Effect of Water and Temperature on Absorption of CO2 by Amine

ARTICLE pubs.acs.org/JPCB

Effect of Water and Temperature on Absorption of CO2 by Amine-Functionalized Anion-Tethered Ionic Liquids Brett F. Goodrich, Juan C. de la Fuente,† Burcu E. Gurkan, Zulema K. Lopez, Erica A. Price, Yong Huang, and Joan F. Brennecke* Department of Chemical and Biomolecular Engineering, University of Notre Dame, Notre Dame, Indiana 46556, United States

bS Supporting Information ABSTRACT: Amine-functionalized anion-tethered ionic liquids (ILs) trihexyl(tetradecyl) phosphonium asparaginate [P66614][Asn], glutaminate [P66614][Gln], lysinate [P66614] [Lys], methioninate [P66614][Met], prolinate [P66614][Pro], taurinate [P66614][Tau], and threoninate [P66614][Thr] were synthesized and investigated as potential absorbents for CO2 capture from postcombustion flue gas. Their physical properties, including density, viscosity, glass transition temperature, and thermal decomposition temperature were determined. Furthermore, the CO2 absorption isotherms of [P66614][Lys], [P66614][Tau], [P66614][Pro], and [P66614][Met] were measured using a volumetric method, and the results were modeled with two different Langmuir-type absorption models. The most important result of this study is that the viscosity of [P66614][Pro] only increased by a factor of 2 when fully complexed with 1 bar of CO2 at room temperature. This is in stark contrast to the other chemically reacted ILs investigated here and all other amino acid-based ILs reported in the literature, which dramatically increase in viscosity, typically by 2 orders of magnitude, when complexed with CO2. The unique behavior of [P66614][Pro] is likely due to its ring structure, which limits the number and availability of hydrogen atoms that can participate in a hydrogen bonding network. We found that water can be used to further reduce the viscosity of the CO2-complexed IL, while only slightly decreasing the CO2 capacity. Finally, from temperaturedependent isotherms, we estimate a heat of absorption of 63 kJ/mol of CO2 for the 1:1 reaction of CO2 with [P66614][Pro], when we use the two-reaction model.

’ INTRODUCTION There are currently about 600 coal-fired power plants in operation in the U.S.1 In response to concerns about global climate change and anticipated future regulations on carbon dioxide (CO2) emissions, there has been increasing attention given to developing less energy intensive alternatives to the conventional technology, which would involve the use of aqueous amine solutions. Ionic liquids (ILs) are an attractive option because they have negligible vapor pressure. This is ideal because volatile compounds cause environmental concerns and raise the cost of operation through solvent replacement.2 ILs are also liquid over a wide range of temperatures, can have decomposition temperatures as high as 300400 °C, and possess virtually limitless chemical tunability.3 Since the first measurements of CO2 solubility in an IL,4 there have been many studies of CO2 + IL systems, with the focus on understanding and increasing the physical solubility of the CO2 in the IL.514 In general, increasing the length of alkyl chains tends to increase CO2 solubility,1316 as does increasing the number of fluorinated alkyl groups in either the cation or the anion.9 Nonetheless, even with these improvements, it is clear that the CO2 capacity of ILs from physical dissolution at the small partial pressures of CO2 in postcombustion flue gas (typically 98%, Sigma Aldrich) was added to the filtered solution and stirred overnight. Cold acetonitrile (ACS grade, Fisher Scientific) was used to crystallize any remaining neutral amino acid, which was removed by vacuum filtration. Solvent and water removal was achieved by pulling a moderate vacuum at 60 °C for approximately one week. 9141

dx.doi.org/10.1021/jp2015534 |J. Phys. Chem. B 2011, 115, 9140–9150

The Journal of Physical Chemistry B [P66614][Pro] was selected for additional studies, and additional quantities were generously supplied by Merck KGaA. The Merck batch of [P66614][Pro] and all ILs synthesized in our laboratory were checked for purity with 1H NMR recorded on Varian 300 spectrometers; these results are shown in Supporting Information. No impurities were identified from the NMR results. Bromide impurity content was determined to be no greater than 6000 ppm using a ColeParmer bromide selective ion conductivity probe, model EW-2750402. All ILs are hygroscopic so water content was measured by a Brinkman 831 Karl Fischer Coulometer. Water content prior to measurements (or intentional addition of water) was below 800 ppm. On the basis of these measurements, we estimate the purity at the time of the measurements to be >98%. Density Measurements. Densities were measured at atmospheric pressure in a DMA 4500 Anton Paar oscillating U-tube densitometer, which includes an automatic correction for the viscosity of the sample, as described previously.25,28 The uncertainty was estimated to be (5 105 g 3 cm3. Two integrated Pt 100 platinum thermometers provided good precision ((0.01 K) in temperature control internally. Glass Transition Temperature Measurements. The glass transition temperature was measured with a Mettler-Toledo differential scanning calorimeter (DSC), model DSC822e, and the data was evaluated using the Mettler-Toledo STARe software version 7.01, as described previously.25 Samples were dried in situ in the DSC. The presence of volatiles significantly affects the glass transition temperature; therefore, samples were dried repeatedly until the phase transition temperatures remained constant. The equipment has an accuracy of (0.3 K when tested with a standard solution, but the presence of trace impurities raises the estimated uncertainty to (1 K. Decomposition Temperature Measurements. Decomposition temperatures were measured with a Mettler-Toldeo TGA/ SDTA 851e/SF/1100 °C thermal gravimetric analyzer, as described previously.25 We report the onset temperature, which is the intersection of the baseline weight, either from the beginning of the experiment or after the drying step, and the tangent of the weight versus temperature curve as decomposition occurs. The samples were run in aluminum pans under a nitrogen atmosphere at a heating rate of 0.17 K 3 s1. When we observed any weight loss from the evaporation of water from the sample, it was further dried in situ at 403 K for 30 min. Reproducibility was verified by running three replicates for each IL. The largest uncertainty is from manually determining the tangent point, which results in an uncertainty in the thermal decomposition temperatures of (2 K. CO2 Absorption Measurements. The CO2 absorption was measured using a calibrated stirred vessel and carefully measuring the CO2 introduced into that vessel from a calibrated chamber. The calibrated chamber (177 mL for setups A and B, 290 mL for setup C) was filled with CO2 (99.99% purity, Praxair) to a pressure of about two bars. The stirred vessel (129 mL for setup A, 257 mL for setup B, 235 mL for setup C) was loaded with a known amount of IL and then evacuated to a pressure of about 5 mbar. The stirred vessel was held at a constant pressure (tolerance 1 mbar) for a minimum of 1 day to ensure that there were no substantial leaks. Then the valve that connected the stirred vessel to the calibrated chamber was briefly opened. The amount of CO2 to enter the stirred vessel was calculated from the pressure drop in the calibrated chamber, and verified by the pressure increase in the stirred vessel. Stirring was activated to mix the IL and the gas (a magnetic stir bar for setups A and B, a MagneDrive

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III Autoclave Engineers stirrer for setup C) and the pressure was recorded at regular time intervals (4 to 8 h for setups A and B, 1 to 2 h for setup C). When the pressure remained constant for a minimum of two time intervals, the system was considered to be at equilibrium. The amount of CO2 absorbed into the IL was calculated from the pressure drop in the stirred vessel between the time when CO2 entered the vessel and when equilibrium was reached. For experiments with added water, the pressure drop used to calculate CO2 uptake took into consideration the initial pressure of the water vapor present before CO2 was added. The room temperature measurements (setup A) were at 22 ( 1 °C. Temperatures below room temperature were reached by placing the stirred chamber in setup B in a water bath controlled by a 1016S Isotemp to (1 °C. Temperatures above room temperature were reached by placing the calibrated chamber and stirred cell in setup C inside a Yamato DKN602 oven, which controlled the temperature to (0.5 °C. Viscosity Measurements. The viscosity of the ILs was measured with a Brookfield DV-III Ultra (cone and plate) rheometer, as described previously.25,28 Although the nominal uncertainty of the instrument is (2%, the actual uncertainties are higher due to slight changes in the concentration of water in the samples, as described in the next paragraph. Repeat experiments showed a precision of (5%. Temperature was controlled with a precision of (0.1 K by means of a TC-602 bath thermostat with a Brookfield temperature controller unit attached. The measurement chamber for the rheometer had either nitrogen or CO2 flowing over the sample depending on if the sample was neat or saturated with CO2. The sample was exposed to the atmosphere briefly during loading and unloading, which introduces an opportunity for the sample to absorb moisture from the atmosphere. Karl Fischer measurements on samples before and after the viscosity measurements typically showed a rise in water content for the neat ILs from about 200 ppm to about 1000 ppm. The CO2 saturated samples were too viscous to measure the water content using the Karl Fischer Coulometer, except for [P66614][Pro], which will be discussed later. However, the high viscosity of the samples would impede diffusion, so it was assumed that the CO2 saturated samples absorbed less water than their neat counterparts during transfer of the samples into the viscometer. Henry’s Constant Measurements. In order to determine reaction equilibrium constants from the CO2 uptake isotherms, it is preferable to have independent measurements of the physical solubility of the CO2 in the IL. For the most thoroughly investigated IL, the [P66614][Pro], we obtained estimates of the Henry’s Law constants for the physical CO2 solubility using infrared spectroscopy. Samples of [P66614][Pro] were dried overnight so the measured water content was below 500 ppm for the all of the experiments performed in the React-IR system. Details of the system have been published previously.9 The system containing the IL sample was thoroughly evacuated prior to measurements to remove any dissolved gases. After charging the vessel with CO2 (99.99% purity, Praxair) and allowing equilibrium to be achieved, the peak area of the physically dissolved CO2 signature band in the 23702310 cm1 region was recorded. The values of peak area obtained at equilibrium were used to determine the concentration of the dissolved CO2 in the IL samples using a calibration curve developed for CO2 in 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide and 1-hexyl-3-methylpyridinium bis(trifluoromethylsulfonyl)imide. These ILs do not react with CO2 so the CO2 solubility at various temperatures and pressures is known independently from gravimetric measurements. 9142


The Journal of Physical Chemistry B

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The primary assumption in estimating the Henry’s Law constants in the chemically reacting ILs is that the absorption crosssection is the same as in the ILs that just physically dissolve CO2. The maximum equilibrium pressures were 1.3, 1.7, and 4.1 bar at 10, 25, and 55 °C, respectively. The applied CO2 pressure was not increased any further at 10 and 22 °C due to the high viscosity of the IL at these temperatures, which increases the equilibration time. Accordingly, 2, 3, and 4 equilibrium pressures were used in the estimation of the Henry’s Law constants at 10, 22, and 55 °C respectively.

’ RESULTS AND DISCUSSION Density. Density measurements of pure ILs were carried out at temperatures ranging from 10 to 70 °C. The experimental density data obtained is reported in Figure 2 for all of the pure ILs studied. The density was observed to decrease in the order [P66614][Tau] ∼ [P66614][Gln] ∼ [P66614][Asn] > [P66614][Met] > [P66614][Pro] > [P66614][Thr] > [P66614][Lys]. Gardas et al. reported a similar trend using several of the same anions with both the tetrabutylphosphonium cation ([P4444][Tau] > [P4444][Pro] > [P4444][Thr] > [P4444][Lys]) and the tributylmethylammonium cation ([N4441][Tau] > [N4441][Thr] > [N4441][Lys]).30 The densities of the [P66614][Tau], [P66614][Pro], [P66614][Thr] and [P66614][Lys] were on average 7.3 ( 0.8% less than their [P4444]+ counterparts, which supports prior claims that the density of amino acid based ILs decreases with an increase in alkyl chain length.21,25,30,31 Glass Transition and Thermal Decomposition Temperatures. Table 1 shows the glass transition temperatures of the [P66614]+ amino acid ILs, which varied from 78 to 51 °C.

Figure 2. Density of [P66614][Tau] (stars), [P66614][Gln] (dashes), [P66614][Asn] (open diamonds), [P66614][Met] (filled triangles), [P66614][Pro] (Merck batch, open triangles), [P66614][Thr] (pluses) and [P66614][Lys] (filled circles) with temperature.

These are lower than the values for the same anions with the [P4444]+,3 3-(aminopropyl)tributylphosphonium [aP4443]+,21 and 1-ethyl-3-methylimidazolium [emim]+32 cations, likely due to the bulkier alkyl chains on the cation. [Asn] and [Gln] have higher glass transition temperatures than [Lys], [Pro], [Thr], and [Met] for all four cations, possibly due to the amide functional group. [Met] had the lowest glass transition temperature out of this group of amino acid ILs for all four cations. No melting points, freezing points, or cold crystallization temperatures were observed for any of these ILs except [P66614][Pro], which had a melting temperature of 21 °C. The decomposition temperatures were observed to decrease in the order [Tau] > [Gln] > [Pro] > [Met] > [Lys] > [Thr] > [Asn] (Table 1). These values are higher than those for the smaller [P4444]+ cation.3,20 The [P66614]+ cation is a very stable cation with a decomposition temperatures as high as 420 °C achieved when paired with a more stable anion.33 We suspect asymmetry of the [P66614]+ cation increases its stability. Although multiple groups have demonstrated that the introduction of asymmetry to the phosphonium cation affects the viscosity and glass transition temperature,31,33 to our knowledge the effect on decomposition temperature has not yet been reported and should be the subject of further research. CO2 Absorption. [P66614][Lys], [P66614][Pro] (Merck batch), and [P66614][Tau] were selected to experimentally measure the CO2 uptake at room temperature. The absorption experiment for [P66614][Lys] had 2 wt % water added because the neat IL became too viscous to stir upon the addition of CO2. This phenomenon will be addressed in more detail later. The isotherms are plotted in Figure 3 with previously reported data for [P66614][Met],24

Figure 3. CO2 capacity in CO2:IL molar ratio at room temperature (22 °C) for [P66614][Lys] (filled diamonds), [P66614][Gly] (filled circles),25 [P66614][Ile] (filled squares),25 [P66614][Pro] (Merck batch, open triangles), [P66614][Sar] (open circles),25 [P66614][Met] (stars),24 [P66614][Tau] (dashes) and [P66614][Ala] (pluses).25

Table 1. Glass-Transition and Decomposition Temperatures compound

molecular weight

glass transition temperature

decomposition temperature

[P66614][Asn]

614.98

51 °C

244 °C

[P66614][Gln]

629.01

61 °C

345 °C

[P66614][Lys]

629.05

68 °C

296 °C

[P66614][Met] [P66614][Pro] (Merck batch)

632.08 597.99

78 °C 72 °C

300 °C 327 °C

[P66614][Tau]

608.01

75 °C

374 °C

[P66614][Thr]

601.98

74 °C

264 °C

9143



Figure 4. CO2 capacity in CO2:IL molar ratio of [P66614][Pro] at room temperature (22 °C) for the batch synthesized in house (circles)24 and the Merck batch (triangles).

[P66614][Gly],25 [P66614][Ile],25 [P66614][Sar],25 and [P66614][Ala].25 There are three distinct groupings of data in Figure 3: [P66614][Ala] and [P66614][Tau] reach a saturation level just above 0.5 mol of CO2 per mole of IL, [P66614][Met], [P66614][Sar], [P66614][Ile], and [P66614][Pro] reach a saturation level just under 1 mol of CO2 per mole of IL, and [P66614][Gly] and [P66614][Lys] reach a saturation level above 1 mol of CO2 per mole of IL. All of the isotherms rise above 0.5 mol of CO2 per mole of IL, which would be the capacity if the 1:2 reaction mechanism was dominant. [P66614][Lys] rises above 1 mol of CO2 per mole of IL, but there are two amines on [Lys], so its theoretical limit for the 1:1 absorption mechanism is 2 mol of CO2 per mole of IL. [P66614][Gly] is the only IL to exceed the theoretical limit for the 1:1 absorption mechanism. Previous experiments showed the formation of a precipitate and indicate that an irreversible side reaction occurs.34 It is apparent that additional chemistry is occurring with [P66614][Gly], which has not yet been fully investigated. The deviation between the theoretical limit of the 1:1 absorption mechanism (eq 3) and the measured isotherms is consistent between batches, as shown in Figure 4, which compares the isotherms of two separate batches of [P66614][Pro], one synthesized in our laboratory and one provided by Merck KGaA. There is an increased uncertainty at higher pressures where there is increased capacity due to physical absorption, but at low pressures, where chemical absorption is dominant, our data shows excellent consistency between the two batches. Nonetheless, it is clear that the [P66614][Pro] does not fully chemically complex 1 mol of CO2 per mole of IL. Moreover, the CO2 uptake by [P66614][Tau] is well below 1:1 and the CO2 uptake by [P66614][Lys] is well below the theoretical limit of 2:1 (recall that there are two amines in each lysinate anion). The NMR spectra do not show any significant impurities, and the residual halide content (0.5 bar), where the contribution from the physically dissolved CO2 becomes somewhat more important. Moreover, when we fit the CO2 uptake isotherms, the values of the Henry’s Law constant, H, can change by an order of magnitude with no substantial impact on the other parameters or the visual quality of the fit. This is because there simply is not enough physically dissolved CO2 at the relatively low pressures investigated to get a good fit of the Henry’s Law constant, H. From our previous experience we would anticipate that the Henry’s law constants for physical dissolution of CO2 in ILs at room temperature to be in the range of 30 to 200 bar.511 Moreover, our estimates of the Henry’s law constant for CO2 in [P66614][Pro] at room temperature from infrared spectroscopy is 9144



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Table 2. Fit Parameters for the Two-Reaction Model with Henry’s Law Constants Fixeda H (bar)

k1 (bar1)

[P66614][Lys]

100

24

15

[P66614][Tau] [P66614][Pro] Merck batch

100 57

29 550

4.3 1.9 2.5

compound

k2 (unitless)

[P66614][Pro]b synthesized in house

57

560

[P66614][Met]b

147

540

19

[P66614][Ile]b

100

130

0.5

[P66614][Sar]b

100

9700

100

a

Calculated parameters for the CO2 absorption model in eqs 57. b Parameter values obtained from the literature.25 Figure 7. Calculated concentrations of total [CO2] (solid line), [RNH2CO2]2/+ (dash-dot line), [RNHCO2]2 (dashed line), [CO2](phys) (dotted line), and experimental total [CO2] (open squares) per mole of IL for [P66614][Tau].

Figure 6. Calculated concentrations of total [CO2] (solid line), [RNH2CO2]2/+ (dash-dot line), [RNHCO2]2 (dashed line), [CO2](phys) (dotted line), and experimental total [CO2] (open squares) per mole of IL for [P66614][Pro] (Merck batch). 24

57 bar. Therefore, we refit the experimental CO2 uptake data for [P66614][Lys], [P66614][Pro] (Merck batch), and [P66614][Tau] with the two-reaction model, using the Henry’s law constant at room temperature from the literature24 for [P66614][Pro] and set H equal to 100 bar for [P66614][Lys] and [P66614][Tau] because it is in the middle of the range of expected values for physical dissolution of CO2 in ILs. The results of the refits (with H values set) are shown in Table 2. Despite changing the Henry’s law constant by an order of magnitude in some cases, there was little difference in the quality of the fit between using all three parameters freely fit and only fitting the reaction equilibrium constants with a fixed H value, as shown by the dotted lines in Figure 5. This indicates that the physical absorption has a negligible impact on the absorption isotherms at pressures below 1 bar. Table 2 lists the calculated parameters for the two-reaction model with Henry’s constant fixed for the data presented in this paper and for previously reported data.25 Good agreement is observed between the two batches of [P66614][Pro] with k1 values of 550 and 560 bar1 and k2 values of 1.9 and 2.5. Among the fitted isotherms, [P66614][Sar] stands out for having a k1 value an order of magnitude greater than the other isotherms, and [P66614][Tau] and [P66614][Lys] stand out for having k1 values an order of magnitude smaller than the other isotherms. As reported previously,25 we believe the k1 value of [P66614][Sar] is an artifact of there being only one data point in the lowest pressure portion of the isotherm and a lack of data points in the curved region

Figure 8. Calculated concentrations of total [CO2] (solid line), [RNH2CO2]2/+ (dash-dot line), [RNHCO2]2 (dashed line), [CO2](phys) (dotted line), and experimental total [CO2] (open squares) per mole of IL for [P66614][Lys].

before the isotherm flattens out, rather than any intrinsic preference of the [Sar] for CO2. [P66614][Tau] and [P66614][Lys] not only have a low k1 value, but absorb significantly less than 1 mol of CO2 per mole of amine, much lower than the other ILs, indicating a stronger effect of the 1:2 mechanism (remember that [P66614][Lys] has two amines on each anion). Perhaps the best way to examine the relative contributions of eq 3 (the 1:1 reaction) and eq 4 (the 1:2 reaction) is to plot the concentration of [RNH2CO2]2/+ (the product of the 1:1 reaction) and [RNHCO2]2 (the product of the 1:2 reaction). Figures 6, 7, and 8 show the relative amounts of physically dissolved CO2, [RNH2CO2]2/+, and [RNHCO2]2 in [P66614][Pro] (Merck batch), [P66614][Tau], and [P66614][Lys] using the parameters in Table 2. As is clear in Figure 6, in [P66614][Pro] (Merck batch) the contribution of the 1:2 mechanism and the subsequent concentration of the dianion species, [RNHCO2]2, is very small. On the other hand, the contribution of the 1:2 mechanism for [P66614][Tau], and [P66614][Lys] is substantial, as evident from the high concentration of [RNHCO2]2 at all pressures. At 1.5 bar of CO2, [P66614][Pro] (Merck batch), [P66614][Tau], and [P66614][Lys] have 0.04, 0.19, and 0.56 mol of [RNHCO2]2 per mole of IL. A reasonable explanation for varying contributions of the 1:2 9145



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Table 3. Fit Parameters for the Deactivated IL Model with Henry’s Law Constant Fixed, except for [P66614][Lys] and [P66614][Tau]a H (bar)

k1 (bar1)

C3 (unitless)

[P66614][Lys]

6.1

390

1.18

[P66614][Tau]

7.6

160

0.64

[P66614][Pro] Merck batch

57

220

0.92

[P66614][Pro]b synthesized in house

57

380

0.88

[P66614][Met]b [P66614][Ile]b

147 100

93 160

0.86 0.91

[P66614][Sar]b

100

2900

0.86

compound

a

Calculated parameters for the CO2 absorption model in eqs 8. b Parameter values obtained from the literature.25 Figure 9. CO2 capacity in CO2:IL molar ratio at room temperature (22 °C) for [P66614][Lys] (squares), [P66614][Pro] (Merck batch, triangles), and [P66614][Tau] (diamonds). Each isotherm is fit with the deactivated IL model with all variables free (dashed lines) and with the Henry’s law constant fixed to the values in Table 3 (dotted lines).

mechanism is the relative stability of the dianion in the various ILs. In [Pro] and most of the ILs tested previously, there are two carbons between the amine and the charge carrying functional group; in [Tau] there are three carbons, and in [Lys] there are two carbons between the charge carrying functional group and the first amine, but the second amine is six carbons away. It is possible that this increased distance allows for more stabilization of the dianion (formed by the 1:2 mechanism, eq 4) because the 2 charge is distributed over a larger area. The two-reaction model seems to be a reasonable approach and provides an excellent fit to the experimental CO2 uptake data. Unfortunately, infrared spectroscopy studies of [P66614][Pro] and [P66614][Met] do not support the presence of significant amounts of [RNHCO2]2.24 Therefore, we have developed an alternative model, which assumes that only the 1:1 reaction (eq 3) takes place and that only a fraction of the amine groups are available for complexation. Allowing less than 100% of the ILs to have the capability of reacting with CO2 provides a mechanism to explain the isotherms not reaching a capacity of 1.0 CO2:IL molar ratio (2.0 for [Lys]). Equation 8 is this alternative “deactivated IL” model, derived assuming that the reaction in eq 3 is the only chemical reaction occurring, but that there is an amount of deactivated IL that is independent of the CO2 concentration. C3 is the molar ratio of active amines to total IL and the other variables are the same as the previous model. The derivation of the deactivated IL model is shown in the Supporting Information. z¼

P=H k1 PC3 + 1 P=H 1 + k1 P

ð8Þ

Figure 9 compares the quality of this model with all three variables (H, k1, and C3) free versus that with the Henry’s Law constant fixed at 57 bar for [P66614][Pro], and 100 bar for [P66614][Lys] and [P66614][Tau]. There is no significant difference in the quality of the fit for [P66614][Pro] (Merck batch) using either the value of the Henry’s Law constant from IR spectroscopy (57 bar) or the freely fit value (18 bar). So for further discussion we will refer to the fit of the [P66614][Pro] data using H = 57 bar. However, using the reasonable estimates of Henry’s law constants of 100 bar for [P66614][Lys] and [P66614][Tau] (the dotted lines in Figure 9) does not provide good fits of the experimental data using the deactivated IL model. When all three parameters (H, k1, and C3)

are freely fit to the data, the representation (the dashed lines in Figure 9) is excellent. Unfortunately, the Henry’s Law constants needed to produce this result, 7.6 and 6.1 bar, are not at all reasonable. This would yield exceptionally high physical CO2 solubility, which is not supported by any other experimental observations in our laboratory. The single 1:1 reaction deactivated IL model forces essentially all the contribution of the chemical reaction to very low pressures so that any increase in capacity beyond the sharp “corner” has to be taken into account by the physical dissolution of the CO2. Nonetheless, the parameters for the deactivated IL model that provide a good representation of the [P66614][Pro] (Merck batch), [P66614][Lys], and [P66614][Tau] data, along with fits for previously reported systems, are reported in Table 3.25 Some information can be gleaned from the parameter values in Table 3. Using a fixed value of the Henry’s Law constant of 57 bar for CO2 in [P66614][Pro] forces the values of k1 and C3 to be somewhat different for the isotherms for the two different samples. This is because the slope of the data at the higher pressures is a bit larger for the Merck batch than the batch synthesized in house and a common value of H forces k1 and C3 to compensate. [P66614][Lys] has a C3 value of 1.18, which corresponds to 59% of the amine sites available for complexation with CO2. It is likely that the assumption that the two amines behave independently is not correct. The reaction of the first amine with CO2 could change the electronics such that reaction of the second amine on the anion with CO2 is less favorable. And, of course, this is the sample that has 2 wt % water added to reduce the viscosity, which could further complicate the chemistry. [P66614][Tau] has a low C3 value of 0.64, while the rest of the ILs have C3 values in the range of 0.86 to 0.92. The easiest explanation for C3 values less than 1.0 (or less than 2.0 for [P66614][Lys]) would be the presence of impurities in the samples. Although halide impurities of up to 0.6 wt % are possible, this cannot explain C3 values as low as 0.64, or even 0.86 to 0.92. Moreover, the NMR spectra do not give any indication of impurities. In summary, it appears that the two-reaction model provides a better fit of the experimental uptake data for [P66614][Lys] and [P66614][Tau] with more reasonable parameters. Both models perform equally well for [P66614][Pro]. It is possible that some impurities in the IL samples may be reducing the ultimate capacity for CO2, which is better taken into account by the deactived IL model. Thus, we believe that the two models presented here cover the range of possibilities of the chemistries that could be occurring. Viscosity. Viscosity measurements of the pure ILs were carried out at temperatures ranging from 10 to 70 °C under 1 bar of nitrogen. The experimental data obtained is reported in Figure 10. 9146



Figure 10. Viscosity of [P66614][Asn] (filled circles), [P66614][Gln] (open circles), [P66614][Pro] (Merck batch, open triangles), [P66614][Lys] (+’s), [P66614][Tau] (’s), [P66614][Thr] (dashes), and [P66614][Met] (filled triangles) as a function of temperature.

Figure 11. Viscosity of CO2 saturated samples of [P66614][Lys] (triangles), [P66614][Met] (circles), [P66614][Tau] (squares), and [P66614][Pro] (Merck batch, diamonds).

The viscosity of the ILs decreased in the order [Asn] > [Gln] > [Pro] ≈ [Lys] ≈ [Tau] > [Thr] > Met]. The room temperature viscosity of [P66614][Tau], [P66614][Thr], and [P66614][Pro] were about half the viscosity of their [P4444]+ counterparts,30 but this difference decreased with increasing temperature. The exception to this trend is [Lys], which had a viscosity of 730 cP with the [P66614]+ cation and 740 cP with the [P4444]+ cation.30 The viscosity of samples of [P66614][Lys], [P66614][Met], [P66614][Tau], and [P66614][Pro] that have been saturated with CO2 at room temperature and 1 bar are shown in Figure 11. The room temperature viscosity of [P66614][Lys] increased 217-fold upon complexation with CO2, [P66614][Met] increased 94-fold, and [P66614][Tau] increased 29-fold. The CO2 saturated sample of [P66614][Lys] had 2 wt % water, so the true viscosity increase is likely even higher than this value. The increase in viscosity of [P66614][Tau], [P66614][Met], and [P66614][Lys] correlates with CO2 capacity. These observations of dramatic increases in viscosity of amino acid IL when they are reacted with CO2 are consistent with all other reports in the literature.20,22,25 This large increase in viscosity would make any practical application of ILs for postcombustion CO2 capture quite challenging. Amazingly, the viscosity of [P66614][Pro] at room temperature barely doubled when fully complexed with CO2 at 1 bar pressure. This is not

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Figure 12. Effect of water content on the viscosity of the uncomplexed [P66614][Pro] (mixed batches, triangles) and [P66614][Tau] (diamonds) and of the CO2 complexed [P66614][Pro] (circles) at 25 °C.

because of low CO2 capacity; the capacity of [P66614][Pro] is nearly 1 mol of CO2 per mole of IL at room temperature and 1 bar CO2 pressure (see Figure 4). There is evidence that the large increase in viscosity of amino acid based ILs is the result of the CO2-complexed ILs forming a salt bridge hydrogen bonded network.26 Thus, freely available hydrogen atoms on the anion would be expected to promote hydrogen bonding and increases in viscosity, as is observed here for [P66614][Tau], [P66614][Met], and [P66614][Lys]. [P66614][Pro] is unique out of the ILs studied in that it has a ring structure so that the nitrogen participating in complexation with CO2 is actually a secondary amine. This ring structure translates into less hydrogen atoms available to form a hydrogen bonding network, and it may provide some steric interference as well, hence the very small increase in viscosity. Although other researchers have synthesized prolinate ILs, they were not tested for CO2 uptake.20 Thus, to our knowledge, this is the first report of an amino acidbased IL with high CO2 uptake and an extremely small increase in viscosity. The discovery that [P66614][Pro] does not increase much in viscosity when it is complexed with CO2 and insight from molecular dynamics on hydrogen bonding networks formed in amino acid-based ILs26 has led us to develop a new class of ILs with aprotic heterocylic anions, which have been reported elsewhere.35 Effect of Water. While the discovery of the very small viscosity increase of [P66614][Pro] upon complexation with CO2 is a major advance, we recognize that further reductions in viscosity may be possible with the addition of water. As mentioned above, water would be present in postcombustion flue gas and some amount of water would be absorbed by the IL depending on its hydrophilicity. The effect of water on the viscosity of the uncomplexed [P66614][Pro] and [P66614][Tau] and CO2-complexed [P66614][Pro] at 25 °C is shown in Figure 12 and listed in the Supporting Information. Both neat ILs have a viscosity at 25 °C of about 700 cP. That viscosity drops by about 100 cP with just 0.1 wt % water, and drops a total of 68% at room temperature with 7.1 wt % water. Water has a smaller effect on the viscosity of the CO2 complexed IL at low water contents, but at about 5 wt % water there appears to be a sudden drop off in viscosity, presumably as the water begins to interfere with the IL/CO2 hydrogen bonded network. After this, the viscosity of the CO2-complexed [P66614][Pro] is in the range of 300500 cP at 25 °C and 78 wt % water, which is a 58 9147



Figure 13. CO2 capacity in CO2:IL molar ratio at room temperature (22 °C) for dry [P66614][Pro] (Merck batch, open triangles), [P66614][Pro] with 4 wt % water (Merck batch, filled triangles), dry [P66614][Met] (open diamonds),9 and [P66614][Met] with 14 wt % water (filled diamonds).

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Figure 15. Plot of ln H (squares), ln k1 from the two-reaction model (+’s), ln k1 from the deactivated IL model (’s), and ln k2 (triangles) versus the inverse temperature for [P66614][Pro] (mixed batches).

Table 4. Fit Parameters for the Two-Reaction Model and the Deactivated IL Model with Henry’s Law Constants Fixed for [P66614][Pro] at Different Temperatures two-reaction model

deactivated IL model

temperature H (bar) k1 (bar1) k2 (unitless) k1 (bar1) C3 (unitless) 10 °C

35

10 000

46

1100

0.90

22 °C

57

550

1.9

220

0.92

22 °Ca

57

560

2.5

380

0.88

40 °C (1)

81

46

0.077

63

0.89

40 °C (2)

81

34

0.0098

39

0.95

60 °C

128

38

0.0098

42

0.96

80 °C 100 °C

192 277

17 5.4

0.019 0.0098

23 6.5

0.92 0.94

Figure 14. CO2 capacity in CO2:IL molar for [P66614][Pro] (Merck batch) at 10 °C (diamonds), 60 °C (circles), 80 °C (triangles), and 100 °C (squares).

a

to 75% decrease in viscosity from the dry CO2 complexed sample. Each data point in Figure 12 contains error bars, which are visible with many of the symbols. These correspond to the water content, as measured by Karl Fischer coulometry, at the beginning and end of the viscosity determination. Relatively dry samples will pick up some water, while samples with higher water content can actually lose a bit of water during the measurements. Even with these uncertainties, the trend of decreasing viscosity with increasing water content is clear. Moreover, it is true for both pure IL samples and CO2-complexed IL samples. These are promising results because it shows that water, which would be present in any postcombustion CO2 capture process, will help to reduce the viscosity of the IL in an industrial setting. A major question is whether the presence of water will affect the CO2 capacity of the chemically complexing ILs. Figure 13 shows the effect of 4 wt % water on the CO2 absorption of [P66614][Pro] (Merck batch) and 14 wt % on the absorption of [P66614][Met]. With 4 wt % water, capacity of [P66614][Pro] appears to be slightly decreased at the lower pressures, but at higher pressures the two curves become statistically equivalent. With 14 wt % water, the CO2 capacity of [P66614][Met] is reduced by approximately 0.2 mol of CO2 per mole of IL at 0.25 bar and by 0.1 mol of CO2 per mole of IL at 1 bar. Despite this reduction, the absorption isotherm for the [P66614][Met]water sample still rises significantly above 0.5 mol of CO2 per mole of IL. From these limited

experiments, we conclude that water seems to have a substantial impact on the CO2 absorption capabilities of ILs only when large quantities are present. Effect of Temperature. The absorption isotherm for [P66614][Pro] were measured at 10, 40, 60, 80, and 100 °C as shown in Figure 14 and listed in the Supporting Information. Repeat experiments with different samples of [P66614][Pro] (Merck batch) at 40 °C show similar uncertainty to those observed at 22 °C. The results in Figure 14 are all with the Merck sample. The 22 and 40 °C data have been omitted from Figure 14 for clarity, but are listed in the Supporting Information. Clearly, as the temperature increases, the isotherms increase less rapidly in the low pressure region, and the total capacity at 1 bar is decreased slightly. It is possible to determine the enthalpy and entropy of reaction from the temperature dependence of the reaction equilibrium constants. Therefore, we fit the isotherms to the two models described earlier: the two-reaction model and the deactivated IL model. As mentioned above, it is preferable to have experimental measurements of the Henry’s Law constant (H) instead of relying on the fit of the isotherms to estimate H since it is very sensitive to the slope of the isotherm at higher pressures. H was experimentally measured by IR spectroscopy to be 35 bar at 10 °C and 113 bar at 55 °C. The physical heat of absorption was then calculated to be 20 kJ/mol using these two measurements and the room temperature value of 57 bar,24 by finding the slope 9148

Synthesized in house. All other runs are from the Merck batch.



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Table 5. Summary of CO2 Absorption Results for [P66614] Amino Acid Based ILs at 25 °C and 1 bar and the Impact on Viscosity, ηa two-reaction model anion

total CO2

[CO2](phys)

2/+

[RNH2CO2]

deactivated IL model 2

[RNHCO2]

[CO2](phys)

[RNH2CO2]2/+

0.01

1.21

48 240 217e

η rise

[Ala] b

0.66

[Gly] b

1.26

[Ile] [Lys]

0.97 1.37

0.01 0.01

0.88 0.76

0.06 0.61

0.01 0.19

0.91 1.17

[Met] b

0.88

0.01

0.74

0.14

0.01

0.86

94

[Pro] c

0.96

0.02

0.89

0.05

0.02

0.92

1.7

[Pro] b,d

0.88

0.02

0.88

0.06

0.02

0.89

[Sar] b

0.91

0.01

0.83

0.09

0.01

0.86

189

[Tau]

0.80

0.01

0.56

0.21

0.15

0.64

29

117

a

Concentrations reported in moles per total moles of IL. Total CO2 calculated by linear interpolation of the nearest experimental data points above and below 1 bar. Models used the parameters reported in Tables 2 and 3. η rise is reported in percent increase 100. b Values from the literature.25 c Merck batch. d Batch synthesized in house. e CO2 saturated sample used to calculate the increase had 2 wt % water.

of ln H versus 1/T (Kelvin) (Figure 15). This value of enthalpy is a bit higher than other values for the partial molar enthalpy for CO2 physical dissolution in ILs but still reasonable.8 This was then used to estimate the values of H at 40, 60, 80, and 100 °C (Table 4). Fixing H at these values, the different temperature isotherms were fit to the two-reaction and deactivated IL models (Table 4). At all the temperatures, the values of the equilibrium constants for the 1:1 reaction, k1, are remarkably similar for the two different models. Basically, the equilibrium constant decreases with increasing temperature. As expected, ln k1 is roughly a linear function of 1/T, as shown in Figure 15. The extremely high value of k1 at 10 °C reflects the very steep isotherm in the low pressure region. In the two-reaction model, the values of k2 do not follow a reasonable trend with temperature (Figure 15). One reason is that the contribution from the 1:2 reaction is very small with [P66614][Pro]. This is easily understood by recognizing that all of the isotherms approach 1 mol CO2/mol IL at the higher pressures. This is also clear from looking at the values of C3 from the deactivated IL model, all of which are above 0.89. The reaction enthalpy and entropy for the 1:1 reaction are found to be 63 kJ/mol and 160 J/mol 3 K using the values of k1 from the two-reaction model and 44 kJ/mol and 100 J/mol 3 K using the values of k1 from the deactivated IL model. A calorimetric measurement of the total change in enthalpy resulting from CO2 uptake by [P66614][Pro] is 80 ( 5 kJ/mol.24 Since we have shown that the 1:1 reaction is the dominant mechanism for CO2 uptake by [P66614][Pro], the estimate of the enthalpy from the temperature dependence of k1 should be consistent with the measurement from calorimetry. The somewhat better agreement between the calorimetric measurements and the enthalpy we obtain from the two-reaction model lends some support to that approach for the reaction of CO2 with [P66614][Pro]. Clearly, the greatest uncertainty in the enthalpies determined from the temperature-dependent CO2 isotherms is the choice of the reaction model.

’ CONCLUSIONS A series of seven amine-functionalized anion-tethered ILs with tetra-alkylphosphonium cations were synthesized. Their physical properties, including density, viscosity, glass transition temperature,

and thermal decomposition temperature were determined. These ILs have low glass transition temperatures in the range of 51 to 78 °C and thermal decomposition temperatures in the range of 244374 °C. The CO2 absorption isotherms of [P66614][Lys], [P66614][Pro], [P66614][Tau], and [P66614][Met] were measured at room temperature using a volumetric method. All four ILs achieved capacities greater than theoretically possible if a 1:2 reaction mechanism was dominant. The uptake isotherms were fit with two models. The two-reaction model included both 1:1 and 1:2 reaction mechanisms. The deactivated IL model assumed only a 1:1 reaction but included a parameter for the number of active sites per IL in order to account for capacities lower than one CO2 per mole of IL. Both models fit the experimental data very well and represent the range of possible chemistries that may occur between the amino acid ILs and CO2. The chemical absorption of CO2 dramatically increased the viscosity of all ILs except [P66614][Pro], likely due to its ring structure limiting hydrogen atoms available to form a hydrogen bonding network. The presence of water significantly reduced the viscosity of both the pure and CO2 saturated ILs, and caused only a slight decrease in the CO2 capacity. Multiple isotherms for absorption of CO2 into [P66614][Pro] between 10 and 100 °C were used to estimate equilibrium constants as a function of temperature and the heat of reaction for the 1:1 reaction. We estimate an enthalpy of the 1:1 reaction between CO2 with [P66614][Pro] of roughly 63 kJ/mol, using the two-reaction model. A summary of all the result presented here compared with ILs presented previously25 can be found in Table 5, where we list each IL, the total capacity at 1 bar CO2 pressure, the contributions to the CO2 uptake from chemical and physical mechanisms as calculated by the tworeaction model and the deactivated IL model, as well as the viscosity increase observed upon complexation with CO2.

’ ASSOCIATED CONTENT

bS

Supporting Information. The Supporting Information file contains details on the NMR analysis of the ILs, derivation of the models used, a table of density data for all the ILs, and all the CO2 uptake data in tabular form. This information is available free of charge via the Internet at http://pubs.acs.org/.

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’ AUTHOR INFORMATION Corresponding Author

*Tel: (574) 631-5847. Fax: (574) 631-8366. E-mail: [email protected]. Present Addresses †

Departamento de Ingeniería Química y Ambiental, Universidad Tecnica Federico Santa María, Valparíso, Chile.

’ ACKNOWLEDGMENT We thank Kathleen Stanley and Michael Glaser for their contributions in synthesis, and David Zadigian, Samantha Miller, Florence Chen, and Serena Mathews for their contributions in data collection. This material is based upon work supported by the Department of Energy under Award Number DE-FC26-07NT43091. ’ REFERENCES (1) Electric Power Annual; U.S. Energy Information Administration (EIA): Washington DC, 2010. Available online at http://www.eia.doe. gov/cneaf/electricity/epa/epa_sum.html (accessed Jul 31, 2010). (2) Aaron, D.; Tsouris, C. Separation of CO2 from Flue Gas: A Review. Sep. Sci. Technol. 2005, 40, 321. (3) Kagimoto, J.; Fukumoto, K.; Ohno, H. Chem. Commun. 2006, 2254. (4) Blanchard, L. A; Brennecke, J. F. Ind. Eng. Chem. Res. 2001, 40, 287. (5) Kamps, A. P.; Tuma, D.; Xia, J.; Maurer, G. J. Chem. Eng. Data 2003, 48, 746. (6) Husson-Borg, P.; Majer, V.; Costa Gomes, M. F. J. Chem. Eng. Data 2003, 48, 480. (7) Baltus, R. E.; Culbertson, B. H.; Dai, S.; Luo, H.; DePaoli, D. W. J. Phys. Chem. B 2004, 108, 721. (8) Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. Acc. Chem. Res. 2007, 40, 1208. (9) Muldoon, M. J.; Aki, S. N. V. K.; Anderson, J. L.; Dixon, J. K.; Brennecke, J. F. J. Phys. Chem. B 2007, 111, 9001. (10) Carlisle, T. K; Bara, J. E.; Gabriel, C. J.; Noble, R. D.; Gin, D. L. Ind. Eng. Chem. Res. 2008, 7005. (11) Jalili, A. H.; Mehdizadeh, A.; Shokouhi, M.; Sakhaeinia, H.; Taghikhani, V. J. Chem. Thermodyn. 2010, 42, 787. (12) Shiflett, M. B.; Yokozeki, A. J. Phys. Chem. B 2007, 111, 2070. (13) Shariati, A; Peters, C. J. J. Supercrit. Fluids 2004, 29, 43. (14) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2004, 30, 139. (15) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2003, 25, 109. (16) Shariati, A.; Peters, C. J. J. Supercrit. Fluids 2005, 34, 171. (17) Bates, E. D.; Mayton, R. D.; Ntai, I.; Davis, J. H., Jr. J. Am. Chem. Soc. 2002, 124, 926. (18) Hanioka, S.; Maruyama, T.; Sotani, T.; Teramoto, M.; Matsuyama, H.; Nakashima, K.; Hanaki, M.; Kubota, F.; Goto, M. J. Membr. Sci. 2008, 314, 1. (19) Myers, C.; Pennline, H.; Luebka, D.; Ilconich, J.; Dixon, J. K.; Maginn, E. J.; Brennecke, J. F. J. Membr. Sci. 2008, 322, 28. (20) Zhang, J.; Zhang, S.; Dong, K.; Zhang, Y.; Shen, Y; Lv, X. Chem. —Eur. J. 2006, 12, 4021. (21) Zhang, Y.; Zhang, S.; Lu, X.; Zhou, Q.; Fan, W.; Zhang, X. Chem.—Eur. J. 2009, 15, 3003. (22) Jiang, X. C.; Nie, Y.; Li, C. X.; Wang, Z. Fuel 2008, 87, 79. (23) Mindrup, E. M.; Schenider, W. F. ACS Symposium Series; Seddon, K., Rogers, R., Plechkova, N., Eds.; American Chemical Soceity: Washington, D.C., 2009. (24) Gurkan, B. E.; de la Fuente, J. C.; Mindrup, E. M.; Ficke, L. E.; Goodrich, B. F.; Price, E. A.; Schneider, W. F.; Brennecke, J. F. J. Am. Chem. Soc. 2010, 132, 2116.

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Effect of Water and Temperature on Absorption of CO2 by Amine

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