, it can be derived from the ideal gas law (p = [n/V]RT = Cg·RT). Correlating the two expressions of 7 and 8, H′ is obtained as D
DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 3. Solubility Data of SO2 in the Liquid with m1 = 6.648 mol·kg−1 at T from (298.15 to 318.15) K and ptotal = 122.7 kPaa T
a
mSO2
K
10 ·ΦSO2
298.15 298.15 298.15 298.15 298.15 298.15 303.15 303.15 303.15 303.15 303.15 303.15
155 371 481 593 746 851 190 336 490 623 757 880
6
−3
10
pSO2
mol·kg
−1
9.55 16.78 19.88 22.63 27.02 30.98 8.80 12.60 16.05 18.29 21.83 23.81
T
mSO2
Pa
K
10 ·ΦSO2
18.9 45.5 59.0 72.8 91.5 104 23.3 41.3 60.1 76.4 92.8 108
313.15 313.15 313.15 313.15 313.15 313.15 318.15 318.15 318.15 318.15 318.15 318.15
198 429 486 649 727 917 209 407 574 686 830 911
6
−3
10
pSO2
mol·kg
−1
6.98 10.94 11.81 14.67 15.36 19.01 7.77 8.97 11.50 13.07 14.37 15.59
Pa 24.3 52.6 59.6 79.6 89.2 112 25.6 49.9 70.4 84.1 102 112
Standard uncertainties are u(T) = 0.01 K, u(p) = 0.1 Pa, and u(mSO2) = 0.2·10−3 mol·kg−1.
Table 4. Solubility Data of SO2 in the Liquid with m1 = 6.648 mol·kg−1 at 308.15 K and ptotal from (128.0 to 138.7) kPaa p
a
mSO2
kPa
10 ·ΦSO2
128.0 128.0 128.0 128.0 128.0 128.0 128.0 133.3 133.3 133.3
89.4 170 290 413 555 680 856 221 273 383
6
10
−3
mol·kg
pSO2 −1
6.49 7.57 11.43 13.16 15.84 17.92 21.21 8.40 10.13 12.47
p
mSO2
Pa
kPa
10 ·ΦSO2
11.4 21.8 37.2 52.8 71.0 87.1 110 29.5 36.4 51.1
133.3 133.3 133.3 133.3 138.7 138.7 138.7 138.7 138.7 138.7
521 667 768 850 177 211 375 641 710 847
6
−3
10
mol·kg
15.41 18.10 20.09 21.73 8.13 8.83 11.77 16.53 17.92 20.09
pSO2 −1
Pa 69.5 88.9 102 113 24.6 29.2 51.9 88.9 98.4 117
Standard uncertainties are u(pSO2) = 0.1 Pa, u(p) = 0.1 kPa, and u(mSO2) = 0.2·10−3 mol·kg−1.
Figure 6. Densities of the urea + EG systems versus the molality of urea in the mixture (m1) at the temperatures from (298.15 to 318.15) K: □, 298.15 K; ○, 303.15 K; △, 308.15 K; ▽, 313.15 K; ☆, 318.15 K.
Figure 5. Desorption curve of SO2 out of the solution with m1 = 6.648 mol·kg−1 at 338.15 K and under N2 flow rate of 2 L·min−1.
H′ = HLC·ρ /RT
(9)
where R is the ideal gas constant (8.314 Pa·m3·mol−1·K−1), T is the absolute temperature (K), and ρ is the density of solvent (kg·m−3). Actually, for our systems investigated here, the phase equilibrium equation of SO2 is described as
SO2 (g) ↔ SO2 (L)
where CSO2 is the molality of SO2 absorbed in the liquid phase and p is the SO2 partial pressure in the equilibrium state. So the Henry’s law constant (H′) is the reciprocal of dimensionless equilibrium constant. The standard Gibbs free energy change (ΔG) when SO2 reaches phase equilibrium can be acquired by
(10)
ΔG = −RT ln(1/H′)
Thus, the equilibrium constant (K) is expressed as
K = CSO2 /p
(12)
Based on the van’t Hoff equation (ΔG = ΔH − TΔS), the following expression is obtained:
(11) E
DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 5. Experimental Densities (ρ) of Urea (1) + EG (2) at T from (298.15 to 318.15) K and ptotal = 122.7 kPaa ρ/g·cm−3
m1b mol·kg
−1
0.000 0.877 1.852 2.939 4.162 5.549 7.137
T/K = 298.15
T/K = 303.15
T/K = 308.15
T/K = 313.15
T/K = 318.15
1.1102 1.1184 1.1279 1.1372 1.1463 1.1553 1.1644
1.1066 1.1148 1.1244 1.1337 1.1429 1.1519 1.1611
1.1030 1.1111 1.1207 1.1300 1.1391 1.1481 1.1574
1.0993 1.1078 1.1174 1.1268 1.1359 1.1449 1.1541
1.0958 1.1039 1.1136 1.1229 1.1321 1.1410 1.1503
Standard uncertainties are u(T) = 0.01 K, u(ρ) = 0.0001 g·cm−3, and u(mSO2) = 0.2·10−3 mol·kg−1. bm1 is the molality of urea in binary mixtures of urea + EG.
a
Table 6. Experimental Viscosities (η) of Urea (1) + EG (2) at T from (298.15 to 318.15) K and ptotal = 122.7 kPaa η/mPa·s
m1b mol·kg
−1
0.000 0.877 1.852 2.939 4.162 5.549 7.137
T/K = 298.15
T/K = 303.15
T/K = 308.15
T/K = 313.15
T/K = 318.15
16.2 17.0 18.5 20.4 22.1 24.0 25.7
14.1 14.9 16.0 17.5 18.8 20.2 21.5
11.6 12.1 13.0 14.1 15.1 16.1 17.1
9.75 10.1 10.8 11.6 12.4 13.2 13.9
7.91 8.13 8.68 9.31 9.86 10.4 11.0
Standard uncertainties are u(T) = 0.01 K, u(η) = 0.02 mPa·s, and u(mSO2) = 0.2·10−3 mol·kg−1. bm1 is the molality of urea in binary mixtures of urea + EG. a
Figure 7. Vicosities of the urea + EG systems versus the molality of urea in the mixture (m1) at the temperatures from (298.15 to 318.15) K: □, 298.15 K; ○, 303.15 K; △, 308.15 K; ▽, 313.15 K; ☆, 318.15 K.
ln H′ = ΔH /RT − ΔS /R
Figure 8. Temperature dependence of Henry’s law constant (H′) for the absorption process of dilute SO2 in the solution with m1 = 6.648 mol·kg−1 under the system pressure of 122.7 kPa.
(13)
Table 7. Henry’s Law Constant (H′) and Standard Gibbs Free Energy (ΔG) for the Absorption Processes of SO2 in Urea + EG Systems at 308.15 K and 122.7 kPa
This equation describes the dependence of H′ on temperature, where ΔH and ΔS represent the enthalpy and entropy of the phase change from the gas phase to the liquid phase, respectively. Because the temperature range here is small, so the enthalpy and entropy of the phase change are independent of temperature, and their values can be obtained from the slope and the intercept of the plotted curve on ln H′ versus 1/T (see Figure 8), respectively. All of these thermodynamical functions values are given in Tables 7 and 8. It is found that ΔG are always below zero and more than −40 kJ·mol−1, which reveals hat the absorption processes of SO2 are spontaneous and reversible under the given conditions. The results indicate that SO2 can be captured by these liquids from a mixed gas stream and also recovered by heating. Moreover, it is also found that the calculated values of the enthalpy and entropy are negative, which indicates that the
m1a/mol·kg−1 0 0.834 1.644 2.799 4.096 5.833 6.648 a
103 H′
ΔG/kJ·mol−1
± ± ± ± ± ± ±
−12.7 −13.8 −14.2 −14.3 −15.2 −15.4 −15.5
7.00 4.59 3.84 3.75 2.61 2.41 2.34
0.69 0.42 0.13 0.15 0.08 0.04 0.13
m1 is the molality of urea in binary mixtures of urea + EG.
absorption process of SO2 is an exothermic process. All of above these results are very valuable for a liquid as a candidate used for absorbing SO2 from a gaseous stream. F
DOI: 10.1021/je5009119 J. Chem. Eng. Data XXXX, XXX, XXX−XXX
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Table 8. Henry’s Law Constant (H′), Standard Gibbs Free Energy (ΔG), Phase Change Enthalpy (ΔH), and Phase Change Entropy (ΔS) for the Absorption Processes of SO2 in the Liquid with m1 = 6.648 mol·kg−1 at the Temperatures from (298.15 to 318.15) K and the System Pressure of 122.7 kPa T/K
298.15
303.15
308.15
313.15
318.15
103H′ ΔG/kJ·mol−1 ΔH/kJ·mol−1 ΔS/J·mol−1·K−1
1.42 ± 0.04 −16.3
1.93 ± 0.06 −15.8
2.34 ± 0.13 −15.5 −32.1 −53.3
2.46 ± 0.07 −15.6
3.46 ± 0.21 −15.0
3.5. Spectral Analysis. This section focuses on the interaction between molecules in the urea + ethylene glycol + sulfur dioxide system. Figure 9 displays the results of Fourier
Figure 10. FTIR spectra of the solution with m1 = 6.648 mol·kg−1 before and after absorbing SO2: blue line, before absorbing SO2; black line, after absorbing SO2; red line, after desorbing SO2.
namely, 1147 and 1325 cm−1. According to the refs 38, they are the symmetrical and asymmetrical stretching of SO band, respectively, and disappear in the spectrograms after desorption. These IR results reveal the absorbed SO2 molecules maintain itself intrinsic structure in the mixture solution and can be desorbed from solution, indicating there are a weak interaction between SO2 and the mixture. Except for the results above, the stretching vibration peak of CO (1666 cm−1) has no change after absorbing SO2, indicating the interaction between urea and SO2 is not located at the O atom of carbanyl group. The changes like 881 cm−1 (CH2 rocking band), 1042 cm−1 (C−C stretching band), 1084 cm−1 (C−O stretching band), and about 3400 cm−1 (hydrogen bond), may result from the hydrogen bond of urea with EG, which could be disrupted by the absorption of SO2. To further investigate the mechanism between SO2, urea, and EG, 1H NMR spectra experiments were also performed at normal temperature. The results are given in Table 9 and
Figure 9. FTIR spectra of EG, urea, and mixture of urea + EG with m1 = 6.648 mol·kg−1.
transform infrared spectroscopy (FTIR) on urea, ethylene glycol, and the solution with m1 = 6.648 mol·kg−1. It is seen that the absorption peaks of urea after mixed with EG are greatly different. Obviously, the characteristic peaks of urea molecular structure at (1153, 3344, and 3443) cm−1, assigned to the rocking, symmetrical stretching and asymmetrical stretching of NH34 respectively, disappear after mixed. For the peak of 1464 cm−1 attributed to CN asymmetrical stretching, its intensity turns weaker and can hardly be observed. In addition, the stretching peak of carbonyl group in urea shifts from (1682 to 1666) cm−1, which may result from the new hydrogen bond formed between the oxygen atom of carbonyl group or hydrogen atom of amine group in urea and the hydroxyl groups in EG, that is, OHglycol···Ocarbonyl and NHamide···Oglycol. In Figure 9, it is also found that the absorption peaks of C−C (1040 cm−1) and C−O (1084 cm−1), which are attributed to their stretching vibration35 in EG, have a blue shift of 2 cm−1 after mixing. This means that the addition of urea destroys first the original hydrogen bond of head-to-tail ligation between EG molecules,36,37 and then the new hydrogen bond between EG and urea forms. Compared to OHglycol···Oglycol (the original hydrogen bond between EG molecules), the electron cloud density of O in the new hydrogen bond (NHamide··· Oglycol) grows up. Thus, the above two peaks shift toward a higher wavenumber. For the OH band (a wide and strong peak 3424 cm−1 in EG), it shifts to 3408 cm−1 in the mixture solution, resulting from the new hydrogen bond. Figure 10 gives the information on the changes of IR spectrograms of the binary mixture with m1 = 6.648 mol·kg−1 before and after absorbing SO2. Clearly, there are two new absorption peaks in the spectrogram after absorbing SO2,
Table 9. Chemical Shifts and Attribution of 1H in Urea, EG, Urea + EG, Urea + SO2, and Urea + EG + SO2 in CD3SOCD3 Solutions system
NH
urea EG urea + EG urea + EG + SO2 after desorbing SO2
5.422 5.469 5.488 5.461
OH
CH2
4.418 4.467 4.479 4.477
3.402 3.401 3.405 3.400
Figure 11. In Table 9, the proton chemical shifts of NH in urea and OH in EG shift toward the lower field when mixing two solvents. This phenomenon is caused by the formed hydrogen bond (OHglycol···Ocarbonyl and NHamide···Oglycol) in the mixture solution which lead to the two protons exposed and their deshielding effect increase. In addition, the values of all proton chemical shifts move downfield after absorbing SO2 and recover G
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Figure 11. 1H NMR absorption spectra of urea (a), EG (b), urea + EG (c), urea + SO2 (d), and urea + EG + SO2 (e).
Figure 12a and b illustrate the UV spectrograms of urea + EG and urea + EG +SO2 systems, respectively. As shown in Figure 12a, the maximum absorption peak presents a red shift with the concentration of urea increase in binary mixture, and its intensity enlarges simultaneously. Owing to a big p−π conjugated system in urea molecule,39,40 the cause of red shift stems from the hydrogen bond between urea and EG which can reduce the electron cloud density of π system. Based on the literature,41 there are two electron transition bands, namely, π → π* (200 to 230 nm) and n → π* (280 to 310 nm). In Figure 12b, the two absorption bands change with the change of SO2 concentration in the ternary system. Specifically, with SO2 concentration growth, the π → π* band shows a red shift, and its intensity ascends; meanwhile, the n → π* band changes in intensity only. These results show that the absorbed SO2 molecules can also stay in the mixture solution in the form of free state except for the interaction with urea or EG.
after desorbing SO2 except for the proton of hydroxyl groups. From these results the following conclusion may be deduced. First, the hydrogen bond NHamide···Oglycol between urea and EG is subjected to break by the incoming SO2. Second, S atom with empty orbit in SO2 is attracted by N atom with lone electron pair in urea, resulting in the proton of NH2 more exposed and the deshielding effect increase. At the same time, the hydrogen bond OHglycol···OSO between EG and SO2 forms and brings about the proton chemical shift of OH moving downfield. Finally, because the energy of charge-transfer interaction between urea and SO2 are weaker than the hydrogen bond between EG and SO2, SO2 molecules absorbed by chargetransfer interaction can be easily desorbed and that absorbed by the bond OHglycol···OSO can be hardly escaped at the experimental condition. Based on the desorption result, SO2 is captured by the weak charge-transfer interaction mainly. H
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REFERENCES
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Figure 12. (a) Ultraviolet absorption spectra of urea + EG (EG as reference solution). (b) Ultraviolet absorption spectra of urea + EG + SO2 (EG + urea as reference solution).
4. CONCLUSION The solubility of dilute SO2 in a binary mixture of urea + EG, and the change rules of the density and viscosity of binary mixture with the increasing of urea molality and measurement temperature were studied in the paper. The interaction mechanism was analyzed using modern spectrum methods. The results demonstrated that the addition of urea can improve the capability of EG absorbing SO2. In spite of the temperature of 318.15 K, the solution with m1 = 6.648 mol·kg−1 has good performance for SO2 absorption as well as low viscosity. In addition, SO2 is captured by physical processes. The physical absorption is included two kinds of force, that is, the chargetransfer interaction (S atom with empty orbit in SO2 and N atom with lone electron pair) and the hydrogen bond (OHglycol···OSO). According to the desorption result, more than 90 % SO2 can be released from the absorption solution. In conclusion, the solution with m1 = 6.648 mol·kg−1 is a promising candidate since it can absorbed SO2 by physical process as well as weak chemical interaction and be reused through desorption SO2 at a higher temperature.
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AUTHOR INFORMATION
Corresponding Author
*Tel.: +86-010-62751529; fax: +86-010-62670662. E-mail:
[email protected] (X. H. Wei). Funding
This work was supported by Boyuan Hengsheng HighTechnology Co., Ltd., Beijing, China. Notes
The authors declare no competing financial interest. I
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