with Sarcosine Salt in Aqueous Solution - American Chemical Society

Sep 19, 2012 - ABSTRACT: Aqueous sarcosine salts are fast carbon dioxide (CO2) absorbents suitable for use in postcombustion CO2 capture in coal-fired...
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Kinetics of the Reversible Reaction of CO2(aq) and HCO3− with Sarcosine Salt in Aqueous Solution Qunyang Xiang,†,‡ Mengxiang Fang,† Hai Yu,*,‡ and Marcel Maeder*,§ †

State Key Laboratory for Clean Energy Utilization, Zhejiang University, Hangzhou, 310027, P. R. China CSIRO Energy Centre, 10 Murray Dwyer Circuit, Mayfield West, NSW 2304, Australia § Department of Chemistry, The University of Newcastle, Callaghan, NSW 2308, Australia ‡

S Supporting Information *

ABSTRACT: Aqueous sarcosine salts are fast carbon dioxide (CO2) absorbents suitable for use in postcombustion CO2 capture in coal-fired power plants. We have developed a detailed reaction scheme including all the reactions in the sarcosine−CO2−water system. All unknown rate and equilibrium constants were obtained by global data fitting. We investigated the temperature-dependent rate and equilibrium constants of the reaction between aqueous CO2 and sarcosine using stopped-flow spectrophotometry, by following the pH changes over the wavelength range 400−700 nm via coupling to pH indicators. The corresponding rate and equilibrium constants ranged from 15.0 to 45.0 °C and were analyzed in terms of Arrhenius, Eyring, and van’t Hoff relationships. The rate constant for the reaction between CO2 and sarcosine to form the carbamate at 25.0 °C is 18.6(6) × 103 M−1 s−1, which is very high for an acyclic amine; its activation enthalpy is 59(1) kJ mol−1 and the entropy is 33(4) J mol−1 K−1. In addition, we investigated the slow reaction between bicarbonate and sarcosine using 1H nuclear magnetic resonance spectroscopy and report the corresponding rate and equilibrium constants at 25.0 °C. This rate constant is 5.9 × 10−3 M−1 s−1.

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INTRODUCTION Carbon dioxide (CO2) emissions from coal-fired power plants and other sources are significant contributors to global climate change.1 The capture and sequestration of CO2 is one way to significantly reduce CO2 emissions and mitigate the effects of climate change. Postcombustion capture (PCC) technology using chemical absorbents is thought to be the most promising way to limit CO2 emissions from existing coal-fired power plants.2 PCC entails the separation of CO2 from the other flue gases with subsequent sequestration in geological structures. The most advanced techniques for PCC are based on the reversible absorption of CO2 in aqueous amine solutions. Gaseous CO2 is selectively absorbed into the amine solution in an absorber column at low temperature, the CO2-rich amine solution is then moved to a stripper column where a fraction of the CO2 is desorbed at a higher temperature and released as a relatively pure gas whereas the CO2-lean amine solution is recycled to the absorber. Dissolved CO2 reacts in two ways, it reacts with water and hydroxide to form carbonic acid and bicarbonate,3 and it also reacts with some amines to form the carbamate. With most amines this second reaction is much faster and thus advantageous as it reduces the reaction times for absorption and thus the size of the absorber column. To improve the efficiency of the PCC process, the investigation of the carbamate forming reaction is crucially important. A substantial © 2012 American Chemical Society

collection of amines has been investigated and it is important to expand this collection and include alternatives such as amino acid salts.4,5 Amino acid salts are proven alternative absorbents for CO2 capture.6−9 Traditional amine-based absorbents, such as monoethanolamine (MEA), are fairly volatile; they suffer from thermal and oxidative degradation, both of which can produce toxic compounds.10 In comparison, amino acid salt solutions exhibit low volatility, have good oxidative and thermal stability and are more environmentally friendly.11,12 Amino acid salts and MEA have comparable reaction rates with CO2 and are significantly faster than aqueous ammonia, which is another important robust CO2 absorbent. Compared with other amino acid salt solutions, aqueous sarcosine salts have fast kinetics with CO2.13 This is confirmed by our recent work using a wetted wall column, in which we compared the mass transfer coefficients for CO2 absorption in a number of amines and amino acid salts, including MEA, ammonia, and sodium salts of sarcosine, glycine, and taurine.14 We have found that sarcosine has a much faster CO2 absorption rate than the two other amino acid salts and MEA. Simons et al.15 used a gas−liquid contactor to study the kinetics of CO2 with sarcosine salts and obtained the rate constant for carbamate formation reaction. Van Holst et al.16 Received: June 12, 2012 Revised: September 19, 2012 Published: September 19, 2012 10276

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Figure 1. General reaction scheme including all reactions in the sarcosine−CO2−water system. Dashed arrows represent reactions for which constants are known; solid arrows represent reactions for which constants are determined in this work.

(SELBY) were used as obtained. Sarcosine was premixed with an equal molar amount of sodium hydroxide to form sodium sarcosinate (SAR−). Ultrahigh-purity Milli-Q water was boiled to remove CO2 and used to prepare all aqueous solutions. All samples and solutions were stored in an N2-purged glovebag. Stopped-Flow Spectrophotometric Study: Formation and Decomposition Reactions of Sarcosine Carbamate. The kinetics of the fast reactions between CO2(aq) and SAR− to form carbamate and the sarcosine carbamate/carbamic acid decomposition reaction were monitored on an Applied Photophysics DX-17 spectrophotometer equipped with a J&M Tidas MCS 500−3 diode-array detector. We observed the pH changes over the wavelength range 400−700 nm via coupling to pH indicators. Samples were thermostated at temperatures of 15.0, 25.0, 35.0, and 45.0 °C within ±0.1 °C; the exact temperature was recorded by a thermocouple located within the stopped-flow instrument. Each stopped-flow measurement was repeated at least four times, allowing comprehensive statistical analysis of the results, i.e., rather than relying on the estimates for the standard deviations of the fitted parameters as supplied by the fitting software, we have four independently determined values for each parameter, defining the standard deviation in a more realistic way. Carbamate formation reaction between SAR− and CO2(aq) was initiated by mixing an aqueous solution of SAR− and thymol blue (1:1 v/v) with an aqueous solution of CO2. The initial concentrations of SAR− after mixing ranged from 3.0− 16.0 mM and the initial concentrations of CO2(aq) after mixing ranged from 5.2 to 13.4 mM. All reactions were performed in the presence of 12.5 μM thymol blue indicator. Similarly, the decomposition of sarcosine carbamate at low pH was initiated by mixing aqueous solutions of sarcosine carbamate with solutions of hydrochloric acid ([H+] = 45.0− 55.0 mM after mixing). Sarcosine carbamate solutions were generated by mixing SAR− with HCO3− ([SAR−]total = 25 mM, [HCO3−]total = 50 mM) and leaving the mixed solution to reach equilibrium in a water bath at reaction temperature for at least 24 h. The composition of the solution (specifically, the concentration of carbamate in the equilibrated solutions) was determined by 1H NMR spectroscopy. Because there was a large pH change during the decomposition reaction of carbamate, a combination of 2 pH indicators (alizarin red S (50 μM) and methyl orange (25 μM)) were preadded to the hydrochloric acid to follow the reaction. 1 H NMR Studies: Kinetics of the Reaction between SAR− and Bicarbonate (Formation and Decomposition

studied the physiochemical properties of aqueous potassium sarcosinate solutions, including its density, viscosity, and physical solubility of CO2 in aqueous sarcosine solutions. Aronu et al.17 tested a mixture of sarcosine and 3-propylamine with a laboratory pilot plant and found the solution had a lower desorption heat than MEA. However, the detailed reaction mechanisms involved for the reaction between sarcosine (SAR −, CH3 NHCH2COO−, sarcosinate, deprotonated form) and CO2 are not well investigated. One reason for this is that the reaction of SAR− with CO2 is very fast and is coupled with many other reactions. Recently, Maeder et al.18−20 applied a stopped-flow spectrophotometry technique to the fast CO2 amine reaction system. They elucidated the detailed mechanisms involved in the reaction of aqueous CO2 with MEA, ammonia and several other amines, and also determined the reaction rate and equilibrium constants, which were unknown in the systems studied. This study is the extension of their previous work on the reaction system involving sodium sarcosinate (SAR−) and CO2(aq). On the basis of the studied reactions of CO2(aq) with amines, we have developed a detailed reaction scheme for SAR− and CO2(aq) in the aqueous solution. We studied the fast, reversible reaction between SAR− and CO2(aq) using a stopped-flow spectrophotometry technique from 15.0−45.0 °C by following the pH changes over the wavelength range of 400−700 nm via coupling to colored acid−base indicators. We obtained the corresponding rate and equilibrium constants over the temperature range, allowing us to determine activation and thermodynamic values. We studied the slow reaction between SAR− and bicarbonate (HCO3−), using 1H nuclear magnetic resonance (NMR) spectroscopy measurements at 25.0 °C to complete the reaction scheme. The variation of species concentrations versus time was extracted from NMR peak integration. This study determines all unknown rate and equilibrium constants for the reactions in a CO2−sarcosine−water system, providing a strong incentive for further study and modeling of sarcosine-based PCC process. To our knowledge, this is the first time that a detailed study of the reactions between SAR− and CO2 has been reported in the literature.

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EXPERIMENTAL SECTION High-purity CO2 gas (BOC), N2 (Coregas), sodium bicarbonate (BDH), sodium hydroxide (Merck), sarcosine (SigmaAldrich), hydrochloric acid (AJAX), thymol blue sodium salt (Sigma-Aldrich), alizarin red S (BDH), and methyl orange 10277

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Figure 2. (a) Absorbance at 590 nm versus time in the reaction of 8.2 mM CO2(aq) with different initial concentrations of SAR− at 25.0 °C in the presence of 12.5 μM thymol blue. (b) Absorbance at 590 nm versus time in the reaction of 8 mM SAR− with three different initial concentrations of CO2(aq) at 15.0 °C in the presence of 12.5 μM thymol blue. Solid markers are measured data and the lines are calculated profiles. k2

of Sarcosine Carbamate at High pH). The kinetics of the slow reversible reaction between SAR− and HCO3− to form carbamate and carbamate decomposition at a high pH was studied using 1H NMR on a Bruker Avance III 600, at a frequency of 600 MHz at 25.0 °C. Reactions between SAR− and HCO3− were initiated by mixing SAR− (or SAR− with OH−) solutions with HCO3− solution; the initial concentrations of SAR− and HCO3− after mixing were 50 and 100 mM, respectively. Sodium hydroxide (0 and 40 mM) was preadded to SAR− to adjust the pH. The solution was stirred for about 10 s before being transferred to the NMR instrument. Approximately 6 min after the starting time of the reaction, the spectra were recorded every 2 min in the first 40 min and every 10 min after that. The slow decomposition of sarcosine carbamate at high pH began upon mixing a pre-equilibrated solution ([carbamate]0 = 17.0 mM, [SAR−] = 33.0 mM and [HCO3−]0 = 133.0 mM, after mixing) with sodium hydroxide solution (150 mM after mixing). The carbamate decomposition reaction occurred even slower than the formation reaction. Spectra were recorded for approximately 7 h after mixing. A glass insert containing TSP (3-(trimethylsilyl)propionic acid-d4, sodium salts) as a reference in D2O acted as the locking agent. It was introduced into the NMR tube for all samples before they were inserted in the NMR instrument. Detailed experimental conditions for stopped-flow and 1H NMR measurements are listed in Table S1 (Supporting Information).

CO2 (aq) + OH− XooY HCO3− k −2

K3

CO32 − + H+ ⇄ HCO3−

(4)

K5

OH− + H+ ⇄ H 2O

(5)

The green section of Figure 1 is the protonation equilibria of sarcosine. Sarcosine exists in three differently protonated forms. Only the completely deprotonated form, SAR−, has a deprotonated amino group and thus is a Lewis acid, which is reactive toward CO2 (aq) and bicarbonate. Sarcosine is preneutralized with an equal molar amount of sodium hydroxide in this work. Most of the sarcosine is in the deprotonated form (SAR−); little is in the single protonated form (SARH). SAR− has two basic sites: the amino group and the carboxylate group. The amino group is much more basic and thus the single protonated form is CH3NH2+CH2COO−. Equation 6 is the protonation equilibrium of the amino group, and eq 7 is the protonation equilibrium of the carboxyl group. The temperature-dependent protonation constant K6 has been reported by several researchers; K6 values used in this work are taken from Bunting et al.25 and Datta et al.26 The values of K10 are only published at 25.0 °C and the published value of log(K10) is 2.2 at 25.0 °C.27 CH3NHCH 2COO− (SAR−) + H+

DATA ANALYSIS A reaction scheme including all the reactions in sarcosine− CO2−water system is displayed in Figure 1. It includes CO2(aq) hydration, sarcosine carbamate/carbamic acid formation, and all protonation equilibria reactions. The reactions of CO2(aq) with water and hydroxide ions are indicated by a red dashed line in eqs 1−5. These reactions have been well studied previously. All the rate and equilibrium constant values used in this work are obtained from published literature.3,21−24

K6

⇄ CH3NH 2+CH 2COO−(SARH)

(6)

CH3NH 2+CH 2COO− (SARH) + H+ K10

XooY CH3NH 2+CH 2COOH(SARH 2+)

(7)

Carbamate/carbamic acid formation reactions and the carbamate protonation equilibrium displayed in Figure 1 are provided as eqs 8−10. We used global data fitting of a series of measurements from stopped-flow experiments to study k7, k−7, and K8. For each temperature, five forward reaction measure-

k1

k −1

(3)

K4

HCO3− + H+ ⇄ H 2CO3

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CO2 (aq) + H 2O XooY H 2CO3

(2)

(1) 10278

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Figure 3. Calculated species concentration profiles and pH in the reaction of 6 mM SAR− with 8.2 mM CO2(aq) at 25.0 °C: (a) initial 0.1 s; (b) total 13 s.

Figure 4. (a) Absorbance at 520 nm versus time in the reaction of pre-equilibrated solution of carbamate containing 7.52 mM carbamate, 17.48 mM SAR− and 42.48 mM HCO3− with 50 and 55 mM hydrochloric acid in the presence of 50 μM alizarin red S and 25 μM methyl orange indicators at 25.0 °C. Total concentrations of sarcosine and bicarbonate are 25 and 50 mM, respectively. Solid markers are measured data and lines are calculated profiles. (b) Calculated species concentration profiles and pH versus time in the reaction between pre-equilibrated solution of carbamate with 50 mM hydrochloric acid at 25.0 °C.

ments and three backward reaction measurements are used for data fitting. We made a series of NMR measurements to study the slow carbamate formation and decomposition reaction (eq 10) at high pH. As a consequence of the principle of microscopic reversibility, k −9 is defined by other constants as k−9=k9*K8*k−7*k1/(k7*k−1*K4) in the data fitting program.

The protonation constants of three acid−base indicators are not published at all temperatures. Therefore, in this case, they need to be fitted to the data. In this study, the aqueous CO2 solution was prepared by bubbling the mixture gas of N2 and CO2 into Milli-Q water and the initial concentrations of CO2(aq) were adjusted by changing the N2 and CO2 gas flow rate. The accuracy of the CO2(aq) concentrations is limited by the accuracy of the gas flow meter, and therefore the initial concentrations of CO2(aq) are also fitted for the final analyses.

CO2 (aq) + CH3NHCH 2COO− (SAR−)

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k7

XooY CH3N(CO2 H)CH 2COO− (SAR−(CO2 H)) k −7

(8)

RESULTS Formation Reaction of Carbamate Followed by Stopped-Flow Spectrophotometric Studies. Figure 2a displays the absorbance change (at 590 nm) versus reaction time at various initial concentrations of SAR− ([SAR−]0 = 3−12 mM) and aqueous CO2 ([CO2]0 = 8.2 mM) in the presence of 12.5 μM thymol blue at 25.0 °C. The solid markers are measured data, and the lines are calculated profiles from data fitting. The absorbance drops rapidly in the first 0.04 s, following by much slower absorbance change. For most

CH3N(CO2−)CH 2COO− (SAR−(CO2−)) + H+ K8

⇄ CH3N(CO2 H)CH 2COO− (SAR−(CO2 H))

(9)

HCO3− + CH3NHCH 2COO− (SAR−) k9

XooY CH3N(CO2−)CH 2COO− (SAR−(CO2−)) ( +H 2O) k −9

(10) 10279

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Table 1. Calculated Rate and Equilibrium Constants for the Reaction of SAR− with CO2(aq) and the Protonation Constants of Sarcosine Carbamate at Different Temperatures k7

CO2 (aq) + CH3NHCH 2COO− XoooY CH3N(CO2 H)CH 2COO−

(R1-1)

k−7 K8

CH3N(CO2−)CH 2COO− + H+ ⇄ CH3N(CO2 H)CH 2COO−

(R1-2)

reaction

constant

15.0 °C

25.0 °C

35.0 °C

45.0 °C

R1-1

k7 (M−1 s−1) k−7 (s−1) K7 (M−1) log(K8)

8.1(1) × 103 35(2) 2.3(1) × 102 7.5(1)

18.6(6) × 103 88(4) 2.1(1) × 102 7.5(1)

41.7(6) × 103 2.1(4) × 102 2.0(3) × 102 7.6(1)

91(9) × 103 5.0(6) × 102 1.8(3) × 102 7.8(1)

R1-2

Figure 5. (a) Arrhenius plots of ln(k7) and ln(k−7) versus 1/T for the forward and backward reaction of SAR− with CO2(aq). (b) van’t Hoff plots of ln(K7) and log(K8) versus 1/T for the reaction of SAR− with CO2(aq) and carbamate protonation.

solution decreases from 10.8 to 8.5 over time, SARH concentration increases. Decomposition Reaction of Carbamate Followed by Stopped-Flow Spectrophotometric Studies. We also studied the decomposition reaction of carbamate at low pH over the temperature range of 15.0−45.0 °C. The pH range during this reaction is about 3−6.2: larger than the forward reaction. Figure 4a displays the absorbance change (at 520 nm) observed during the reaction between the pre-equilibrated carbamate solution (7.52 mM carbamate, 17.48 mM SAR−, and 42.48 mM HCO3−) with hydrochloric acid (50 and 55 mM) in the presence of 50 μM alizarin red S and 25 μM methyl orange indicators at 25.0 °C. Figure 4b shows the calculated species concentration profiles versus time for 50 mM hydrochloric acid measurement. During this reaction, the carbamate is mainly in the protonated form of SAR−(CO2H) and decomposes via the k−7 pathway to form CO2(aq) and SAR−: SAR− is instantly protonated to form SARH. All these reactions occur in approximately 0.2 s. A small amount of sarcosine is present in the double-protonated form of SARH2+ when the pH is about 3, and is deprotonated to SARH when the pH increases. Global data fitting combined carbamate formation and decomposition measurements to fit three unknown parameters: k7, k−7, and K8. The corresponding rate and equilibrium constants are listed in Table 1 over the temperature range 15.0−45.0 °C. Temperature Dependence. Eyring, Arrhenius, and van’t Hoff relationships were used to study the temperature dependence of the rate and equilibrium constants. Figure 5a

measurements the absorbance change is recorded for about 10 s and the initial 0.1 s is the most critical time range. Figure 2b shows the absorbance change (at 590 nm) versus reaction time at three different initial concentrations of CO2(aq) ([CO2]0 = 8.37−13.1 mM) and a fixed initial concentration of SAR− ([SAR−]0 = 8 mM) at 15.0 °C. The time axis is in common logarithm scale. The time interval between two markers is 0.0025 s, which is the smallest integration time of the diode-array detector. Calculated concentration profiles and pH in the reaction of 6 mM SAR− with 8.2 mM CO2(aq) at 25.0 °C are displayed in Figure 3. Figure 3a shows the concentration profiles in the initial 0.1 s of the reaction. The concentration of SAR− drops dramatically, accompanied by a rapid increase of SARH and SAR− (CO2−) concentrations and decrease in CO2(aq) concentration. This is due to the fast reaction of CO2(aq) with SAR−, which leads to the formation of carbamic acid, which is instantly deprotonated to carbamate. Most SAR− converts to SARH during the reaction. Figure 3b shows the concentration profiles in the complete 13 s stopped-flow investigation time; the time axis is in common logarithm scale. After 0.1 s, most SAR− is consumed; the concentration of SAR− is very low and remains almost unchanged. Due to the low concentration of SAR−, its contribution to the consumption of CO2(aq) is less important than the relatively slow reactions between CO2(aq) and OH−/water to form bicarbonate. Decreasing CO2(aq) concentration causes the decomposition of carbamate as a shift of eq 8 to left. The pH values during the reaction are shown on a secondary y-axis. As the pH of the 10280

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shows the Arrhenius plots for fitted rate constants k7 and k−7. The linear correlation coefficients (R2) for ln(k7) versus 1/T and ln(k−7) versus 1/T are both greater than 0.98, showing good linearity. Figure 5b shows the van’t Hoff plots for the calculated equilibrium constants K7 and K8. The corresponding activation parameters, enthalpy and entropy values over the temperature range of 15.0−45.0 °C are listed in Table 2.

10.22 during the reaction. A large amount of both deprotonated and single-protonated forms of sarcosine exist in the solution, because the pKa value of the amino group of sarcosine is 10.05, which is quite close to the pH range during this reaction. The concentration profiles of sarcosine carbamate versus time at 25.0 °C in the carbamate decomposition reaction at high pH are displayed in Figure 7a. The corresponding calculated species concentration profiles and pH versus time in this reaction are shown in Figure 7b. Carbamate is relatively stable at high pH, and thus the decomposition reaction is slow, taking approximately 8 h to reach equilibrium. By using global data fitting with the 1H NMR and stoppedflow measurement results, we obtained rate and equilibrium constants (k9/k−9) for the direct reaction between SAR− and HCO3− at 25.0 °C. The parameters required for data fitting were obtained from published values and the former stoppedflow fitting results (k7, k−7, and K8). The corresponding rate and equilibrium constants are listed in Table 3.

Table 2. Calculated Activation Parameters, Enthalpies, and Entropies for the Reaction of SAR− with CO2(aq) and Sarcosine Carbamate Protonation k7

CO2 (aq) + CH3NHCH 2COO− → CH3N(CO2 H)CH 2COO−

(R2-1)

k−7

CH3N(CO2 H)CH 2COO− ⎯⎯→ CO2 (aq) + CH3NHCH 2COO−

(R2-2) K7

CO2 (aq) + CH3NHCH 2COO− ⇄ CH3N(CO2 H)CH 2COO− K8

CH3N(CO2−)CH 2COO− + H+ ⇄ CH3N(CO2 H)CH 2COO− Arrhenius reactions R2-1 R2-2 R2-3 R2-4

(R2-3)

■

(R2-4)

DISCUSSION This work investigated two reaction pathways of the formation and decomposition reactions of sarcosine carbamate in detail. These are the k7/k−7 pathway (reversible reaction between SAR− and CO2(aq) to form carbamate and carbamate decomposition) and the k9/k−9 pathway (reversible reaction between SAR− and HCO3− to form carbamate and carbamate decomposition). SAR− + CO2(aq): k7/k−7 and K7. From a PCC point of view, the formation of carbamate from the k7/k−7 pathway is critical. The rate constant k7 between amine and CO2 to form carbamate is one of the most important parameters for absorbent selection. Fast absorbents with a large k7 are favorable for the PCC process. In this work, the rate constant k7 obtained by global data fitting was 18.6(6) × 103 M−1 s−1 at 25.0 °C. No similar study for the rate constant of SAR− with CO2 has been published to our knowledge. The same rate constants for several other amines have been reported in the literature. The published rate constants k7 of aqueous ammonia, aqueous MEA and aqueous piperazine in the reactions with CO2(aq) to form carbamate at 25.0 °C are as follows: aqueous ammonia, 215−450 M−1 s−1;19,28 aqueous MEA, 4900−6000 M−1 s−1;20,29 aqueous

Eyring/van’t Hoff

Ea (kJ mol−1)

A

61(1) 67(2)

9.5 × 1014 4.0 × 1013

ΔH‡/ΔH⌀ (kJ mol−1)

ΔS‡/ΔS⌀ (J mol−1 K−1)

59(1) 64(2) −5(3) 19(4)

33(4) 7(8) 26(9) 210(14)

Formation and Decomposition Reactions of Carbamate via the Bicarbonate (k9/k−9) Pathway, Followed by 1 H NMR Spectroscopy. The slow reactions between SAR− and bicarbonate (eq 10) were followed by 1H NMR spectroscopy at 25.0 °C. Figure 6a shows the concentration profiles of sarcosine carbamate versus time for the direct reaction between SAR− and HCO3− in the presence and absence of additional sodium hydroxide. Under these conditions, a substantial amount of carbamate is formed via the HCO3− (k9/k−9) pathway. Figure 6b shows the calculated concentration profiles versus time in the reaction of 50 mM SAR− with 100 mM HCO3− in the presence of 40 mM sodium hydroxide. The pH values are shown on a secondary y-axis and increase slightly from 10.18 to

Figure 6. (a) Carbamate concentration versus time in the reaction of 50 mM SAR− with 100 mM HCO3− in the presence of 0 and 40 mM sodium hydroxide at 25.0 °C. Solid markers are measured data and lines are calculated profiles. (b) Calculated species concentration profiles and pH versus time in the reaction of 50 mM SAR− with 100 mM HCO3− in the presence of 40 mM sodium hydroxide at 25.0 °C. 10281

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Figure 7. (a) Sarcosine carbamate concentration profiles versus time in the decomposition reaction of pre-equilibrated carbamate solution with 150 mM sodium hydroxide. Total concentrations of sarcosine and bicarbonate of the pre-equilibrated carbamate solution are 50 and 150 mM, respectively. Solid markers are measured data and lines are calculated profiles. (b) Calculated species concentration profiles and pH versus time in this reaction.

confirm that sarcosine has a very fast CO2 absorption rate. This is an important factor in absorbent selection. The temperature-dependent expressions for k7 and k−7 in Arrhenius forms are k7 = 9.5 × 1014 exp(−7348/T); k−7 = 4.0 × 1013 exp(−8004/T). Over the 30.0 °C temperature range studied, k7 and k−7 values increase by 11 and 14 times, respectively. At other temperatures, k7 and k−7 can be estimated using these two equations. At 60.0 °C, k7 was calculated as 2.5 × 105 M−1 s−1. In a study by Bishnoi et al.,30 the corresponding k7 value for piperazine is 2.2 × 105 M−1 s−1. This shows that sarcosine has a CO2 absorption rate comparable to that of piperazine at high temperatures. Sarcosine could therefore have PCC potential at relatively high temperatures. K7 values were calculated on the basis of the corresponding forward and backward rate constants of k7 and k−7. The temperature dependence of K7 can be expressed in the van’t Hoff equation; the error limit is larger than the rate constants k7 and k−7. K7 did not change significantly with temperature, decreasing from 2.3(1) × 102 M−1 to 1.8(3) × 102 M−1 over the temperature range 15.0−45.0 °C. The corresponding enthalpy and entropy of the reaction can be obtained by the van’t Hoff analyses. Data fitting for each repeat measurement will contribute up to ±20% error to the rate and equilibrium constants. This limits the accuracy of enthalpy and entropy of the equilibrium constants listed in Table 2. According to the full reaction scheme for CO2 absorption in aqueous sarcosine solution, SAR−, water and hydroxide ions can all contribute to CO2 absorption. The reaction between CO2(aq) and water is quite slow compared with the other two reactions, and it can be disregarded during the carbamate formation period. The reaction between CO2(aq) and OH− is fast; the corresponding rate constant k2 (12.1 × 103 M−1 s−1 at 25.0 °C) is comparable to the carbamate formation reaction rate constant k7. We estimate that the OH− pathway can contribute up to 5% of total absorption of CO2(aq) at the carbamate formation period under our experimental conditions. After the carbamate formation period, CO2(aq) hydration with water and OH− become important for CO2 absorption in aqueous solutions. This leads to the decomposition of carbamate, due to the reductions in pH and CO2(aq) concentration. In the PCC process, the amine concentration

Table 3. Calculated Rate and Equilibrium Constants for the Reaction of SAR− with Bicarbonate at 25.0 °C k9

HCO3− + CH3NHCH 2COO− XoooY CH3N(CO2−)CH 2COO− ( +H 2O) k−9

(R3-1) reaction

constant

values at 25.0 °C

R3-1

k9 (M−1 s−1) k−9 (s−1) K9 (M−1)

5.9 × 10−3 6.0 × 10−4 9.9

piperazine, 53700−70000 M−1 s−1.30,31 The rate constant k7 for SAR− at 25.0 °C is about 3 times larger than MEA and 40 times larger than aqueous ammonia. It is smaller only than piperazine, a well-known fast CO2 absorbent. This result agrees well with our previous wetted wall column results.14 The low solubility of piperazine in low-CO2 loading solutions limits the industrial use of concentrated piperazine. However, sarcosine does not have such solubility issues, and is more environmentally friendly than piperazine, MEA and aqueous ammonia, making it a promising solvent for PCC. The kinetics of CO2 absorption in aqueous sarcosine salt solution have been studied by Simons et al.15 and Van Holst et al.13 Both used gas−liquid reactors and applied the zwitterion mechanism to define the CO2 absorption rate shown in eq 11. In the zwitterion mechanism, the reaction order of amine is an empirical value and is obtained by data fitting under specific conditions. Simons et al. reported the reaction order n for potassium sarcosinate as 1.66 at 25.0 °C, and Van Holst et al. reported the value as 1.41 for potassium sarcosinate and 1.49 for lithium sarcosinate at 25.0 °C. In the zwitterion mechanism, the unit of the rate constant k′ of the reaction between amine and CO2 is related to the reaction order n. As a consequence, the values of k′ are not comparable between these two studies or to the k7 value in our work. R CO2 =

k′[Amine][CO2 ] 1+

k −1 ∑ k b[B]

= k′[Amine]n [CO2 ] (11)

Although k′ values from published work and the forward reaction rate constant k7 from this study have different definitions and are not comparable, all published results 10282

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and 19(4) kJ mol−1, respectively. Temperature-dependent protonation constants of sarcosine (eq 6) have been published. Enthalpy values for this reaction are calculated as −40.5 kJ mol−1 by Datta et al.,32 and the enthalpy values for the hydration of CO2(aq) to form HCO3− (+OH−) and H2CO3 (+H2O) have been reported by Wang et al.3 To calculate the overall enthalpy for CO2 absorption by aqueous sarcosine solution, the enthalpies of the protonation of carbonate and hydroxide also need to be included, all of which are published. If the cyclic range of CO2 loading between lean solutions and rich solutions is known, the species concentrations in lean and rich solutions can be simulated by our reaction scheme. Thus, the overall enthalpy change between lean and rich solutions can be derived.

will be much higher, and we expect the influence of the reaction between OH− and CO2 to be even lower in carbamate formation period. However, the CO2 hydration pathway maintains to be important to HCO3− formation and carbamate decomposition. Protonation of Carbamate: K 8 . This work also determined the protonation constant of carbamate K8 by global data fitting. The value of log(K8) was 7.5(1) at 25.0 °C. Determination of K8 tells us which form of carbamate will exist in the solution at different pH levels. In the PCC process with an aqueous sarcosine solvent, CO2 loading ([CO2]/[sarcosine] mol/mol) is usually less than 0.5, and pH of the solution will be greater than 9 at 25.0 °C. Under these conditions, most of the sarcosine carbamate will be in its deprotonated form. SAR− + HCO3−, k9/k−9 and K9. The k9/k−9 pathway is not as important as the k7/k−7 pathway for CO2 absorption in the PCC process. The reaction between SAR− and HCO3− is very slow compared with the reaction between SAR− and CO2(aq). The CO2-lean solution coming from the stripper will contain a substantial amount of HCO3−, and a small amount of carbamate will form via the k9/k−9 pathway. Studies of the k9/k−9 reaction can contribute to building up a full story of the sarcosine reaction mechanism that is meaningful from the chemistry point of view. In this study, the kinetics of the reversible reaction between SAR− and HCO3− were investigated using 1H NMR spectroscopy at 25.0 °C. The forward reaction to form carbamate was investigated at high pH, under which almost no CO2(aq) exists in the solution, and thus the HCO3− pathway becomes important. The backward reaction for carbamate decomposition was also studied at high pH, ranging from 11.7 to 12.3 during the reaction. In this pH range, essentially all of sarcosine carbamic acid is deprotonated, which ensures the importance of the k−9 pathway for carbamate decomposition. The rate constants k9 and k−9 obtained by global data fitting were 5.9 × 10−3 M−1 s−1 and 6.0 × 10−4 s−1, respectively. In this work, the formation and decomposition reactions of carbamate at high pH are thought to occur via both the k7/k−7 pathway and the k9/k−9 pathway. Another possible explanation is that the carbamate formation and decomposition reactions occur via the k7/k−7 pathway only. In the data fitting process, reaction models of both the k7/k−7 and k9/k−9 pathway (full model) and the k7/k−7 pathway only (reduced model) are used for global fitting of NMR and stopped-flow results. The fitting results of the reduced model were inferior, similar to our previous MEA work.20 On the contrary, the fitting results of the full model were much closer to the experimental results. That confirms the important influence of the k9/k−9 pathway on the formation and decomposition of carbamate at high pH. Without the k−9 decomposition reaction, the carbamate would be increasingly stable at high pH as the concentration of the carbamic acid continuously decreases. Thermodynamics. Evaluating absorbents for the PCC process requires knowledge of the energies involved in the absorption/desorption process. Temperature-dependent rate constants give information about activation energy, activation enthalpy and activation entropy. The Arrhenius activation energy for the reaction of CO2(aq) with SAR− (eq 8) is 61(1) kJ mol−1, and the activation enthalpy and activation entropy from the Eyring relationship is 59(1) kJ mol−1 and 33(4) J mol−1 K−1, respectively. Van’t Hoff analyses of the temperature-dependent equilibrium constants of eqs 8 and 9 yield an enthalpy value of −5(3)

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CONCLUSIONS

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ASSOCIATED CONTENT

An important aim in PCC research is to reduce the energy and cost requirements per unit of captured and purified CO2. At this stage of the general development, research is directed toward the gathering of information and the gaining of understanding of the relationships between the structures and the relevant properties of a wide range of amines. There are two aspects that are crucially important, they are (a) the kinetics of the reactions of CO2 in the amine solution and (b) the energy requirements in the stripper. The kinetics of the interactions in the absorber column defines its physical size and thus cost of construction and maintenance. The energy requirements for the PCC process are more complex, they involve several reaction enthalpies but a crucially important aspect is the cyclic capacity of the solution, the amount of solution that is required in the cyclic process for the purification of a unit of CO2. This capacity is determined by the temperature dependence of a series of equilibrium constants. This work investigated aqueous sarcosine as a new absorbent for CO2 capture. We determined the rate and equilibrium constants for all reactions in the sarcosine−CO2−water system over the temperature range 15.0−45.0 °C. Analysis afforded the activation energy (Arrhenius), activation enthalpy and entropy (Eyring), and enthalpy and entropy (van’t Hoff) of the reactions. Of particular interest is the fast reaction of sarcosine with CO2 to form the carbamate. The collection of these fundamental parameters allows the detailed modeling of a PCC plant based on aqueous sarcosine solutions. This work will also support further fundamental studies and applied research on the aqueous sarcosine-based process for CO2 capture.

S Supporting Information *

Experimental conditions for both stopped-flow and 1H NMR measurements, Table S1. This information is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*M.M.: tel, +61 2 4921 5478; fax, +61 2 4921 5472; e-mail, [email protected]. H.Y.: tel, +61 2 4960 6201; fax, +61 2 4960 6021, e-mail: [email protected]. Notes

The authors declare no competing financial interest. 10283

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(31) Derks, P. W. J.; Kleingeld, T.; Aken, C. V.; Hogendoorn, J. A.; Versteeg, G. F. Chem. Eng. Sci. 2006, 61, 6837−6854. (32) Datta, S. P.; Grzybowski, A. K. Trans. Faraday. Soc. 1958, 54, 1188−1194.

ACKNOWLEDGMENTS This project is part of the CSIRO Advanced Coal Technology Portfolio and received funding from the Australian Government under the China-Australia Joint Coordination Group on Clean Coal Technology Research Grant. The views expressed herein are not necessarily the views of the Commonwealth, and the Commonwealth does not accept responsibility for any information or advice contained herein.

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REFERENCES

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