Reactions of CO2 with Aqueous Piperazine Solutions: Formation and

Article pubs.acs.org/JPCA

Reactions of CO2 with Aqueous Piperazine Solutions: Formation and Decomposition of Mono- and Dicarbamic Acids/Carbamates of Piperazine at 25.0 °C William Conway,*,†,‡ Debra Fernandes,† Yaser Beyad,† Robert Burns,† Geoffrey Lawrance,† Graeme Puxty,‡ and Marcel Maeder*,† †

Department of Chemistry, The University of Newcastle, Callaghan, NSW 2308, Australia CSIRO Energy Technology, Mayfield West, NSW 2304, Australia

‡

S Supporting Information *

ABSTRACT: Piperazine (PZ) is widely recognized as a promising solvent for postcombustion capture (PCC) of carbon dioxide (CO2). In view of the highly conflicting data describing the kinetic reactions of CO2(aq) in piperazine solutions, the present study focuses on the identification of the chemical mechanism, specifically the kinetic pathways for CO2(aq) in piperazine solutions that form the mono- and dicarbamates, using the analysis of stopped-flow spectrophotometric kinetic measurements and 1H NMR spectroscopic data at 25.0 °C. The complete set of rate and equilibrium constants for the kinetic pathways, including estimations for the protonation constants of the suite of piperazine carbamates/carbamic acids, is reported here using an extended kinetic model which incorporates all possible reactions for CO2(aq) in piperazine solutions. From the kinetic data determined in the present study, the reaction of CO2(aq) with free PZ was found to be the dominant reactive pathway. The superior reactivity of piperazine is confirmed in the kinetic rate constant determined for the formation of piperazine monocarbamic acid (k7 = 2.43(3) × 104 M−1 s−1), which is within the wide range of published values, making it one of the faster reacting amines. The corresponding equilibrium constant for the formation of the monocarbamic acid, K7, markedly exceeds that of other monoamines. Kinetic and equilibrium constants for the remaining pathways indicate a minor contribution to the overall kinetics at high pH; however, these pathways may become more significant at higher CO2 loadings and lower pH values where the concentrations of the reactive species are correspondingly higher.

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INTRODUCTION The onset of climate change primarily by increasing concentration of carbon dioxide in the atmosphere and the recognition of issues relating to the contributions of anthropogenic carbon dioxide emissions from the generation of electricity are driving a corresponding investment in research and technologies to diminish emissions and improve the environmental impact of coal-fired electricity generation. In view of this, postcombustion capture (PCC) using reactive amine solvents to capture carbon dioxide present in the flue gas before its emission into the atmosphere is widely recognized as a promising technology capable of achieving the required shortterm reduction targets. The modest performance of the current industry standard monoethanolamine, MEA, is, in several respects, driving the refinement of the current process in the pursuit of highly efficient, alternative solvent systems thereby ensuring the success of the technology. The use of amines for CO2 capture processes is attractive due to the unique ability of primary and secondary amines to react directly and reversibly with CO2 to form the respective carbamates. The associated kinetic, equilibrium, and thermodynamic contributions for amines that follow the carbamate pathway are the key chemical criteria that need to be considered in the assessment of novel solvents for the PCC process. The fast formation of carbamate is a clear advantage © 2013 American Chemical Society

but unfortunately is also detrimental to the stoichiometric capacity of an amine solution for CO2 capture, whereby two amine molecules are consumed for each molecule of CO2 absorbedone amine in the formation of the carbamic acid and a second amine for its deprotonation to the carbamate. Conversely, despite the “ideal” 1:1 equilibrium CO2 loading of the alternative pathway (for CO2) via the hydration system of reactions with water in the presence of a nonreactive base, i.e., a tertiary amine, in which only one molecule of amine is consumed for each molecule of CO2, this hydration pathway is kinetically considerably slower. Since the physical size of the absorption and desorption columns, pumping equipment, and the associated capital and ongoing costs are directly linked among several other aspects to the chemical performance of the solvent, a reasonable compromise of the chemical properties must be found. An elegant solution to the above dilemma is in the use of diamine solvents whereby one amine group is involved in a fast reaction with CO2 to form the carbamate while the other amine absorbs the released proton. An example for this scenario is the cyclic diamine piperazine, PZ, which reacts rapidly with CO2 Received: October 24, 2012 Revised: December 23, 2012 Published: January 3, 2013 806

dx.doi.org/10.1021/jp310560b | J. Phys. Chem. A 2013, 117, 806−813

The Journal of Physical Chemistry A

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absorption chemistry.13 The constants associated with reactions 1−5 were taken from the literature.13−15

and thus has attracted extensive interest for its usage in CO2 capture applications, particularly as a rate promoter for CO2 absorption in carbonate and tertiary amine solutions.1,2 The collection of chemical reactions which describe the absorption of CO2 in piperazine solutions is considerably more complex than traditional monoamines, which ensues from the additional reactivity of the second amine group with CO2 and several protonation equilibria including the protonation of the PZ, the monocarbamate, and the dicarbamate species. These additional reactions cannot be ignored and thus must also be accounted for in any chemical mechanism to interpret the kinetic and equilibrium parameters for the system of CO2/diamine reactions. In view of this chemical complexity, it is not unexpected therefore to discover that the published constants derived in previous investigations of CO2/piperazine kinetics are noticeably conflicting. Such inconsistencies within the reported kinetic data are often attributed to the substantially different measurement techniques and nonspecific mechanisms used in the acquisition and interpretation of the data. Furthermore, the contributions of the reactive pathways involving the monoprotonated piperazine (PZH+), and the monocarbamate/carbamic acids (PZCO2−/PZCO2H), with CO2, are essentially unresolved.3−11 Thus, the thorough investigation of the mechanisms and kinetic constants governing the pathways in diamine solutions is still of critical importance. The work herein is focused on two aspects of the aqueous CO2/piperazine system: (a) the determination of the chemically complete mechanism to describe all reactions of CO2(aq) within pure and CO2-loaded piperazine solutions and (b) the quantification of the kinetic and equilibrium constants to describe the reversible formation of carbamates/carbamic acids in aqueous piperazine solutions from the analysis of stoppedflow spectrophotometric and quantitative 1H NMR spectroscopic measurements at 25.0 °C.

k1

CO2 (aq) + H 2O XooY H 2CO3 k −1

(1)

k2

CO2 (aq) + OH− XooY HCO−3 k −2

K3

CO32 − + H+ ↔ HCO−3 K4

HCO−3 + H+ ↔ H 2CO3 K5

OH− + H+ ↔ H 2O

(2) (3) (4) (5)

An additional known part of the overall mechanism includes the two protonation equilibria of piperazine as in eq 6. The corresponding values for the first and second protonation constants of piperazine, log K6a = 9.77 M−1 and log K6b = 5.60 M −1 , at 25.0 °C, respectively, have been determined experimentally via potentiometric titrations of the free amine in our previous investigations.16 K 6a

PZ + H+ ← → PZH+ K 6b

PZH+ + H+ ←→ PZH 22 +

(6)

The reaction mechanism that describes all interactions of PZ with CO2, including all protonation equilibria, is best represented schematically as shown in Figure 1.

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KINETIC MODELS The homogeneous interactions of CO2 in aqueous piperazine solutions incorporate a complex series of parallel and reversible reaction pathways including the hydration system of reactions for CO2 with water and hydroxide and those with the amine to form carbamates. From the interpretation of 1H and 13C NMR data, both the mono- and dicarbamates of piperazine are present in appreciable (quantifiable) concentrations in equilibrium solutions.12 Thus, the collection of kinetic pathways that are used to procure the formation and decomposition of these species must also be accounted for and quantified in the chemical mechanism. It should be noted that we use two-way arrows (i.e., ⇄) to represent equilibiria with kinetically measurable rate constants, while a doubleheaded arrow (↔) represents an instantaneous protonation equilibrium. Once dissolved into solution, CO2(aq) is involved in two reversible reactions with water (H2O), k1, k−1, and hydroxide (OH−), k2, k−2, eqs 1 and 2. Despite the high concentration of the amine in PCC solutions and the alluring assumption that the amine pathway with CO2(aq) will dominate the kinetics, the latter reaction with hydroxide contributes significantly to the observed kinetics, particularly in moderate to highly basic amines (note that the concentration of hydroxide increases with amine basicity) and in tertiary amine solutions. In addition, the protonation equilibria of the carbonate species and that of the hydroxide ion, K3, K4, K5, eqs 3−5, form an integral part of the

Figure 1. Extended kinetic model used to describe the reversible formation of piperazine carbamates from the reactions of CO2(aq) with free piperazine (PZ), protonated piperazine (PZH+), piperazine monocarbamate (PZCO2−), and piperazine monocarbamic acid (PZCO2H). The rate and equilibrium constants are given in this scheme; see the Results section for more information. Note that the CO2 hydration reactions do not appear in the figure for clarity.

The mechanism includes the following set of reactions of CO2(aq) with unprotonated and monoprotonated piperazine (PZ/PZH+) and with the unprotonated and monoprotonated monocarbamates of piperazine (PZCO2−/PZCO2H): k7

PZ + CO2 (aq) XooY PZCO2 H k −7

(7)

k8

PZH+ + CO2 (aq) XooY PZ(H+)CO2 H k −8

807

(8)



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k9

PZCO−2 + CO2 (aq) XooY PZ(CO−2 )CO2 H k −9

in this work are outlined in detail in our previous contributions.13,18−22 In order to globally determine the kinetic and equilibrium constants for the parallel reactive pathways for CO2 in piperazine solutions, as in Figure 1, the kinetic formation and decomposition of piperazine carbamates have been investigated in a series of measurements over a range of initial concentrations observing both the formation and decomposition reactions. All kinetic measurements were performed on an Applied Photophysics DX-17 spectrophotometer equipped with a J&M Tidas MCS 500-3 diode-array detector, observing the pH changes over the wavelength range 400−700 nm via coupling to colored acid−base indicators. All reactions were performed at 25.0 °C. Samples were thermostatted by a circulating Julabo F20 water bath at 25.0 (±0.1) °C, and the exact temperature was reported by a thermocouple located in the stopped-flow instrument mixing chamber. Formation Reactions of Piperazine Carbamates. In these experiments solutions of CO2(aq) were mixed in the stopped-flow instrument with a second solution containing indicator and various concentrations of PZ. The concentrations of reagents were chosen such that after 1:1 mixing the initial concentrations were [CO2]0 = 4.35 mM, [PIPZ]0 = 1.0−6.0 mM, and 12.5 μM thymol blue indicator. To investigate the reactions at higher pH, sodium hydroxide ([OH−]0 = 1.0−4.0 mM) was added to a PZ solution ([PZ]0 = 4.0 mM), using the same initial concentrations of CO2 and indicator. In order to observe the reaction of CO2(aq) with monoprotonated piperazine, the pH of a series of piperazine solutions was adjusted down by the addition of HCl. After mixing, initial concentrations were [CO2]0 = 5.1 mM, [PZH+]0 = 1.0−3.0 mM, [HCl]0 = 1.0−3.0 mM, with 12.5 μM thymol blue. Decomposition Reactions of Piperazine Carbamates. The reversibility of the reactions was studied during the decomposition of an equilibrated solution of piperazine carbamate following the addition of hydrochloric acid. A solution containing 0.15 M HCO3− and 0.05 M piperazine was first prepared and equilibrated for 24 h at 25.0 °C. Under these conditions carbamates are formed slowly, and the concentrations of the amine and carbamate(s) species were determined quantitatively from the corresponding 1H NMR spectrum of the equilibrated solution, as shown in Figure 2. This equilibrated carbamate solution (NMR concentrations = 14.0 mM monocarbamate, 4.8 mM piperazine dicarbamate, 126.4 mM HCO3−, and 31.2 mM piperazine) was transferred into one of the stopped-flow syringes and mixed with an equal volume of a hydrochloric acid solution ([H+] = 90.0−100.0 mM) containing 0.025 mM methyl orange and 0.05 mM alizarin red S indicators. As with our previous investigations, we have employed global analysis methods using ReactLab (www.jplusconsulting.com) and modified in-house extensions of it in which the entire collection of individual stopped-flow absorption measurements are analyzed together in a single unit. Reactlab is based on traditional nonlinear least-squares fitting methods.22−24 In this way the rate and equilibrium constants describing the reactions are common to all measurements throughout the fitting process, resulting in a robust set of rate and equilibrium constants for all reactions. Debye−Hückel approximations for species activities were incorporated during the analysis for all charged species. Hence, all rate and equilibrium constants are thermodynamically correct and valid at zero ionic strength.

(9)

k10

PZCO2 H + CO2 (aq) XoooY PZ(CO2 H)CO2 H k −10

(10)

The following protonation equilibria for the mono- and dicarbamate complete the reaction scheme: K11

PZCO−2 + H+ ↔ PZ(H+)CO−2

(11)

K12

PZ(H+)CO−2 + H+ ← → PZ(H+)CO2 H K13

PZ(CO−2 )CO−2 + H+ ↔ PZ(CO−2 )CO2 H K14

PZ(CO−2 )CO2 H + H+ ← → PZ(CO2 H)CO2 H

(12) (13) (14)

This reaction scheme requires the determination of eight rate constants for eqs 7−10, and four protonation constants, eqs 11−14. The principle of microscopic reversibility allows the determination of two rate constants based on other constants in the two closed reaction loops. Hence, the rate constant k−8 is calculated based on other constants in a loop as k−8 = (k8k−7K6b)/(K12k7). Similarly, k−10 is calculated based on the constants in a loop as k−10 = (k10k−9K11)/(K14k−9). The equilibrium constants K7 and K9 were determined based on the relationship between the formation and decomposition rate constants, i.e., Ki = ki/k‑i. All constants were updated simultaneously and automatically throughout the fitting procedure so as to maintain microscopic reversibility. In view of the difficulties in the determination of the protonation constants of the carbamate species due to the low concentration of the acid forms in solution, the second protonation constant of the dicarbamic acid, log K14, was fixed to a value of 0.5 log units lower than the first protonation constant, log K13. Such an assumption eliminates the high correlation between the two parameters in the fitting and is consistent with the relationship observed between the first and second protonation constant of similar diacid species (e.g., EDTA, adipinic acid, dibenzoic acid).17 Published values for the first and second protonation constants of PZ, as well as the rate and equilibrium constants describing the hydration reactions, reduce the total number of parameters to be determined in the analysis here to 10.

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EXPERIMENTAL SECTION Chemicals. High-purity CO2 gas (BOC), N2 (Coregas), potassium bicarbonate (BDH), sodium hydroxide (Merck), piperazine (Sigma-Aldrich), potassium hydrogen phthalate (AJAX), hydrochloric acid (AJAX), thymol blue sodium salt (Sigma-Aldrich), alizarin red S (BDH), and methyl orange (SELBY) were all used as obtained without further purification. Stock solutions of piperazine, NaOH, and HCl solutions were prepared and standardized by potentiometric titration. Ultrahigh purity Milli-Q water was boiled to remove dissolved CO2 gas and used to prepare all solutions inside a CO2 free atmosphere glovebag. Stopped-Flow Measurements. The kinetic and equilibrium reactions of CO2(aq) in aqueous piperazine solutions have been investigated by observing the absorption change of colored acid−base indicators (and subsequently changes in pH) upon mixing of solutions in a stopped-flow spectrophotometer. The procedure for the stopped-flow measurements performed 808



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which were calculated based on a statistical analysis of five repeats of the stopped-flow measurement data. Note that all rate and equilibrium constants are also displayed in Figure 1. From the constants the reactivity of the four piperazine species toward CO2 follows the trend PZ > PZCO2− > PZH+ > PZCO2H; this order is chemically/structurally consistent; i.e., the positive charge on the PZH+ ion will reduce the Lewis basicity of the free amine group more than that of the carbamate substitution with a negative charge. The relative ratios between the dominant pathway for CO2(aq) via a reaction with free PZ, k7, and the minor pathways k8, k9, and k10 are approximately 47, 4, and 200 000, respectively. Excluding the k10 pathway which is some 5 orders of magnitude slower, the remaining pathways are not overly dissimilar to that of the free PZ. As a result, at higher loadings and lower pH conditions the constants for the minor pathways amount to a substantial contribution to the overall kinetics given that the corresponding concentrations of the reactive piperazine species are accordingly higher under these conditions. As the rate constants are not much smaller, the rate of the reaction is maintained over an extended period of time/pH. It is also interesting to compare the rate constants for the decomposition pathways. The monocarbamic acid can decompose directly via the k−7 pathway or via the additionally amine protonated monocarbamic acid via k−8. From the constants, the latter pathway is the kinetically preferred pathway given the 60-fold difference in the constants. However, this is strongly dependent on the pH of the solution. The situation for the decomposition of the dicarbamates is similar. In parallel with the determination of rate and equilibrium constants, the analysis produces the corresponding concentration profiles for all species during the reactions which are shown in Figures S1−S3 of the Supporting Information. The ability of a single piperazine molecule to react with CO2 (forming the carbamate) while absorbing the proton released during the formation (of the carbamate) is demonstrated in the right of Figure S1 whereby the concentration of the monocarbamic acid, and not the carbamate, increases with the time of the reaction. Such behavior is testament to the superior reactivity of diamines. Literature Values. Several studies report kinetic constants for the reactions of CO2 in piperazine solutions. Typically, the fast reaction of CO2 with the free amine, the k7 pathway here, has been investigated using homogeneous methods, i.e., stopped-flow kinetics or heterogeneous gas−liquid measurements, which additionally incorporate the mass transfer of CO2 gas into solution. Ideally, both methods should produce similar results within the error limits. The values for k7 determined here at 25.0 °C, together with published values, are shown in Table 2. For those studies that do not report values at 25 °C a range of values at different temperatures are provided. Considerable variation is observed in the constants which are spread over some 3 orders of magnitude. The value determined here for k7 (2.43(4) × 104 M−1 s−1) is in good agreement with several of the published values and agrees closely with the values determined by Rayer et al. (1.67 × 104 M−1 s−1)25 and Gordesli and Alper (1.8 × 104 M−1 s−1),27 who similarly determined the kinetics using the stopped-flow technique in their studies of piperazine kinetics and also the value determined by Sun et al. (2.1 × 104 M−1 s−1),7 via gas absorption methods. Several of the remaining published values are substantially higher than the value reported here. These variations can be rationalized to some extent by the differences in measurement techniques; the kinetics of mass transfer is

Figure 2. 1H NMR spectrum of an equilibrated HCO3−/PZ solution at 25.0 °C containing 7.0 mM piperazine monocarbamate, 2.4 mM piperazine dicarbamate, 63.2 mM HCO3−, and 15.6 mM piperazine. Peak integrals and chemical shifts have been removed from the spectrum for clarity.

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RESULTS AND DISCUSSION Kinetic Data. The complete set of rate and equilibrium constants for the four reversible pathways for CO2(aq), i.e., k7, k−7−k10, k−10, and K7−K10, including the protonation equilibria, K11−K14, of the carbamate species, are listed in Table 1 together with standard deviations for the parameters in parentheses Table 1. Rate and Equilibrium Constants for the Reactions of CO2(aq) with Piperazine Species at 25.0 °C, Including the Protonation Constants of Piperazine Carbamatesa reaction

K11

K12

log K12 = 6.35(7) [M−1]

PZ + CO2 (aq) XoooY PZCO2 H k−7

k8

PZH+ + CO2 (aq) XoooY PZ(H+)CO2 H k−8

k9

PZCO−2 + CO2 (aq) XoooY PZ(CO−2 )CO2 H k−9

k10

PZCO2 H + CO2 (aq) XooooY PZ(CO2 H)CO2 H k−10

−

constants k7 = 2.43(4) × 104 M−1 s−1 k−7 = 4.10(6) × 10−1 s−1 K7 = 6.0(2) × 104 M−1 k8 = 5.2(3) × 102 M−1 s−1 k−8 = 23(3) s−1 K8 = 23(4) M−1 k9 = 5.6(1) × 103 M−1 s−1 k−9 = 59(4) s−1 K9 = 96(5) M−1 k10 = 0.13(9) M−1 s−1 k−10 = 0.2(1) s−1 K10 = 0.8(1) M−1 log K11 = 9.60(1) [M−1]

k7

PZ( H+)CO2 + H+ ← → PZCO2 H PZ(H+)CO−2 + H+ ←→ PZ(H+)CO2 H K13

PZ(CO−2 )CO−2 + H+ ←→ PZ(CO−2 )CO2 H K14

PZ(CO−2 )CO2 H + H+ ←→ PZ(CO2 H)CO2 H

log K13 = 7.99(2) [M−1] log K14 = 7.49(2) [M−1]

a

Numbers in brackets represent the error in the last digit (i.e., 2.43(4) × 104 is equivalent to 24300 ± 400 M−1 s−1). 809



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Table S1 of the Supporting Information. The rate constants determined in this study are comparable with those determined previously, with both published values about a factor of 10 higher. As a result of the kinetic analysis here, values for the protonation constants of the carbamates were also generated (see Table 1 for values). The value for K11 determined here, together with several published values for comparison purposes, is given in Table S2 of the Supporting Information. The value determined here is in good agreement with those determined previously, which is again validation of the kinetic model and the fitting procedure. Model Validation. The reaction mechanism presented in Figure 1, and by eqs 7−14, is inherently complex. During the development of the extended mechanism here several simpler/ reduced chemical mechanisms were investigated in parallel, and the results for one representative example are given here. This reduced model includes three reactions of CO2 with the amine, the protonated amine, and the monocarbamate. Each reaction is a combination of the addition reaction followed by the deprotonation of the initially formed carbamic acid.

Table 2. Second-Order Rate Constants for the Reaction of CO2(aq) with Piperazine (PZ) at 25.0 °C author our value Bishnoi/Rochelle4 Seo/Hong5 Sun et al.7 Derks et al.8 Samanta et al.9 Xu et al.3 Bindwal et al.11 Konduru et al.10 Zhang et al.6 Rayer et al.25 Bougie et al.26

Cullinane/ Rochelle2 Gordesli/Alper27

k7 (M−1 s−1)

temp (°C)

2.43(4) × 104 5.35 × 104 4.1 × 103 7.1 × 103 2.1 × 104

25.0 25 30 40 25

7.0 × 104 5.8 × 104 1.3 × 102 2.6 × 104 2.4 × 104 6.1 × 104 1.67 × 104 6.65 × 104 9.80 × 104 1.416 × 105 1.00 × 105 (1 M PZ)

25.15 25 25 30 30 30 25.15 30.15 35.15 40.15 25.0

2.27 × 103 1.13 × 104 1.80 × 104

5 15 25

technique stopped flow wetted wall gas absorption gas absorption stirred cell wetted wall disk column stirred cell stirred cell disk column stopped flow wetted wall

wetted wall

k7

A + CO2 (aq) XoooY ACO−2 + H+

stopped flow

k −7r

(15)

k8

AH+ + CO2 (aq) XoooY A(CO2 )H + H+ k −8r

relatively slow, and thus the analysis of the observations in terms of a much faster reaction is always difficult. Several researchers do not take into account the parallel reactions of CO2 with water, and in particular the reaction with hydroxide, and this thus lends to an overestimation of the reaction with PZ. In addition, the reversibility of the reactions and the impact of charged species on the reactivity (ionic strength) are often underestimated or overlooked in the analysis of mass transfer processes. Such effects have been noted under certain conditions to accelerate the overall rates of reaction, leading to higher contributions of the amine to the absorption kinetics.2 Both the reversibility and ionic strength are accounted for in the current study. Few values are reported for the minor kinetic pathways involving the reactions of CO2 with the monoprotonated piperazine, PZH+, k8, and with the piperazine carbamate, PZCO2−, k9, to form the mono- and dicarbamic acids. To our knowledge the constants determined in this work for the reversible reaction of CO2(aq) via the monocarbamic acid pathway, k10, k−10, are novel. Published kinetic data for the reaction of CO2(aq) via the k8 and k9 pathways are available in

(16)

k9

ACO−2 + CO2 (aq) XoooY [A(CO2 )2 ]2 − + H+ k −9r

(17)

The forward reactions and their rate constants are the same as those of the full mechanism; this includes the rate constants k7, k8, k9. However, the back reactions are different, as now they are defined as second-order reaction with the proton as one reagent. The resulting best fits are significantly inferior; two example plots of measured data with fits by the complete and the reduced mechanisms are shown in Figure 3 which shows the measurement traces at 590 nm for the reaction of 3.5 mM CO2(aq) with 4.0 mM piperazine, as well as that for the acid decomposition of a piperazine carbamate solution at 550 nm containing 7.0 mM piperazine monocarbamate and 2.4 mM piperazine dicarbamate, together with the calculated traces produced from the optimized analysis of the data using the reduced and extended kinetic models. Note that both plots feature a logarithmic time axis, and the differences are much less obvious if plotted with a linear time axis. A series of

Figure 3. (Left) Measurement and fitted traces at 590 nm in the reaction of 3.5 mM CO2(aq) with 4.0 mM piperazine in the presence of 12.5 μM thymol blue indicator; (right) measurement and fitted traces at 550 nm in the acid decomposition of an equilibrated carbamate solution containing 7.0 mM piperazine monocarbamate and 2.4 mM piperazine dicarbamate in the presence of 0.5 mM alizarin red S and 0.025 mM methyl orange indicators. Markers: ⧫ = measured data, dashed line = reduced model, solid line = complete model. 810



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Figure 4. Brønsted plot of the rate constant for the formation of carbamic acid from the reaction of CO2(aq) with amine, log k7, against protonation constant of the amine, log K6. The k7, k8, and k9 data for PZ are indicated as PZ, PZH+, and PZCO2−, respectively. ◇ = nonsterically hindered, ○ = sterically hindered, □ = NH3, ◆ = N-PIPZ.

representative measurement and fitted traces using the extended model for the reactions of CO2(aq) with free piperazine (PZ), partially protonated piperazine (PZH+), as well as the acid-decomposition of piperazine carbamate(s), are also shown in Figures S1−S3 of the Supporting Information. More important than the inferior fit are chemical reasons for the inadequacy of the reduced model. The pH dependence of the back reaction for the reduced model, a simple second-order reaction between the appropriate carbamate and a proton, is different from the pH dependence of a pre-equilibrium between carbamic acid and carbamate with decomposition of the carbamic acid. As the measurements cover a large pH range, this effect is clearly discernible. An additional consideration is that at very low pH the dicarbamate is completely protonated to the dicarbamic acid which is not reactive in the reduced mechanism and thus should be very stable. This observation is in clear contrast with the knowledge of carbamate stability under acidic conditions. It should be noted that HCO3− is also involved, in parallel, in the same series of reactions with each of the piperazine species.20,21 Ideally, a complete mechanism would include the HCO3− subset of reactions in addition to those examined in this work. However, the reactions with HCO3− are substantially slower (∼10−3 M−1 s−1) and do not contribute significantly to the kinetics on the millisecond/second time scale of the reactions investigated here. Brønsted Relationship. The relationship between the structure of the reactive amine species and its reactivity with CO2(aq) is not immediately apparent from the kinetic data alone. An obvious method of rationalizing the kinetic constants is to relate them to the protonation constants of the amine in a Brønsted type relationship. The protonation constant is a relative indication of the Lewis basicity which, due to the small size of the proton, is essentially not affected by steric hindrance. For a series of cyclic amines, there is a clear linear relationship between the log of the protonation constant and the log of the rate constant.21 Steric effects due to the larger size of the CO2 molecule, compared with the proton, manifest in lower than expected rate constants. Figure 4 demonstrates the relationship;

the dashed line is defined by a series of nonhindered amines (◇). The figure includes a selection of sterically hindered amines (○ for 1-amino-2-propanol (1-AP), sec-butylamine (SBA), and diethanolamine (DEA)) for which the rate constant is clearly lower than expected in the absence of steric effects. Ammonia also lies below the line (□), but here we assign the effect to strong and unique solvation in water. For PZ we include the rate constants for the uprotonated, k7, singly protonated, k8, of PZ as well as k9 for its monocarbamate (△). From the plot the kinetics of free piperazine, k7, and protonated piperazine, k8, are some ∼0.7 logarithmic units above the Brønsted values predicted by their basicities based on the trendline, while the monocarbamate, k9, reacts in line with the Brønsted trend. Molecular explanations for this type of behavior are not immediately obvious; however, it is possible that such increases are related to the role of the cyclic/diamine structure in the stabilization of intermediate transition state geometries for the carbamate species, which could result in lower activation energies and faster reaction rates. Charge interactions (i.e., repulsive/attractive interaction of charged groups in protonated piperazine and piperazine carbamate upon approach to the positively charged carbon in CO2) may also be a contributing factor. Mass Transfer Validation. The kinetic constants for the extended kinetic model determined here have been evaluated in a diffusion/chemical reaction model of CO2 absorption by a falling liquid film. The model and its implementation were as described in previous work.28 The CO2 absorption flux (NCO2) in piperazine solutions over a range of piperazine concentrations (0.45−1.5 M PZ) and CO2 loadings (0.0−0.019) was calculated and compared to literature data measured using a wetted wall contactor. The calculated data was compared to experimentally determined flux data at low CO2 partial pressures (PCO2 = 34−448 Pa) performed by Cullinane et al.2 This data set was selected as low CO2 loadings were used and the film properties (exposure time and thickness of the liquid film) could be estimated with reasonable confidence based on the density and viscosity data of pure piperazine solutions.2,29 Diffusion coefficients for CO2, piperazine, and all other 811



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Figure 5. Parity plot of calculated against measured absorption flux in piperazine solutions at 25 °C from wetted wall data at low CO2 partial pressures and loadings.

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chemical species and the Henry coefficient for CO2 were used as given in the study by Cullinane et al. The overall correlation of the calculated and experimentally determined flux data was assessed in a parity plot, which is given in Figure 5. Given the simple estimations for the physical parameters of the solutions in the mass transfer model, the agreement between the values for the calculated and observed flux is reasonable. The solution containing 0.6 M PZ and 0.15 M KOH at zero loading resulted in the highest agreement with the measured flux. It is likely the overall agreement would improve given better estimations for the solution physical properties and diffusion coefficients which begin to play a significant role at higher fluxes; however, the determination of these properties is outside the scope of this work. Nevertheless, the agreement is within the limits of around ±20%. The mass transfer behavior in solutions at higher CO2 loadings and lower pH conditions remains to be validated.

Notes

The authors declare no competing financial interest.

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REFERENCES

(1) Cullinane, J. T.; Rochelle, G. T. Fluid Phase Equilib. 2005, 227, 197−213. (2) Cullinane, J. T.; Rochelle, G. T. Ind. Eng. Chem. Res. 2006, 45, 2531−2545. (3) Xu, G. W.; Zhang, C. F.; Qin, S. J.; Wang, Y. W. Ind. Eng. Chem. Res. 1992, 31, 921−927. (4) Bishnoi, S.; Rochelle, G. T. Chem. Eng. Sci. 2000, 55, 5531−5543. (5) Seo, D. J.; Hong, W. H. Ind. Eng. Chem. Res. 2000, 39, 2062− 2067. (6) Zhang, X.; Zhang, C.-F.; Qin, S.-J.; Zheng, Z.-S. Ind. Eng. Chem. Res. 2001, 40, 3785−3791. (7) Sun, W.-C.; Yong, C.-B.; Li, M. H. Chem. Eng. Sci. 2005, 60, 503− 516. (8) Derks, P. W. J.; Kleingeld, T.; van Aken, C.; Hogendoorn, J. A.; Versteeg, G. F. Chem. Eng. Sci. 2006, 61, 6837−6854. (9) Samanta, A.; Bandyopadhyay, S. S. Chem. Eng. Sci. 2007, 62, 7312−7319. (10) Konduru, P. B.; Vaidya, P. D.; Kenig, E. Y. Environ. Sci. Technol. 2010, 44, 2138−2143. (11) Bindwal, A. B.; Vaidya, P. D.; Kenig, E. Y. Chem. Eng. J. 2011, 169, 144−150. (12) Ermatchkov, V.; Kamps, A. P.-S.; Maurer, G. J. Chem. Thermodyn. 2003, 35, 1277−1289. (13) Wang, X.; Conway, W.; Burns, R.; McCann, N.; Maeder, M. J. Phys. Chem. A 2010, 114, 1734−1740. (14) Harned, H. S.; Scholes, S. R. J. Am. Chem. Soc. 1941, 63, 1706− 1709. (15) Maeda, M.; Hisada, O.; Kinjo, Y.; Ito, K. Bull. Chem. Soc. Jpn. 1987, 60, 3233−3239. (16) Fernandes, D.; Conway, W.; Wang, X.; Lawrance, G.; Maeder, M.; Puxty, G. J. Chem. Thermodyn. 2012, 51, 97−102. (17) Academic Software, United Kingdom, www.acadsoft.co.uk. (18) McCann, N.; Phan, D.; Wang, X.; Conway, W.; Burns, R.; Attalla, M.; Puxty, G.; Maeder, M. J. Phys. Chem. A 2009, 113, 5022− 5029.

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CONCLUSIONS The kinetic formation and decomposition of piperazine carbamates has been comprehensively investigated by stopped-flow spectrophotometry. Global analysis of the measurement data using an extended kinetic model which includes all possible pathways for CO2(aq) in piperazine solutions resulted in the complete set of rate and equilibrium constants, including the protonation constants of the mono- and dicarbamate species. CO2(aq) reacts via the free piperazine at high pH in piperazine solutions while the lower kinetic constants for the remaining pathways testifies to a lower contribution to the overall kinetics, except at higher CO2 loadings and lower pH conditions where the concentrations of the respective reactive piperazine species are correspondingly higher.

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AUTHOR INFORMATION

ASSOCIATED CONTENT

S Supporting Information *

Absorbance traces (Figures S1−S3) and rate constants (Table S1) and protonation constants (Table S2). This material is available free of charge via the Internet at http://pubs.acs.org. 812



Article

(19) Wang, X.; Conway, W.; Fernandes, D.; Lawrance, G.; Burns, R.; Puxty, G.; Maeder, M. J. Phys. Chem. A 2011, 115, 6405−6412. (20) Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M. J. Phys. Chem. A 2011, 115, 14340−14349. (21) Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M. Environ. Sci. Technol. 2012, 46, 7422−7429. (22) Conway, W.; Wang, X.; Fernandes, D.; Burns, R.; Lawrance, G.; Puxty, G.; Maeder, M. Environ. Sci. Technol. 2012, DOI: 10.1021/ es3025885. (23) Maeder, M.; Zuberbühler, A. D. Anal. Chem. 1990, 62, 2220− 2224. (24) Maeder, M.; Neuhold, T. M. Practical Data Analysis in Chemistry; Elsevier: Amsterdam, The Netherlands, 2007. (25) Rayer, A. V.; Sumon, K. Z.; Henni, A.; Tontiwachwuthikul, P. Energy Procedia 2011, 4, 140−147. (26) Bougie, F.; Lauzon-Gauthier, J.; Iliuta, M. C. Chem. Eng. Sci. 2009, 64, 2011−2019. (27) Gordesli, F. P.; Alper, E. Int. J. Global Warm. 2011, 3, 67−76. (28) Puxty, G.; Rowland, R. Environ. Sci. Technol. 2011, 45, 2398− 2405. (29) Derks, P. W. J.; Hogendoorn, K. J.; Versteeg, G. F. J. Chem. Eng. Data 2005, 50, 1947−1950.

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Reactions of CO2 with Aqueous Piperazine Solutions: Formation and

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