CO2 Absorption into Aqueous Solutions Containing 3

Article pubs.acs.org/IECR

CO2 Absorption into Aqueous Solutions Containing 3‑Piperidinemethanol: CO2 Mass Transfer, Stopped-Flow Kinetics, 1 13 H/ C NMR, and Vapor−Liquid Equilibrium Investigations William Conway,*,† Yaser Beyad,‡ Marcel Maeder,§ Robert Burns,§ Paul Feron,† and Graeme Puxty† †

CSIRO Energy Flagship, Mayfield West, New South Wales 2304, Australia CSIRO Energy Flagship, Clayton, Victoria 3169, Australia § Department of Chemistry, The University of Newcastle, Callaghan, New South Wales 2308, Australia ‡

S Supporting Information *

ABSTRACT: Global efforts to reduce carbon dioxide emissions stemming from the combustion of fossil fuels have acknowledged and focused on the implementation of post combustion capture (PCC) technologies utilizing aqueous amine solvents to fulfill this role. The cyclic diamine solvent piperazine has received significant attention for application as a CO2 capture solvent, predominantly for its rapid reactivity with CO2. A thorough investigation of alternative but simpler cyclic amines incorporating a single amine group into the cyclic structure may reveal further insight into the superior kinetic performance of piperazine and the wider applicability of such cyclic solvents for PCC processes. One such example is the cyclic monoamine 3piperidinemethanol (3-PM). To facilitate the evaluation of 3-PM as a capture solvent requires knowledge of the fundamental chemical parameters describing the kinetic and equilibrium of the reactions occurring in solutions containing CO2 and 3-PM. Additionally, in parallel with the preceding, experimental measurements of CO2 absorption into 3-PM solutions, including mass transfer and vapor−liquid equilibrium measurements, can be used to validate the CO2 absorption performance in 3-PM solutions and compared to that of monoethanolamine (MEA) under similar conditions. The present study is focused in two parts on (a) determination of fundamental kinetic and equilibrium constants via the analysis of stopped-flow kinetic and quantitative equilibrium measurements via 1H/13C nuclear magnetic resonance (NMR) spectroscopy and (b) experimental measurements of CO2 absorption into 3-PM solutions via wetted wall column kinetic measurements, vapor−liquid equilibrium measurements, and corresponding physical property data including densities and viscosities of the amine solutions over a range of concentrations and CO2 loadings. Fundamental kinetic rate constants describing the reaction of CO2 with 3-PM are significantly faster than MEA at similar temperatures (3-PM = 32 × 103 M−1 s−1, extrapolated to 40 °C from kinetic data between 15.0 and 35.0 °C; MEA = 13 × 103 M−1 s−1, 40 °C). Conversely, the equilibrium constants describing the reaction between bicarbonate and amine, often termed carbamate stability constants, are significantly lower for 3-PM than MEA at similar temperatures. Overall CO2 absorption rates in 3.0 M solutions of 3-PM and MEA, assessed in overall CO2 mass transfer coefficients, are lower in the former case over the entire range of CO2 loadings from 0.0 to 0.4 mol of CO2 per mol of amine. The reduced absorption rates in the 3-PM solutions can be attributed to higher solution viscosities and thus corresponding reductions in CO2 diffusion. CO2 absorption and cyclic capacities in 3.0 M solutions of 3-PM and MEA were found to be significantly higher in the case of 3-PM. The larger CO2 capacities are attributed to the lower stability 3-PM carbamate and the formation of larger amounts of bicarbonate compared to MEA. Overall, the larger CO2 absorption capacity, cyclic capacity, and rapid kinetics with CO2 position 3-PM as an attractive CO2 capture solvent.

1. INTRODUCTION

Currently, the most viable and technologically capable process which can effectively achieve the required and immediate reductions in CO2 emissions from fossil fuel power generation is post combustion CO2 capture (PCC) with subsequent storage of the captured CO2 gas in geological formations. Alternatively, the capture CO2 gas can be utilized for the purposes of enhanced oil recovery3 or as feedstock in the manufacture of high value chemicals.4 An overview of the current status of PCC technology indicates two common roadblocks: large capital costs relating to the manufacture of the

The generation and release of carbon dioxide stemming from the combustion of fossil fuels for the production of electricity, and the potential impacts (of these emissions) on the global climate system, are among the most significant environmental challenges of the modern era.1 It is generally acknowledged that the polluting processes contributing to these underlying emissions must be addressed in order to ensure future threats, impacts, and disruptions, to the environment and society, are minimized.2 In response to these issues, a conscious global research effort is underway targeting methods which are positioned to contribute to significant and immediate reductions in the evolution of parasitic greenhouse gas emissions. © 2014 American Chemical Society

Received: Revised: Accepted: Published: 16715

August 11, 2014 September 9, 2014 September 22, 2014 September 22, 2014 dx.doi.org/10.1021/ie503195x | Ind. Eng. Chem. Res. 2014, 53, 16715−16724

Industrial & Engineering Chemistry Research

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absorption equipment and large and parasitic energy requirement for the PCC process which demands up to ∼30% of the base-load output of a power station when PCC is deployed (devoted to the operation of the capture equipment based on a benchmark technology employing 30% (w/w) monoethanolamine solvent). 5−7 Any industrial realization of PCC technologies will require improvements in the practical design, size, and operation of current-generation PCC processes. Of these, the simplest improvements to date have arisen from the investigation of the chemical solvent. Higher efficiency and rapid reacting solvents are favored in view of replacing typical capture solvents such as monoethanolamine (MEA), methyldiethanolamine (MDEA), 1-amino-2-methyl-1-propanol (AMP), ammonia (NH3), and piperazine (PZ).8,9 The latter example involving PZ, a cyclic diamine, has been extensively investigated for its rapid reactivity with CO2.10,11 While kinetically attractive, several limitations are apparent including low aqueous solubility, limited CO2 capacity at low PZ concentrations, and issues relating to operation of the solvent at high concentrations due to precipitation. Despite the promising kinetic potential of PZ, the preceding issues indicate that alternatives are required.11,12 In view of the promising kinetic performance of PZ a selection of alternatives can be derived from the similar subfamily of “cyclic monoamines” which incorporate aspects of the parent piperazine structure. In a recent screening study13 in which a series of 76 amines were examined for their CO2 absorption rates and absorption capacities, the cyclic monoamine 3-piperidinemethanol (3-PM) was identified as an interesting and promising CO2 capture solvent. 3-PM shares a cyclic structure similar to that of PZ incorporating a single amine moiety into the ring structure; however, unlike PZ, a methanol group is additionally attached to the ring in the case of 3-PM. In view of the large absorption capacity and CO2 absorption behavior, the preceding work has been extended here to include a fundamental investigation of the chemical reactivity in 3-PM solutions and investigations of CO2 absorption in 3-PM solutions via wetted wall covering a range of concentrations and CO2 loadings. 1.1. Outline. CO2 absorption into a series of 3-PM solutions from 0.1 to 3.0 M and CO2 loadings in a 3.0 M 3PM solution from 0.0 to 0.4 mol of CO2 per mol of amine have been investigated using a wetted wall column gas−liquid contactor at 40 °C. CO2 solubility in a 3.0 M 3-PM solution from 40 to 80 °C has been investigated using a vapor−liquid equilibrium apparatus. The fundamental kinetic reactivity of 3PM in the absence of physical mass transfer phenomena, and in equilibrium solutions, has been investigated using rapid stopped-flow kinetic and 1H/13C NMR spectroscopic equilibrium measurements, respectively. 1.2. Kinetic/Equilibrium Reaction Set. The absorption of CO2 in aqueous amine solutions incorporates consecutive contributions from the physical dissolution of CO2 (gas to liquid phase) together with chemical reactions occurring in the liquid phase which act to chemically consume dissolved CO2(aq). The former process is described and quantified by the temperature-dependent Henry’s constant in HCO2 =

pCO

2

cCO2

cCO2 is the dissolved concentration of CO2 in the liquid (aqueous) phase.14,15 Once dissolved into solution, a series of parallel reactions ensue between CO2(aq) and water (H2O), hydroxide (OH−), and amine (R1R2NH where R1 and R2 are typically side groups made up of carbon/hydrogen/oxygen/sulfur/nitrogen atoms in the case of primary and secondary amines) to form carbonic acid (H2CO3), bicarbonate (HCO3−), carbonate (CO32−), and carbamic acid/carbamate (R1R2NCO2H/R1R2NCO2−), respectively. The kinetic reactions are described in eqs 2−4. k1

CO2 (aq) + H 2O XooY H 2CO3 k −1

(2)

k2

CO2 (aq) + OH− XooY HCO3− k −2

(3)

k7

R1R 2NH + CO2 (aq) XooY R1R 2NCO2 H k −7

(4)

Equilibrium reactions describing the protonation(s) of chemical species in solution including those of hydroxide (OH−), carbonate (CO32−), bicarbonate (HCO3−), carbamate (R1R2NCO2−), and amine (R1R2R3N) are described in eqs 5−9. K3

CO32 − + H+ ↔ HCO3− K4

HCO3− + H+ ↔ H 2CO3 K5

OH− + H+ ↔ H 2O K6

R1R 2NH + H+ ↔ R1R 2NH 2+ K8

R1R 2NCO2− + H+ ↔ R1R 2NCO2 H

(5) (6) (7) (8) (9)

An alternative pathway leading to the formation of the carbamate exists for primary and secondary amines via a reaction with HCO3− as described in k9

R1R 2NH + HCO−3 XooY R1R 2NCO2− ( +H 2O) k −9

(10)

The kinetics of this reaction has also been observed by 1H NMR spectroscopy.16,17 The significantly slower rate of this reaction in competition with reactions of CO2 with hydroxide and amine (eqs 3 and 4) was found to have a negligible impact on the analysis of absorption flux data obtained from absorption measurements in a wetted wall contactor and kinetic stopped-flow measurements.18 Thus, the reaction does not appear in the kinetic model for the analysis of stopped-flow and wetted wall measurements. However, the equilibrium constant for the pathway is typically used to describe the stability of carbamates and appears exclusively in the model for the analysis of 1H NMR data here. 1.3. Chemical Constant Data and Activity Coefficient Corrections. Values for the kinetic and equilibrium constants for the reactions described in eqs 2, 3, and 5−8 were taken from the literature.19−21 Kinetic constants for the formation of the carbamic acid in eq 4, and the carbamate stability constant in eq 10, were regressed from the stopped-flow kinetic and NMR equilibrium measurements here, respectively. Values for the protonation constant of the carbamic acid, log K8, eq 9, were determined in the kinetic analysis using the equilibrium

= 2.82 × 106e−2044/ T (1)

where HCO2 represents the Henry’s constant for CO2 solubility in pure water at infinite dilution (kPa M−1), p is the partial pressure of CO2 in the gas (vapor) phase above the liquid, and 16716

dx.doi.org/10.1021/ie503195x | Ind. Eng. Chem. Res. 2014, 53, 16715−16724


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constant data for eq 10, log K9, determined in the independent 1 H NMR analysis, and the equilibrium constants log K1, log K4, and log K7 (calculated as K7 = k7/k−7) from eqs 2, 6, and 4, respectively, via microscopic reversibility as K8 = K7K4/K1K9. Activity coefficient corrections were applied in the calculations using the Debye−Hückel approximation and a reduced version of the specific interaction theory (SIT) for charged species. The former was used in the analysis of fundamental kinetic and 1H NMR measurement data while the latter was used in the analysis of wetted wall column and vapor−liquid equilibrium (VLE) measurements. All calculated values correspond to standard state conditions at infinite dilution at zero ionic strength. The Debye−Hückel and SIT expressions are shown in eqs 11 and 12.

log γi =

log γi =

respectively. The chemical shift of the TSP standard was assumed constant and unaffected by temperature over the narrow range of temperatures investigated in the NMR measurements here. All samples were tightly capped and sealed with para film tape so as to minimize any loss of CO2 gas from the samples during equilibration and measurement. The concentrations were chosen here to ensure the solution CO2(aq) concentration did not exceed the solubility at normal pressure. Solutions were thermostated in a temperature controlled water bath until equilibrium was completely reached (determined when no further change in the spectrum was observed). The set point temperature was maintained internally within the instrument (±1 °C) during the acquisition of spectra. 2.2.2. Spectral Processing and Peak Integration. 1H and 13 C NMR spectra were processed individually using Bruker Topspin 3.0.1.b software. A baseline correction (polynomial) was applied to all spectra prior to the integration of peaks. Spectral peaks corresponding to amine and carbamate were integrated in the topspin software, and the total concentrations of 3-PM and its carbamate calculated from the total concentrations of amine ([3-PM] tot ) and bicarbonate ([HCO3−]tot) in the solution. The average integrals of 13C signals for the aliphatic carbons in 3-PM and its carbamate were used in the equilibrium analysis of the low concentration samples at 25.0 °C for comparison to the 1H data. 2.2.3. NMR 3-PM/Na2CO3/HCl Titration Samples. A series of equilibrium samples were prepared in 5 mm diameter NMR tubes containing the TSP/D2O capillary insert by dosing various amounts of a concentrated HCl solution (4.0 M) into 1.5 mL of an equilibrated carbamate solution initially containing 0.5 M 3-PM and 1.0 M Na2CO3. The series of additions were designed to cover the approximate range of pH conditions over which carbamates are typically observed (pH 12−8). A duplicate series of samples containing 1.0 M 3-PM and 1.8 M Na2CO3 dosed with various volumes of HCl were prepared and analyzed in parallel. 2.3. Stopped-Flow Kinetics. Fast reaction kinetics were performed on an Applied Photophysics DX-17 stopped-flow spectrophotometer equipped with a J&M Tidas MCS 500-3 diode array detector.22 The absorbance changes of acid−base indicators, corresponding to changes in the pH of the solutions upon reaction of solutions in the stopped flow, were observed over the wavelength region from 400 to 700 nm. All solutions were thermostated and maintained at the set point temperature ((15.0−35.0) ± 0.1 °C) by a circulating Julabo F20 water bath. The exact temperature of the solution was monitored by a thermocouple located inside the stopped-flow absorption cell block. 2.3.1. Reaction of CO2(aq) with 3-PM and Decomposition of 3-PM Carbamate. First, the formation of 3-PM carbamate was investigated by reacting equal volumes of a CO2(aq) solution (4.35 mM), with a series of 3-PM solutions ([3-PM]0 = 1.0−10.0 mM), in the presence of 0.05 mM alizarin red S and 12.5 μM thymol blue indicators, in the stopped flow. Initially, CO2(aq) solutions were prepared by bubbling a mixed gas containing CO2 and N2 into a water solution located in a small sample reservoir above the drive syringe in the stopped flow. The initial concentration of CO2(aq) in this solution was determined from the CO2 partial pressure in the gas stream (adjusted by gas flow controllers), and the corresponding Henry’s constant for CO2 solubility in water. Ideal behavior and

Az i 2 μ 1+

μ

(11)

Az i 2 μ 1 + 1.5ρ−1/2 μ

(12)

where γi is the activity coefficient, A = (1.8248 × 10 )/ (eT)3/2 is the Debye−Hückel law slope, ρ is the density of water (kg dm−3), μ is the ionic strength (mol dm−3), and zi is the charge of species i. 6

2. EXPERIMENTAL SECTION 2.1. Chemicals. High-purity carbon dioxide gas (CO2, BOC Gases Australia, 99.99% purity), N2 (Coregas, 99.99%), potassium bicarbonate (KHCO3−, BDH), sodium carbonate (Na2CO3, BDH), potassium hydrogen phthalate (AJAX), sodium hydroxide solid (Merck), 3-piperidinemethanol (3PM, Suzhou Rovathin Pharmatech), and hydrochloric acid (AJAX) were all used as obtained. The concentrations of sodium hydroxide and 3-PM were standardized by potentiometric titrations. A stock solution of sodium hydroxide was initially titrated with potassium hydrogen phthalate (KHP) standard, and this solution was used to standardize the concentration of the hydrochloric acid stock solution. A known amount of the standardized hydrochloric acid solution was added to the 3-PM solution and back-titrated to high pH with standardized sodium hydroxide solution. All solutions were prepared using ultrahigh purity milli-Q water which was further boiled to remove dissolved CO2 gas. 2.2. 1H/ 13C NMR Spectroscopy. 2.2.1. Acquisition. Quantitative proton (1H) and carbon (13C) NMR spectra were acquired on a Bruker Avance Ascend 600 NMR spectrometer operating at a frequency of 600.213 MHz. Proton 1 H spectra were obtained as the average of 16 scans with a pulse delay time (D1) of 1 s. Inverse-gated decoupled 13C spectra were obtained at a pulse angle of 30° (zgig30 pulse program, Bruker) as the average of 128 scans with a pulse delay time (D1) of 60 s, corresponding to a value ∼5 times that of the T1 relaxation time of the slowest relaxing carbon (typically carbamate/HCO 3 − ). Additional one-dimensional (1-D) DEPT135 13C and two-dimensional (2-D) HSQC spectra were measured and utilized in the assignment of spectral peaks corresponding to 3-PM and its carbamate. These spectra are available in Figures S1 and S2 in the Supporting Information. TSP (3-(trimethylsilyl)propionic acid-d4, sodium salt) dissolved in D2O in a sealed glass capillary insert was present in the NMR tube acting as the external reference and locking agents, 16717



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logarithmic mean of the inlet and outlet CO2 partial pressures. The linear slope of the NCO2 against PCO2 plot is equal to the overall mass transfer coefficient, KG. The interested reader is directed to Darde et al.24 and Puxty et al.18 for a comprehensive description of the relationships between CO2 diffusion and chemical reaction during absorption of CO2 into an amine liquid and the mathematical principles of the absorption processes. 2.5. Density and Viscosity. Density and viscosities of the amine solutions were determined using a combined Anton Parr DMA-38 density meter (±0.001 g/mL) and AMVn viscometer, at each of the required temperatures. Density and viscosity measurements were repeated in triplicate, and the resulting value is the average of these repeats. Density and viscosity data as a function of 3-PM and MEA concentrations from 0.5 to 6.0 M, and in 3.0 M 3-PM and MEA solutions and CO2 loadings from 0.0 to 0.4 mol CO2/mol amine, are presented in Figures S3−S6 in the Supporting Information. 2.6. Vapor−Liquid Equilibrium. CO2 solubilities in 3.0 M solutions of 3-PM and MEA were determined using a Parr VLE apparatus from 40 to 80 °C.25 The equilibrium distribution of CO2 between gas and liquid phases in each of the amine solutions was investigated by sequentially dosing CO2 gas from a Swagelok CO2 cylinder, into a known volume of amine solution (100 mL), contained in a sealed and thermostated glass reaction vessel (Parr). The precise amount of CO2 dosed into the vessel was determined by variation in mass of the CO2 cylinder before and after dosing, which was suspended from an analytical balance (Mettler Toledo pb4002-s). Initially, and prior to CO2 dosing, the vessel was repeatedly charged with ultrapure N2 gas to ensure all gas impurities were removed from the gas lines and gas headspace of the vessel. Following the N2 flush, stirring of the amine solution was initiated and maintained at ∼500 rpm until steady state conditions were reached as indicated by constant gas pressure reading (Swagelok S model transducer) and constant temperature (thermocouple). Total gas phase pressure and temperature in the reactor were monitored throughout the measurements following each dosing of CO2, until steady state/equilibrium was again reached. Once steady, the proceeding CO2 dosing was initiated. From the total reactor pressure at each steady state, and CO2 dosing, the total amount of CO2 remaining in the gas phase, subtracting contributions to the total reactor pressure from water vapor and initial N2, and the subsequent CO2 loading in the amine liquid were calculated according to eqs 13−15

uniform mixing of the gases in the gas stream was assumed here. Second, the decomposition of 3-PM carbamate at low pH conditions was investigated by reacting a preequilibrated carbamate solution with hydrochloric acid solutions in the stopped flow. Initially, a carbamate solution containing 0.025 M 3-PM and 0.05 M HCO3− was prepared and equilibrated for 24 hours at the desired temperature. The composition of the equilibrated carbamate solution was determined quantitatively using 1H NMR spectroscopy prior to the stopped-flow measurements. This equilibrated carbamate solution was further reacted in the stopped flow with a series of hydrochloric acid solutions to establish the decomposition reaction kinetics. Absorbance changes of 0.05 mM alizarin red S and 0.025 mM methyl orange indicators, and subsequently changes in the pH of the solutions due to the decomposition of 3-PM carbamate, carbonic acid, and bicarbonate, and protonation of amine, were monitored. The preceding series of reactions were repeated over a range of temperatures from 15.0 to 35.0 °C to establish the temperature dependence of the reactions. Measurements were repeated a minimum of four times to accommodate a statistical analysis of the regressed kinetic and equilibrium parameters which were compared to the standard deviations produced in the software. Analysis of the kinetic and equilibrium data was performed using ReactLab Kinetic and Equilibrium software packages (www.jplusconsulting.com) and in-house extensions of the software written in Matlab. 2.4. Wetted Wall Column. Absorption of CO2 into individual aqueous solutions of 3-PM and MEA was performed using a wetted wall column gas−liquid contactor at 40 °C. Briefly, a stainless steel column with an effective height and diameter of 8.21 cm and 1.27 cm, respectively, was used here.23 CO2 absorption flux, NCO2, into approximately 0.6 L of amine solution was measured over a range of CO2 partial pressures spanning 1.0−20.0 kPa. The total liquid flow in the apparatus was maintained at 121.4 mL·min−1 (2.02 mL·s−1) by a Masterflex peristaltic pump. The liquid flow was monitored using a calibrated liquid flow meter. The total gas flow rate in the system was maintained at 5.0 L min−1, and the desired composition of the gas (i.e., CO2 partial pressure) established by variation of Bronkhorst mass flow controllers, calibrated for CO2 and N2 gases, respectively. The composition of the gas stream entering and exiting the wetted wall column, expressed in vol % CO2, was continuously monitored using a Horiba VA3000 IR gas analyzer. The former was measured at each CO2 partial pressure while bypassing the wetted wall column with the gas stream passing directly to the Horiba gas analyzer. 2.4.1. CO2 Absorption Flux (NCO2) and Overall Mass Transfer Coefficients (KG). The amount of CO2 absorbing into the amine liquid was determined from the relative amount of CO2 in the gas stream entering (bottom) and exiting (top) the wetted wall column. The CO2 absorption flux, expressed as millimoles of CO2 absorbed per second, per unit area (mmol· s−1·m−2) of contact between liquid amine and gas, was determined over a range of CO2 partial pressures in each of the amine solutions. The preceding procedure was repeated for each of the amine solutions including those preloaded with different amounts of CO2. Overall mass transfer coefficients, denoted here as KG, were determined from plots of the absorption flux, NCO2, against the applied driving force, PCO2. The latter was determined as the

nCO2(gas) = Z

(Pvessel(tot) − PH2O(vap) − PN2)Vgas RT

(13)

nCO2(liquid) = nCO2 (dose) − nCO2(gas)

(14)

loading CO × α = nCO2(liquid)/namine

(15)

2

where nCO2(gas) is the moles of CO2 gas, Z is the compressibility of CO2, Pvessel(tot) is the total pressure as measured inside the vessel at steady state, PH2O(vap) is the vapor pressure of water, PN2 is the partial pressure of nitrogen, Vgas is the volume of the gas phase (total vessel volume−amine liquid volume in m3), R is the gas constant (8.314 J·(mol·K)−1), T is the temperature (K), nCO2(dose) is the moles of CO2 dosed, nCO2(liquid) is the moles of CO2 dissolved into the amine solution, and namine is the 16718



16719

98(3) 41(8) 63(2)

1.17× 10 76(3)

6.6(1) 6.4(1) 6.3(5) log (K8/M−1) MEA

a

K8

RNHCO2− + H+ ↔ RNHCO2 H

3-PM MEA 3-PM

MEA

16.8(2) 2.7(1) × 103 35(1) 1.6(3) × 102 76(5) 7.52(3) k−7 (s−1) k7 (M−1 s−1) k−7 (s−1) K7 (M−1) K7 (M−1) log (K8/M−1)

73(5) 4.9(2) × 103 90(8) 1.1(1) × 102 54(4) 7.28(2)

2.4(3) × 102 7.6(4) × 103 1.9(3) × 102 81(17) 41(6) 7.06(4)

1.98(5) × 10 8.2(5) × 10 2.6(4) × 10 s )

k7 (M 3-PM k7

k−7

RNH 2 + CO2 (aq) XoooY RNHCO2 H

k denotes rate constant; K denotes equilibrium constant and is calculated as Ki = ki/k−i. Numbers in parentheses represents standard deviation in the last digit. MEA data included for comparison. bk(T) = A exp(−Ea/RT). ck(T) = (k′T/h) exp((ΔH⧧ − TΔS⧧)/RT). dK(T) = exp(−(ΔH°/RT) + (ΔS°/R)).

17(1) 6.3(2) × 10−3

17

−40(10) −43(7) 81(12) 96(3) 38(1) 61(2) 1.15 × 1019 5.8× 1010 1.0× 1013

73(3)

74(8)

112(8) −47(3) −4(7)

−23(3) −23(2) −19(3)

ΔH° (kJ mol−1) ΔS⧧ (J mol−1 K)

17

A Ea

4

35.0 °C

3

25.0 °C 15.0 °C

3

−1 −1

constant reaction

Arrheniusb

Table 1. Kinetic and Thermodynamic Data Describing the Reversible Formation of 3-PM Carbamate/Carbamic Acida

Eyringc

3. RESULTS AND DISCUSSION 3.1. Kinetics. 3.1.1. Stopped-Flow Studies. The complete series of kinetic measurements, including measurements of carbamate formation and decomposition, were analyzed in a single global fit at each temperature. The resulting analysis produced kinetic and equilibrium constants for the reversible formation of 3-PM carbamate, k7 and k−7, eq 4, and the protonation constant of 3-PM carbamate, K8, eq 9, at each temperature. Corresponding equilibrium constants for the carbamate pathway in eq 4 were also calculated from the ratio of the formation and decomposition rate constants as K7 = k7/k−7. Subsequent analysis of the kinetic constants and their temperature dependencies using the Arrhenius, Eyring, and van’t Hoff relationships resulted in standard activation energies, Ea, enthalpies, ΔH⧧/ΔH°, and entropies, ΔS⧧/ΔS°, of the reactions, respectively. The calculated kinetic and thermodynamic data are presented in Table 1. From the data provided in Table 1 all kinetic values were found to increase with temperature. A comparison of the kinetic constants, k7, and the equilibrium constant, K7, for 3-PM and MEA reveals consistently larger values for 3-PM at similar temperatures. The underlying reason for this trend relates to the relative basicities of the amines, 3-PM being ∼1.0 unit larger than MEA at similar temperatures. We have previously investigated the intimate relationships26 between the kinetic and equilibrium constants of carbamates and the protonation constant of the amine. It was typically found that amines with larger protonation constants resulted in large kinetic constants in interactions with CO2. The resulting data for 3-PM here are consistent with this trend where the larger protonation constant of 3-PM correspondingly renders it as a faster reacting amine by promoting stronger interactions with CO2. This is largely due to the apparent size of the lone electron pair on the nitrogen (amine group) which, in turn, increases the effective nucleophilic reactivity of the amine group toward CO2 molecules. Conversely, the stability of 3-PM carbamate, as outlined in the equilibrium constants log K9, is substantially lower than the corresponding MEA carbamate. Molecular explanations for the reduced carbamate stability in the case of 3-PM are complex and may reside from a combination of electronic and steric contributions, the latter forming the most probable explanation where weak steric hindrance imposed by the relatively close proximity of the large and bulky ethanol tail to the amine group, and thus the carbamate when formed, may dissociate under particular conditions. Furthermore, it has been suggested that the formation of intramolecularly hydrogen bonded ring structures may contribute to decreased carbamate stability.16 The Arrhenius activation energy for 3-PM determined here is significantly larger than the corresponding value for MEA and some other aliphatic amines, the values of which are typically in the range of 40−65 kJ/mol. Chemical explanations for the large values are unclear, however activation energies reported in the literature for the cyclic amines such as piperdine and pyrrolidine, 75.9 and 79.9 kJ/mol, respectively, are similarly larger and close to the activation energies reported here for 3-PM.27

ΔH⧧ (kJ mol−1)

van’t Hoffd

moles of amine in solution. Given the large molecular weight and polarity of 3-PM, and the presence of the polar hydroxyl group which can participate in hydrogen bonding interactions, it was assumed that the amine vapor pressure is too low to provide an observable contribution to the total measured pressure.

ΔS° (J mol−1 K)

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3.1.2. CO2 Absorption Flux, NCO2, and Overall Mass Transfer Coefficient(s), KG. Overall mass transfer coefficients are critical parameters of amines in the assessment of their CO2 absorption performance. Typically, KG values relate directly to the absorption area and packing requirements for the capture process. Overall CO2 absorption rates into 3-PM and MEA solutions over a range of amine concentrations and CO2 loadings from 0.0 to 0.4 mol of CO2 per mol of amine were investigated here at 40 °C using a wetted wall column. Although CO2 absorption into MEA solutions has been extensively investigated, a spread in published data relating to experimental variations justifies the repeated measurement of the data here to allow a consistent comparison. Due to the limited solubility of 3-PM in aqueous solution the upper concentration was limited to 3.0 M. Similarly, measurements up to 3.0 M for MEA were completed and data from 3.0 to 5.0 M were taken from our previous study and used in the comparisons here. It should be noted that segments of nonlinearity were observed in the CO2 flux data for 3-PM. In such cases, only the linear region of the flux curve was considered in the calculation of KG values. 3.1.2.1. Effect of 3-PM and MEA Concentration(s) on KG Values. KG data for CO2 absorption into 3-PM and MEA solutions as a function of amine concentration, at zero CO2 loading, are presented in Figure 1. Despite larger kinetic

attributable to the fact that the viscosity increases by a larger extent as the 3-PM concentration is increased, relative to MEA. This increased viscosity inhibitively acts to lower the diffusion coefficients of species in the liquid phase which, in turn, offsets the kinetic benefit of 3-PM and the larger reaction rate constant in its reaction with CO2. Given the considerable structural variations between 3-PM and MEA, specifically the large rigid cyclic ring structure in 3-PM, as opposed to the small free rotating chain structure of MEA, the observed increases in KG are easily justified. It should be emphasized that while it appears the amine concentration has only a minor impact on KG at higher amine concentrations (plateau region in Figure 1), considerable advantages in terms of the overall absorption uptake and CO2 capacity when operating the solvent at higher amine concentrations justifies the pursuit of more concentrated amine solutions. 3.1.2.2. Effect of CO2 Loading on KG in 3.0 M Amine Solutions. A comparison between the CO2 absorption performance in 3-PM and MEA solutions, including solutions over a range of CO2 loadings, was performed here. Selected KG data for individual 3.0 M solutions of 3-PM and MEA and for CO2 loadings from 0.0 to 0.4 mol of CO2 per mol of amine are presented in Figure 2. Overall, CO2 absorption is slower in the

Figure 1. Overall mass transfer coeffeicients at 40 °C, KG, in 3-PM and MEA solutions as a function of amine concentration.

Figure 2. Overall mass transfer coefficients at 40 °C, KG, for 3.0 M solutions of 3.0 M 3-PM and MEA and for CO2 loadings from 0.0 to 0.4 mol of CO2 per mol of amine.

constants for the reaction of CO2(aq) with 3-PM, the corresponding trend from the stopped-flow data where 3-PM is reacting significantly faster than MEA is reversed in the KG data. From Figure 1, CO2 absorption rates were found to be comparable at low amine concentrations up to 1.0 M; however the trend in KG increases more rapidly with increasing amine concentration in the MEA solutions beyond 1.0 M. Interestingly, a plateau in KG is observed along the sequences at higher amine concentrations where KG increases only marginally with progressive increases of the amine concentration above 2.0 and 3.0 M for 3-PM and MEA, respectively. Concomitant changes in the physical properties of the amine solutions as the amine concentration is increased are the dominant causes of the effect. Any significant increases in KG with increasing amine concentration above the aforementioned concentrations are counterbalanced by parallel increases of the solution viscosities which impact diffusion related processes occurring in the solution. Further consideration should be given to the preceding trends with respect to the individual amines. The slower overall CO2 mass transfer in the 3-PM solutions is

3-PM solutions over the entire range of CO2 loadings. Although the KG values decrease with increasing CO2 loading, consistent with decreasing amounts of the free reactive amine for reactions with CO2 in solution, the magnitude of the decrease is reasonably constant and linear with increasing CO2 loading for both 3-PM and MEA. The observed decrease in KG with increasing CO2 loading is induced by the changing chemical and physical properties of the solutions with increasing CO2 loading; first the reduced availability of free amine and second changes in the physical properties of the solutions resulting from the increased concentration of ionic species (carbamate, bicarbonate, carbonate, and protonated amine) and the ensuing interactions of these species. From a theoretical perspective, and in order to compete with MEA at similar concentrations and CO2 loadings, it is possible that the 3-PM solvent could be operated at slightly elevated temperatures which may result in a corresponding increase in the CO2 absorption rate. While the prospect of larger KG values at higher temperatures is attractive, the effect on diffusion and CO2 solubility may, in parallel, impede such efforts. Similar 16720



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determination of fundamental chemical constants and their application to the prediction of realistic absorption flux data are broadly demonstrated. 3.2. Equilibrium Studies. The equilibrium performance and CO2 capacity are critical properties in the selection of an amine solvent which ultimately defines the overall energy demand of the desorption step. Two approaches are typically employed for the investigation of equilibrium behavior in amine solutions; fundamental investigations of chemical equilibrium constants relating to the formation of carbamates and direct measurements of CO2 solubility in a batch type reactor. The former reveals intimate information about the energy related to the formation (and destruction) of the carbamate (C−N bonds). Such information is additionally and fundamentally useful in the selection of novel solvents and in the prediction of equilibrium species distribution in CO2−amine solutions. Direct measurements of the equilibrium CO2 absorption capacity, as a function of increasing CO2 loading, based on measurements of the equilibrium CO2 partial pressure in the gas phase behavior, are principally simpler but do not carry information about the chemical speciation in the solutions. We have investigated both of these aspects here using 1H/13C NMR spectroscopy and direct measurements of CO2 solubility via VLE measurements. The results of these studies are presented and discussed in the following. 3.2.1. 1H/13C NMR Investigations: Carbamate Stability Constants. The robust determination of equilibrium constants describing the formation of carbamates, also commonly termed carbamate stability constants, ideally requires a series of measurements covering a range of pH conditions where the carbamate is present in measurable and quantifiable amounts. Simpler investigations which involve the extraction and analysis of liquid samples from equilibrated amine solutions via NMR, the solutions of which are typically prepared with increasing amounts of CO2 sparged into the amine solutions, are prone to significant error. In favor of this approach we have chosen to perform a series of acid−base titrations of solutions containing 3-PM and Na2CO3, with various amounts of hydrochloric acid, to establish a series of samples covering the typical range of pH conditions where the carbamate is present and at a fixed amine concentration and CO2 loading. This approach allows the accurate preparation and delivery of both the amine and CO2 (as CO32− initially) to the sample without introducing small but often significant errors associated with evaporative losses (both solvent and water) during the CO2 bubbling and liquid sampling processes and gaseous loss of CO2 from the samples. In an effort to establish the existence of any underlying effect(s) of concentration on the resolution and sensitivity of the NMR technique for the analysis of CO2/amine systems, we have prepared duplicate series of samples at both low and high amine and CO32− concentrations, respectively. Additionally, selected samples have been analyzed in parallel via 13C measurements at 25.0 °C and the resulting integrals of the

arguments can be made for the operation of MEA at elevated temperatures; however the penalty to the absorption capacity, which is lower at higher temperatures, does not justify the pursuit for greater mass transfer rates. Conversely, given the larger absorption capacity of 3-PM, it can potentially sustain such a trade-off between larger KG values and CO2 capacity. 3.1.3. CO2 Mass Transfer Predictions. Experimental absorption flux and mass transfer data have been predicted using the kinetic (rate) and equilibrium constants describing the reaction of CO2 with 3-PM from the independent stoppedflow kinetic study. A simple extrapolation of the kinetic constants to 40 °C via the Arrhenius expression was conducted in the absence of measured values. An in-house software tool developed in Matlab was used for the prediction of molar CO2 absorption fluxes occurring in the amine film during absorption based on a comprehensive diffusion and chemical reaction model.18,28,29 In the absence of measured values, diffusion coefficients and CO2 solubility in the solutions were calculated using correlations for these properties and the measured density and viscosity data.29 Experimental CO2 absorption data in the 3-PM solutions were predicted with the overall agreement between the measured and predicted absorption fluxes assessed in a parity plot as presented in Figure 3.

Figure 3. Parity plot of measured and calculated absorption flux in 3PM solutions at 40 °C and zero CO2 loading.

The agreement between measured and model calculated flux data is satisfactory in the amine solutions, in the absence of preloaded CO2, over the entire range of 3-PM concentrations. The average absolute deviation (AAD) over the entire series is 14.3%. Considering the dynamic nature of the experimental measurements where small variations in the expected CO2 loading of the solution from the mass difference method translate into significantly large deviations in the modeling results, and the simple estimations for the film properties (film thickness, exposure time, and diffusion coefficients) in the model, the result can be considered satisfactory. Overall, the

Table 2. Equilibrium Constant Data for the Formation of 3-PM Carbamate Determined from the Analysis of 1H NMR Measurements at 15.0−45.0 °C van’t Hoff reaction K9

RNH 2 + HCO3− ↔ RNHCO2− (+ H 2O)

constant log (K9/M−1)

15.0 °C

25.0 °C

35.0 °C

45.0 °C

ΔH° (kJ mol−1)

ΔS° (J mol−1 K)

1

0.93(1)

0.84(1)

0.76(1)

0.67(1)

−14.8(1)

−34(1)

1

0.95(1)

0.86(1)

0.78(1)

0.71(1)

−13.9(5)

−30(2)

NMR sample H high concn H low concn

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C spectra compared to the analysis of the corresponding 1H spectra. A typical series of 1H and 13C NMR spectra at 25.0 °C in the equilibrium solutions containing 0.5 M 3-PM and 1.0 M Na2CO3 in the presence of additional amounts of HCl are presented in Figures S7 and S8 in the Supporting Information. Analysis of the concentration data obtained from the integration of the carbamate and amine peaks in the NMR spectrum, together with the initial concentrations of the amine and Na2CO3 in the solutions, were analyzed using an equilibrium chemical model incorporating reactions 5−10. The analysis resulted in values for the carbamate stability constant, K9, at each of the temperatures from 15 to 45 °C. Selected fits of the concentration data are shown in Figure S9 in the Supporting Information. The resulting values for K9 at each of the temperatures, together with corresponding thermodynamic values determined from subsequent analysis of the temperature dependencies (of the equilibrium constants) via the van’t Hoff relationship, are shown in Table 2. The equilibrium constants and temperature dependences are similar for the two series of solutions, confirming the effect of concentration on the determination of equilibrium constants is minimal here. Integrals determined from the analysis of 13C data at 25.0 °C are similar to the corresponding 1H data and were generally within ∼30% of each other. While the difference between the 1H and 13C results may appear somewhat large, the discrepancy relates to the sensitivity of the technique to the 13 C signal and relatively low abundance of the 13C isotope (∼1% of the 12C isotope abundance). Thus, the integration of small peaks close to the signal baseline is complex and prone to slightly larger errors. The case is more relevant for amines such as 3-PM where the carbamate is not present at significant concentrations despite the use of high amine concentrations here. However, both 1H and 13C NMR could be considered suitable for the analysis of CO2/amine solutions where the amount of carbamate formed is reasonably high so as to minimize issues relating to the sensitivity of the NMR signal to carbamate concentration. A significant limitation of the NMR method here is the formation of CO2 gas bubbles resulting from the reduced solubility at low pH. Thus, the investigation of equilibrium behavior at significantly higher, and industrially relevant, amine concentrations is not possible without significant modification of the procedure. While the conditions here result in adequate and quantifiable amounts of carbamate, the formation of considerable amounts of CO2 gas (bubbles) in the solutions at lower pH conditions somewhat restrict the method and analysis. Moreover, for this reason the inclusion of a reaction defining the protonation of the carbamate into the chemical model, K8, was found to have little impact on the quality of the data fitting. Thus, this reaction is essentially not defined in the measurements here. 3.2.1.1. Comparison of log K9 Values with MEA and Other Cyclic Amines. A comparison of the carbamate stability constants and thermodynamic values here with the corresponding values for MEA and a series of cyclic amines in our previous work reveals several similarities.16 The average value (from the low and high concentration NMR series) for log K9 = 0.84 at 25.0 °C here is considerably lower than the corresponding value for the unhindered parent cyclic amine piperidine (log K9 = 1.38) and is aligned with the lower end of the range of values (log K9 = 0.79−2.69) for sterically unhindered cyclic secondary amines at similar temperatures. Considering the stability of 3PM carbamate is potentially influenced by the formation of an intramolecularly hydrogen bonded structure, any further

convolution of this behavior on the magnitude of stability constants is highly complex and beyond the scope of this work. The temperature dependence of log K9 here, expressed in the ΔH° values, is within the range of values for similar cyclic secondary amines we have assessed previously. The use of lower amine concentrations in our previous study, and the observation (or not) of carbamates in similar hindered amines to 3-PM with alcohol groups in close proximity to the reactive amine group, further complicates any in-depth analysis of any trends. 3.2.2. Vapor−Liquid Equilibrium. Direct measurement of CO2 solubility in 3.0 M solutions of 3-PM and MEA at 40.0 to 80.0 °C were performed using a vapor−liquid equilibrium apparatus. CO2 solubility data for a 3.0 M 3-PM solution is presented graphically in Figure 4.

Figure 4. Equilibrium CO2 partial pressure as a function of CO2 loading in a 3.0 M 3-PM solution from 40 to 80 °C. The solid black line corresponds to 15 kPa CO2 partial pressure (15% CO2).

From the data in Figure 4 the gas phase CO2 partial pressure increases marginally with increasing CO2 loading at low loadings and rapidly at high loadings above 0.3 at each of the temperatures. The highest CO2 solubility is observed at low temperatures, which decreases significantly with increasing temperature. Further changes in the CO2 partial pressure with increasing CO2 loading above ∼15 kPa are similar at each of the temperatures. While the quality and definition of the data at low CO2 loadings are somewhat discerning, related to the limited sensitivity of the pressure measurements at very low loadings below 20 kPa, the data presented here are useful for the purpose of defining the approximate absorption capacity of the solvents. Such effects are commonly encountered in VLE measurements. Absorption capacities, expressed on a moles per mole basis, at 40 °C and 15 kPa CO2 partial pressure (from the data), together with estimated cyclic capacities calculated from the absolute difference in CO2 loading from 40 to 80 °C at 15 kPa CO2 partial pressure, are presented in Table 3. Absorption capacity at low temperature is some 20% larger for 3-PM than MEA. The larger capacity is a function of the low carbamate stability in the case of 3-PM which chemically favors the formation of larger amounts of bicarbonate and lower amounts of carbamate at CO2 partial pressures similar to those of MEA. The result of this behavior is perpetually higher CO2 loadings beyond 0.5:1.0 mol of CO2 per mol of amine can be achieved. Moreover, the cyclic capacity of 3-PM over the temperature range here is almost double that of MEA for similar reasons. It is important to emphasize that any realistic comparison of the 16722



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NMR equilibrium plots, DEPT-135 and HSQC spectra, and calculated and measured equilibrium concentration profiles from 1H NMR analysis. This material is available free of charge via the Internet at http://pubs.acs.org.

Table 3. Estimated Absorption and Cyclic Capacities from Experimental CO2 Solubility Measurements in 3-PM and MEA Solutions amine (3.0 M)

absorption capacitya

cyclic capacityb

3-PM MEA

0.63/83c 0.51/110c

0.21/28c 0.11/24c

■

Corresponding Author

Moles of CO2 per mole of amine: 40 °C; 15 kPa CO2 partial pressure. bMoles of CO2 per mole of amine; difference in loading between 40 and 80 °C at 15 kPa CO2 partial pressure. cGrams of CO2 per liter of solvent. a

*E-mail: [email protected]. Tel.: + 61249 606098. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was carried out within CSIRO’s Energy Flagship, and it was supported by the Australian Government through the Australia-China Joint Coordination Group on Clean Coal Technology (JCG). 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.

maximum CO2 loading and cyclic capacities of amine solvents should be performed on a mass basis (grams of CO2 per liter of solvent), particularly in the case of solvents which will be operated at significantly different concentrations as is presented here. The higher chemical efficiency of the 3.0 M 3-PM solvent is captured more effectively on a mass basis, and the overall capacity approaches that in a 5.0 M MEA solution. Furthermore, on a mass basis, the larger cyclic capacity of 3PM results in a larger overall CO2 cycle in 3.0 M 3-PM than in 5.0 M MEA. Thus, a process in which the two solvents would be operated would require similar solvent inventories to achieve comparable cyclic capacities despite the significantly lower concentration in the former solution. This, in turn, results in advantages in terms of the capitol and operating costs related to solvent makeup. Detailed process simulations will reveal the optimum operating temperature and CO2 loading ranges of the solvent.

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REFERENCES

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4. CONCLUSION The fundamental chemical and CO2 absorption behavior of 3piperidinemethanol has been investigated via a series of fast kinetic stopped-flow, 1H/13C NMR equilibrium, wetted wall column, and vapor−liquid equilibrium measurements. The fundamental kinetic constants describing the formation of 3PM carbamic acid are significantly larger when compared to corresponding values for MEA at similar temperatures while carbamate stability constants are lower and thus highly favored. 3-PM offers reasonable overall absorption rates compared to MEA at similar concentrations and loadings. The results here have additionally demonstrated that, in the absence of measured absorption flux data, the values can be reasonably predicted using the fundamental chemical data including rate and equilibrium constants and corresponding physical property data of the amine solutions. Significantly larger equilibrium absorption and cyclic capacities were found for 3-PM over MEA at similar conditions. From a strictly chemical point of view the rapid kinetics and low carbamate stability highlight 3-PM as an interesting candidate for CO2 capture processes. The marginally slower absorption rates position 3-PM as a promising candidate for use in a blended solvent which could potentially increase the solubility of 3-PM and overall CO2 absorption rates. A more comprehensive investigation of cyclic amines should be undertaken to exhaustively determine if any further relationships (of these molecules) to the cyclic amine piperazine exist and what is their overall suitability for PCC processes.

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AUTHOR INFORMATION

ASSOCIATED CONTENT

S Supporting Information *

Figures showing density and viscosity data in 3-PM and MEA solutions and a range of CO2 loadings, stacked 1H and 13C 16723



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