ARTICLE pubs.acs.org/IECR
Composition Analysis of CO2NH3H2O System Based on Raman Spectra Qing Zhao, Shujuan Wang,* Feng Qin, and Changhe Chen Department of Thermal Engineering, Key Laboratory for Thermal Science and Power Engineering of Ministry of Education, Tsinghua University, 100084 Beijing, China ABSTRACT: We used Raman spectroscopy to systematically and comprehensively measure the concentrations of the main ion components (NH2COO, HCO3, and CO32) in the CO2NH3H2O system for an initial total ammonia concentration [A] = 0.69, 0.88, 1.09, 1.34, 1.60, and 2.10 mol 3 L1 and CO2 loading [C]/[A] = 0.18, 0.35, 0.43, 0.50, and 0.67, respectively. The experimental data were processed using two existing theoretical methods, and the analysis on results allowed us to propose a number of improvements based on these theoretical methods. Experimental results show that upon increasing [C]/[A], the ratios of [NH2COO] and [CO32] to total carbon decrease linearly, whereas the ratio of [HCO3] to total carbon increases linearly. The reaction process can be divided into three steps. Considering the impacts of experiment errors, [A] has little effect on the relative composition of the system in the range we studied. Results calculated with a revised theoretical method are in better agreement with the experimental data.
’ INTRODUCTION Because of the growing consensus that climate change is due to the greenhouse effect, which is caused by the introduction of huge quantities of CO2 into the atmosphere, the question of how to effectively limit and reduce greenhouse-gas emissions has attracted wide attention as a high priority. The CO2 emitted from coal-fired power plants is a main source of greenhouse gases, and to reduce the concentration of CO2 in coal-fired power-plant flue gas is very important and so has great significance for environmental protection and sustainable development. To address this issue, researchers1,2 have proposed aqueous ammonia to absorb CO2 from power-plant flue gas. The majority of such experimental studies use a mixed gas of CO2 and N2 to simulate the flue gas after desulfurization and denitrification and react this mixed gas with aqueous ammonia as a CO2NH3H2O reaction system. The initial total ammonia concentration ([A]) in the liquid phase gives the total absorption capacity, and the concentration ratio of the total carbon to the initial total ammonia ([C]/[A]) indicates CO2 loading. However, most research on this reaction system is based on the carbonation process in the chemical industry, such as that for the synthesis of NH4HCO3 in fertilizer plants or for the production of soda and ammonium chloride in soda plants, etc. However, conditions (pressure, temperature, and reactant concentration) of these chemicalindustry processes are very different from that in power plants.3 Therefore, it is necessary to study the CO2NH3H2O system systematically and comprehensively for the specific circumstances of power plants. In particular, to optimize the reaction process with regard to [C]/[A] and [A], it is very important to specify the products that are formed in aqueous solution during the reaction between ammonia and carbon dioxide under different experimental conditions. In the CO2NH3H2O system, there may be CO2, H2O, NH3 3 H2O, H2CO3, NH4þ, Hþ, OH, NH2COO, HCO3, and CO32 in the solution.4,5 Because this system is alkaline, r 2011 American Chemical Society
concentrations of CO2, H2CO3, and Hþ are negligible.6 According to mass conservation of nitrogen and carbon, [A] and [C] can be defined as follows: initial total ammonia concentration : ½A ¼ ½NH3 þ ½NH4 þ
þ ½NH2 COO
ð1Þ
total carbon concentration : ½C ¼ ½HCO3 þ ½CO3 2 þ ½NH2 COO ð2Þ Therefore, the principle cation is NH4þ, and our experimental and theoretical study focuses on changes in concentration of the three carbon-containing anions (CCAs; NH2COO, HCO3, and CO32) in the liquid phase. Although there are many theoretical methods for calculating concentrations of such a multicomponent gasliquid two-phase reaction system, all have inherent problems. To date, no satisfactory theoretical method has been published to evaluate the relative concentration of species that occur in the CO2NH3H2O system.7,8 Thus, this is not a reliable way to investigate the distribution of CCAs in the CO2NH3H2O system. We introduced two main methods herein and provided detailed discussions, calculations, and comparisons in the body of this paper. Both nuclear magnetic resonance (NMR) and Raman spectroscopy can be used simultaneously to detect all three CCAs. Holmes et al.9 measured CCAs at 293 K by NMR in a series of Received: May 4, 2010 Accepted: February 27, 2011 Revised: January 25, 2011 Published: March 15, 2011 5316
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Industrial & Engineering Chemistry Research CO2NH3H2O systems with [A] and [C]/[A] ranging discretely from 0.69 to 8.95 mol 3 L1 and 0.35 to 0.72, respectively, but no systematic measurement. Mani et al.8 used NMR to determine the relationship between the concentration ratio of each CCA to [C] and the pH of the solution using aqueous ammonia (2.5 mol 3 L1) reacting with a mixed gas of CO2 (10 vol %) and N2 at 293 K, also to investigate the relationship between the concentration ratio of each CCA to [C] and [NH3]/[NH4þ] of the solution using a series of 0.24 L aqueous ammonia (0.85, 2.5, 5.08, 7.5, and 10 mol 3 L1, respectively), reacting with ca. 0.1 mol of CO2 (equivalently, [C]/[A] ≈ 0.490, 0.167, 0.082, 0.056, and 0.042). Because [A] and [C]/[A] are both variable, the effect of [A] or [C]/[A] exerted on ion concentrations of CCAs could not be known. Using Raman spectroscopy, Wen and Brooker7 measured the concentration of each CCA in ammonium carbonate solutions (0.65383.451 mol 3 kg1) at 295 K and in the same solutions (2.9047 mol 3 kg1) at 295373 K; then they investigated the effects of temperature and [A] exerted on concentrations of CCAs. It can be calculated that when [C]/[A] (≈0.5) was unchanged and [A] increased from 0.6538 to 3.451 mol 3 kg1, the absolute change values of the concentration ratio of CCAs to [C] were þ6.86, 5.91, and 0.96%, respectively. For the process of absorbing a mixed gas of CO2 (12.81 vol %) and N2 by aqueous ammonia (13 wt % or, equivalently, 7.22 mol 3 L1), Kim et al.10 measured CCA concentrations by Raman spectroscopy at different absorption phases (37 h) at 298 K. From related information given in this paper, it could be inferred that when [C]/[A] < X1, [NH2COO] > [CO32] > [HCO3]; when X1 < [C]/[A] < X2, [NH2COO] > [HCO3] > [CO32]; and when [C]/[A] > X2, [HCO3] > [NH2COO] > [CO32]. In the absence of more data, we can only know that X1 is slightly greater than 0.35 and X2 is slightly smaller than 0.45. In summary, existing quantitative studies on CCAs using NMR or Raman spectroscopy are only for the CO2NH3H2O system with limited [A] and [C]/[A]. Therefore, a more systematic and comprehensive study of CCAs in the CO2NH3H2O system with different [A] and [C]/[A] (01) is necessary. Because Raman spectroscopy offers the advantage of rapid scan speed, and is preferable for aqueous solutions, we chose it to investigate the distribution of CCAs at room temperature, with different [A] and during the reaction progress. By comparing the experimental results with the theoretical results, we proposed an improved theoretical treatment.
’ EXPERIMENTAL METHODS The initial total ammonia concentration [A] is selected according to the solubility of ammonium salts8 (NH4HCO3, 2.78 mol 3 L1; (NH4)2CO3, 3.33 mol 3 L1) and the limitation of the Raman spectrometer. The CO2 loading [C]/[A] is selected to simulate the entire process of CO2 absorption by aqueous ammonia. When the total amount of ammonia and carbon are equivalent, the CO2NH3H2O system formed though CO2 absorption by aqueous ammonia and preparation by interrelated chemicals are all the same.11 According to this basic law, all of the solutions we used in the study were prepared by ammonium carbonate, ammonium bicarbonate, and 25 wt % ammonia solution (AR grade, Beijing Modern Eastern Fine Chemicals Co., Ltd.). After chemicals were weighed and they were dissolved with deionized water, solutions were stored in volumetric flasks
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Figure 1. Raman spectrogram (T = 25 °C).
(250 mL) for 24 h standing culture to reach chemical equilibrium. As NH4HCO3 is disproportionally dissolved (solution composition phase is inconsistent with solid composition) between 20 and 50 °C at 1 atm,12 all of the solutions were titrated by an automatic potentiometric titrator (809 Titrando, Metrohm, Zurich, Switzerland), and averages of titration values were treated as true values of [A] (0.69, 0.88, 1.09, 1.34, 1.60, and 2.10 mol 3 L1) and [C]/[A] (0.18, 0.35, 0.43, 0.50, and 0.67) for subsequent analysis. Absolute errors of titration values for [C] and [A] are both 0.019 mol 3 L1. At the same time, we sampled some solution from the volumetric flasks and sealed it in capillary glass tubes and then measured the solution with a confocal Raman spectrometer (Renishaw, Wotton under Edge, UK RM2000-type). Instrument parameters were as follows: laser wavelength, 632.8 nm; spectral resolution, 1 cm1; scan time, 30 s; accumulated times, 10. Experimental data were processed with professional software GRAMS/32. The experiments were performed at 25 °C and 1 atm. The Raman quantitative measurement formula can be expressed as Ii ¼ Ji ci
ð3Þ
where Ii is the relative scattering intensity, Ji is the molar scattering intensity for each molecule and is characteristic of the designated band for the given measurement conditions and medium, and ci is the concentration of the chosen molecule. Although the concentration of the chosen molecule can be calculated using eq 3, it is in practice very difficult to obtain a linear relationship between Ii and ci experimentally because the Raman intensity also depends on many instrumental and sample factors, such as the stability of the light source, the monochromator slit width, the size of the sample pool, self-absorption in the sample, refractive index changes due to different sample concentrations, and background noise due to the solvent. Some of these factors are difficult to control, such as samples that may partially absorb scattered light, especially for the method of resonance-enhanced Raman spectroscopy. Therefore, a substance (NaClO4 3 H2O) is added when preparing the solutions to serve as an internal measurement standard to eliminate such factors. To ensure the reliability and accuracy of experimental data, the whole process of experiments (weighing, dissolution, titration, Raman spectra, and data processing) were repeated 35 times in every set of condition and averages of multigroup data were taken as a basis for the following analysis. The Raman spectrogram of Figure 1 shows the relationship between relative scattering intensity and frequency shift which relates to the type of molecule or ion. According to the studies of Wen and Brooker7 and Kim et al.,10 the frequency shifts for 5317
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ClO4, HCO3, NH2COO, and CO32 are 935, 1017,1034, and 1065 cm1, respectively. The experiment involves the three steps given as follows: (1) Determination of molar scattering intensity J of HCO3/ CO32. A number of pure solutions of Na2CO3/NaHCO3 with different concentrations were prepared with NaClO4 3 H2O serving as the internal standard. Because ClO4 is stable in aqueous solution, its concentration, denoted as c935 (mol 3 L1), is equivalent to the mass-calculated value. The concentrations of HCO3, NH2COO, and CO32 are denoted as c1017, c1034, and c1065 (mol 3 L1), respectively. The integrals of the Raman peaks for each component were obtained using professional software (GRAMS/32) analysis tools and are denoted as A935, A1017, A1034, and A1065, respectively. Because the integral A of the Raman peak is proportional to the relative scattering intensity I in the same spectral range, we can infer that Ii Ai ¼ ð4Þ I935 A935 where i can refer to HCO3, NH2COO, or CO32. Because Na2CO3 and NaHCO3 can ionize completely, the HCO3 and CO32 concentrations can be calculated on the basis of their mass. The molar scattering intensity J of HCO3 and CO32 can be obtained from the following formula: Ii
¼ Ji
I935
ci
ð5Þ
c935
(2) Calculation of molar scattering intensity J of NH2COO. Because solid NH2COONa can react with water, it is not possible to obtain the corresponding pure solution. Therefore, the molar scattering intensity J of NH2COO must be calculated indirectly. To do this, the first step is to prepare (NH4)2CO3 solutions with similar NaClO4 3 H2O concentration, as in the previous step, and then to measure the Raman relative scattering intensity I of CCAs. The concentrations of HCO3 and CO32 can be derived by making use of the molar scattering intensity calculated in the previous step. cHCO3 ðaqÞ ¼
HCO3
J
standard deviation
0.188
0.0042
NH2COO
0.224
0.0390
CO32
0.302
0.0145
cCO3 2 ðaqÞ ¼
ICO3 2 ðaqÞ JCO3 2 ðaqÞ
ð6Þ
Because the concentrations of CO2 and H2CO3 are very low and can therefore be neglected, the concentration of NH2COO can be calculated by using the law of carbon conservation, cNH2 COO ðaqÞ ¼ cðNH4 Þ2 CO3 ðaqÞ cHCO3 ðaqÞ cCO3 2 ðaqÞ ð7Þ Finally, the molar scattering intensity J of NH2COO can be calculated using the following formula: JNH2 COO ðaqÞ ¼
INH2 COO ðaqÞ cNH2 COO ðaqÞ
ð8Þ
(3) Calculation of the concentration of each component by making use of J. Equations 4 and 5 are used for the calculation. This study focuses on the concentration ratio of ion i to [C], as given in eq 9, so the exact mass of NaClO4 3 H2O is not needed. ci Ii =Ji Ai =Ji ¼ ¼ ½C Ii =Ji Ai =Ji
Table 1. Molar Scattering Intensity J for Each Component and the Corresponding Standard Deviations (T = 25°C) component
IHCO3 ðaqÞ , JHCO3 ðaqÞ
∑i
∑i
ð9Þ
1. Experimental Results. Table 1 shows the experimentally determined values for the molar scattering intensity J for each component and the corresponding standard deviations.
Figure 2. Experimental concentration ratios of CCAs to [C] as a function of [C]/[A] ([A] = 1.34 mol 3 L1, T = 25 °C). 5318
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Figure 3. Experimental concentration ratios of CCAs to [C] as a function of [A] ([C]/[A] = 0.67, T = 25 °C).
The result of the measurements, which are shown in Figure 2, indicate that, with increasing [C]/[A], the ratio of [HCO3] to [C] increases linearly, whereas those of [NH2COO] and [CO32] decrease linearly. The curves intersect at points (X1, Y1), which is where [HCO3] = [CO32], and (X2,Y2), which is where [HCO3] = [NH2COO]. The result is essentially similar to that obtained by the research of Mani.8 Errors may arise from weighing, dissolution, Raman light intensity measurements, and data processing. Absolute errors of concentration ratios of CCAs to [C] are 9.0, 7.5, and 2.6%, respectively. The measurement results and error limits of the CO 2 NH3H2O system ([A] = 0.692.10 mol 3 L1, [C]/[A] = 0.67) are shown in Figure 3. By fitting to the experimental data, we find ostensibly that upon slightly increasing [A], the ratio of [NH2 COO] to [C] increases, whereas the ratios of [HCO3] and [CO32] decrease. However, the impact of errors on the relationship between the concentration ratios of the three CCAs to [C] and [A] could be positive or negative or have no obvious impact. Such errors may lead to results similar to that obtained in the research of Wen and Brooker7 with 0.65383.451 mol 3 kg1 ammonium carbonate solutions. However, the linear correlation coefficients (R) are very small, so under the experimental conditions used for the present study, we assume that [A] has no effect on the concentration ratios of the three CCAs to [C]. To determine more accurate trends, more precise measurements are required. The data with the same [C]/[A] were averaged for different values of [A], and the results were fitted to a line. The resulting slopes and intercepts are shown in Table 2. On the basis of the analysis of the experimental results, the entire low-concentration ammoniaCO2 reaction process can be divided into the following steps: (1) When [C]/[A] < X1, [NH2COO] > [CO32] > [HCO3]; (2) when X1 < [C]/[A] < X2, [NH2COO] > [HCO3] > [CO32]; (3) when [C]/[A] > X2, [HCO3] > [NH2COO] > [CO32].
Table 2. Parameters Used To Fit to the Data of Figure 2 (T = 25 °C) HCO3
NH2COO
CO32
slope
1.14
0.63
0.50
intercept
0.01
0.58
0.41
R2
0.994
0.997
0.987
Upon changing [A], the position of the intersection point remains basically unchanged. The intersection point of HCO3 and CO32 occurs at X1 = 0.240, Y1 = 28.4%, and the intersection point of HCO3 and NH2COO occurs at X2 = 0.323, Y2 = 37.9%. 2. Theoretical Calculations. The current methods to calculate the concentrations of each component in the CO2NH3H2O system are based on mass conservation, charge conservation, and other basic laws, combined with the chemical reaction equilibrium constant K. According to types of K, there are two main calculation methods, which are discussed below. a. Experimental Equilibrium Constant Method. With this method, equilibrium constants are obtained by experiments, including the concentration equilibrium constant, the equilibrium constant calculated from experimentally measured activity factors, and the mixed equilibrium constant used in analytical chemistry. The constant can be used to directly calculate ionic concentration, but it varies with temperature, pressure, and composition, so it is not universal. Van Krevelen et al.13 measured the partial vapor pressures of CO2 and NH3 in the CO2NH3H2O system (2060 °C, [A] = 0.492 mol 3 L1, [C]/[A] = 0.190.67) and calculated the concentrations of CCAs, and the experimental equilibrium constants of reactions 10 and 11 were derived by fitting to these data. Zheng et al.6 used the same method and the same values of K to calculate the concentration of each component in the CO2NH3H2O system (2060 °C, [A] = 213 mol 3 L1, 5319
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Figure 4. Calculated concentration ratios of CCAs to [C] as a function [C]/[A]. The equilibrium constants were experimentally determined (T = 25 °C).
[C]/[A] = 0.150.45). NH3 þ HCO 3 T NH2 COO þ H2 O
ð10Þ
2 þ NH3 þ HCO 3 T CO3 þ NH4
ð11Þ
½CO3 2 ¼
Accordingly, the method is as follows: (1) mass conservation of nitrogen and carbon, eqs 1 and 2; (2) charge conservation, þ
½NH4 ¼ ½HCO3 þ 2½CO3 þ ½NH2 COO 2
equilibrium constant of reaction 10: ½NH2 COO ½NH3 ½HCO3
ð13Þ
equilibrium constant of reaction 11: K2 ¼
½NH4 þ ½CO3 2 ½NH3 ½HCO3
ð14Þ
Because the amount of CO2 dissolved in the CO2NH3H2O system but not yet reacted is very low,6 and the calculation under the conditions in this experiment shows that [CO2] in the liquid phase is no more than 1% of the total carbon amount, it can be ignored. From eqs 1 to 2 and eqs 12 to 14, we calculated the following expressions (eqs 1519) for the concentrations of each component: ð½A 2½CÞðK1 þ ðK2 =½CÞÞ þ 1 ½HCO3 ¼ 2K1 þ ð2K2 =½CÞ qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ½ð½A 2½CÞðK1 þ ðK2 =½CÞÞ þ 12 þ 4½CðK1 þ ðK2 =½CÞ
þ
2K1 þ ð2K2 =½CÞ
ð15Þ ½NH3 ¼ ½HCO3 þ ½A 2½C
qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ½C2 þ 4K2 ½HCO3 ½NH3 2
ð17Þ
½NH4 þ ¼ ½C þ ½CO3 2
ð18Þ
½NH2 COO ¼ ½C ½HCO3 ½CO3 2
ð19Þ
ð12Þ
(3) chemical reactions,
K1 ¼
½C þ
ð16Þ
b. Standard Equilibrium Constant Method. In this method, the standard equilibrium constant involves only temperature. The constant can be calculated from thermodynamic data or measured by, for example, electrochemical methods. However, knowledge of the activity coefficients is needed to calculate the concentrations of ions. Unfortunately, the activity coefficient of single ionic species cannot be accurately determined by strict thermodynamic methods, so some nonthermodynamic assumptions and mathematical approximations must be introduced, which lead to errors. In addition, two different standard equilibrium constants for reaction 10 are found in the literature, both with unknown origins. Edwards et al.14 and Chen and Han15 used the K = 3.1 from Faurholt16 and from personal communications in 1977 with Harris and Mason. Kawazuishi and Prausnitz17 and Xu and Tang18 used K = 2.2, which was based on a communication in 1982 with Mason. The discrepancy between these two K values makes it difficult to determine the concentrations of system components. In this paper, we used the data from Edwards et al.14 to establish a simulation model for a multicomponent gasliquid two-phase reaction system.19
’ RESULTS AND DISCUSSION 1. Theoretical Results. The results of the two types of theoretical calculations are shown in Figures 4 and 5, using experimental and standard equilibrium constant method, respectively, which show that, with increasing CO2 absorption, the ratio of [HCO3] to [C] increases linearly, whereas the ratios of 5320
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Figure 5. Calculated concentration ratios of CCAs to [C] as a function of [C]/[A], derived using standard equilibrium constants (T = 25 °C).
Figure 6. Plot of log K1 and log K2 vs T1.
[NH2COO] and [CO32] to [C] decrease linearly, in agreement with the experimental results. When [A] increases from 0.69 to 2.10 mol 3 L1, the ratios of [HCO3] and [CO32] to [C] decrease, whereas that of [NH2COO] increases. In Figure 4, the curves for [CO32-] and [HCO3] intersect at the same ratio [C]/[A] ≈ 0.25 for both values of [A]; and the curves for [NH2COO] and [HCO3] intersect at a ratio [C]/[A] of about 0.3 and about 0.5 for [A] = 0.6 mol 3 L1 and [A] = 2 mol 3 L1, respectively. In Figure 5, the ratio of [CO32] to [C] is the lowest over the entire region measured. Upon changing [A] from 0.69 to 2.10 mol 3 L1, the curves for [HCO3] and [NH2COO] intersect at a [C]/[A] ratio of about 0.25 and about 0.4, respectively.
Theoretical results show that [A] has a great impact on the distribution of ions, which is in contrast to the experimental results. 2. Improvement of Theoretical Calculation. The second calculation method involves many complicated parameters such as activity coefficients and so is relatively difficult to modify. Therefore, in this study, we restrict ourselves to improving only the first calculation method. The experimental equilibrium constants K1 and K2 in the first method are not related to [A] and [C]/[A] but are only functions of temperature, and this relationship is plotted in Figure 6. Wen and Brooker7 measured the amount of each CCAs at 22 °C in solutions of (NH4)2CO3 and assumed [NH3] ≈ [HCO3], or 5321
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Table 3. Related Data Used To Fit to the Carves of Figure 7 ([C]/[A] = 0.5, T = 25 °C) [A], mol 3 L1
[HCO3], mol 3 L1
[NH2COO], mol 3 L1
[CO32-], mol 3 L1
K1
K2
1.31
0.38
0.17
0.10
1.21
0.52
1.99
0.58
0.27
0.15
0.80
0.52
2.70
0.78
0.36
0.21
0.59
0.52
3.43
0.10
0.46
0.26
0.46
0.52
4.18
1.21
0.56
0.32
0.38
0.52
4.98
1.44
0.66
0.38
0.32
0.52
5.81
1.69
0.78
0.44
0.27
0.52
6.90
2.00
0.92
0.53
0.23
0.52
Figure 7. Experimental equilibrium constants as a function of [A] (T = 25 °C).
Figure 8. Concentration ratios of CCAs to [C] as a function of [C]/[A], derived using the improved experimental equilibrium constants (T = 25 °C). 5322
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Figure 9. Concentration ratio of [HCO3] to [C] as a function of [C]/[A] derived using different methods (see figure legend, T = 25 °C).
Figure 10. Concentration ratio of [NH2COO] to [C] as a function of [C]/[A] derived using different methods (see figure legend, T = 25 °C).
[C]/[A] = 0.5. The reaction equilibrium constants K1 obtained in that study are strongly associated with [A]. For a more convenient comparison, the experimental data (measured at 25 °C) were refitted with the same condition of [C]/[A] = 0.5; the related data are in Table 3. The recalculated values of K1 and K2 were plotted versus [A] in Figure 7 along with the original data from Wen and Brooker.7 The similar relationships constitute strong evidence that K1 and [A] are related. That the values of K1 and K2 derived from experimental data are larger than those from Wen and Brooker7 is tentatively attributed to experimental conditions, and in particular to temperature differences. After improvement, [A] does not affect the concentration ratios of various CCAs to [C]. The results of the calculation are shown in Figure 8.
To facilitate a comprehensive comparison of different theoretical calculation methods, Figures 911 show average values of the ion concentration ratio for the same [C]/[A], but with different [A] ([A] = 0.69, 0.88, 1.09, 1.34, 1.60, and 2.10 mol 3 L1) calculated by method 1, method 2, and the improved method 1, respectively, together with average results of experimental data and the experimental error limits. As is shown, the estimates of [HCO3] to [C] based on methods 1 and 2 are accurate, but those for [NH2COO] and [CO32] exceed the experimental error limits, so both theoretical methods 1 and 2 fail to simultaneously give concentrations of all. Only estimates of improved method 1 do not exceed the experimental error limits for all CCAs. 5323
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Figure 11. Concentration ratio of [CO32] to [C] as a function of [C]/[A] derived using different methods (see figure legend, T = 25 °C).
Table 4. Average Relative Errors between Improved Method 1 and Methods 1 and 2 (T = 25 °C) compared with method 1
be divided into the following steps: ½C=½A < X1 ,
compared with method 2
X1 < ½C=½A < X2 ,
[C]/[A] HCO3 NH2COO CO32 HCO3 NH2COO CO32 0.18
2.65
14.09
0.35
7.09
6.72
21.21
0.43
8.57
2.34
20.46
0.50
9.01
9.36
19.85
0.67
5.15
20.35
8.09
½NH2 COO > ½CO3 2 > ½HCO3 ;
18.93 36.58
4.65
28.84
4.36
16.47
34.85
0.29
19.24
32.01
1.50
19.80
26.30
0.90
17.54
0.33
For accurate comparison, we calculated the average relative errors between the improved method 1, method 1, method 2, and experimental values on the basis of the same ratio [C]/[A] (with [A] = 0.69, 0.88, 1.09, 1.34, 1.60, and 2.10 mol 3 L1). Using the following formula, the results are calculated and are shown in Table 4. improved method 1 experiment data average relative error ¼ experiment data
method 1=2 experiment data experiment data
ð20Þ
½C=½A > X2 ,
½NH2 COO > ½HCO3 > ½CO3 2 ;
½HCO3 > ½NH2 COO > ½CO3 2
In these steps, X1 is the [C]/[A] ratio where [HCO3] = [CO32] = Y1 and X2 is the [C]/[A] ratio where [HCO3] = [NH2COO] = Y2. (2) In this experiment, in the range of 0.692.10 mol 3 L1, considering impacts of experiment errors, [A] has little effect on the relative concentrations of HCO3, NH2COO, and CO32. Upon changing [A], the position of the intersection point remains essentially unchanged. More specifically, when X1 = 0.240, Y1 = 28.4%, and when X2 = 0.323, Y2 = 37.9%. (3) When [C]/[A] is low, the deviations between theoretical results and experimental data are larger and more sensitive to variations in [A]. (4) Upon improving the theoretical method 1, the average relative errors decrease, and [A] does not affect the relative concentrations of HCO3, NH2COO, or CO32. Furthermore, this result agrees with experimental results.
’ AUTHOR INFORMATION Corresponding Author
Except for individual cases, the improved method 1 reduces the average relative errors—the most significant reduction is 36.6%. Furthermore, [A] does not influence the calculated results, which is consistent with the experimental data.
’ CONCLUSIONS (1) Experimental results show that upon increasing [C]/[A], the ratios of [NH2COO] and [CO32] to [C] decrease linearly, whereas the ratio of [HCO3] to [C] increases linearly. The process of CO2 absorption by low-concentration ammonia can
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’ ACKNOWLEDGMENT We gratefully acknowledge support for Project 50876051 from the National Natural Science Foundation of China. ’ NOMENCLATURE [A] = initial total ammonia concentration (mol 3 L1) Ai = integral of the Raman peak of ion i 5324
dx.doi.org/10.1021/ie1010178 |Ind. Eng. Chem. Res. 2011, 50, 5316–5325
Industrial & Engineering Chemistry Research [C] = total carbon concentration (mol 3 L1) [C]/[A] = CO2 loading ci = concentration of ion i (mol 3 L1) Ii = Raman relative scattering intensity of ion i Ji = molecular characteristic constant of ion i K = chemical reaction equilibrium constant T = temperature (X1,Y1) = coordinates of intersection point [HCO3] and [CO32] when plotted as a function of [C]/[A] (X2,Y2) = coordinates of intersection point of [HCO3] and [NH2COO] when plotted as a function of [C]/[A]
ARTICLE
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