Experimental Measurement of the Solubility and Diffusivity of CO 2 in Room-Temperature Ionic Liquids Using a Transient Thin-Liquid-Film Method

Experimental Measurement of the Solubility and Diffusivity of CO2 inRoom-Temperature Ionic Liquids Using a Transient Thin-Liquid-Film Method

Ying Hou and Ruth E. Baltus*

Department of Chemical and Biomolecular Engineering, Clarkson UniVersity, Potsdam, New York 13699-5705

In this paper, results from an experimental investigation of carbon dioxide (CO2) solubility and diffusivity inthe ionic liquids 1-n-butyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([bmim][Tf2N]), 1,2-dimethyl-3-propylimidazolium bis(trifluoromethylsulfonyl)imide ([pmmim][Tf2N]), 1-butyl-3-methylpyridiniumbis(trifluoromethyl sulfonyl)imide ([bmpy][Tf2N]), 1-(3,4,5,6-perfluorohexyl)-3-methylimdazolium bis(tri-fluoromethyl sulfonyl)imide ([perfluoro-hmim][Tf2N]), and 1-n-butyl-3-methylimidazolium tetrafluoroborate([bmim][BF4]) are reported. A transient thin liquid film method was developed, which enables one to determinethe Henry’s law constant and the diffusivity at low pressure simultaneously. Measurements were performedat temperatures in the range of 283-323 K. Henry’s law constants were in the range of 25.5-84 bar andwere in general agreement with measurements reported by other researchers for these and similar ionic liquids.The entropies and enthalpies of absorption were determined to increase as gas solubility decreased. Themeasured CO2 diffusion coefficients in the five ionic liquids were∼10-6 cm2/s, which is an order of magnitudesmaller than the coefficients for CO2 diffusion in traditional organic solvents. In contrast to the gas solubilityresults, measured diffusion coefficients were determined to be dependent on the ionic liquid cation as well asthe anion. In addition, CO2 diffusion coefficients were considerably more sensitive to temperature than wereCO2 solubilities in these ionic liquids. Results were used to develop a correlation relating CO2 diffusion toionic liquid properties and system temperature.

1. Introduction

Room-temperature ionic liquids (RTILs) are organic salts thatmelt near room temperature. RTILs generally have bulky low-symmetry organic cations with a delocalized charge, which leadsto their low melting point.1-4 The ionic liquid anions are moresymmetrical and are generally smaller than the cations. Commonanions include hexafluorophosphate (PF6

-), tetrafluoroborate(BF4

-), and (bis(trifluoromethylsulfonyl)imide) (Tf2N-). RTILshave negligible vapor pressure;5-7 therefore, they are consideredto be “green” solvents that may potentially replace conventionalvolatile organic compounds (VOCs) in many reaction andseparation processes.8 The properties of RTILs, such as meltingpoint, viscosity, density, and gas and liquid solubility, can beadjusted via the appropriate choice of cation and anion.9,10RTILsare also nonflammable, have high thermal stability, and have awide liquid range.11,12

To design ionic-liquid-based processes, knowledge of thesolubilities and diffusivities of gases in room-temperature ionicliquids is needed. However, only limited information about theseproperties is available in the literature. Carbon dioxide (CO2)has been used in many previous studies of solubility in ionicliquids, because it is of interest for a variety of ionic liquidapplications. Therefore, it was a convenient solute choice forthis study. The objective of this work was to develop anexperimental approach to measure both solubility and diffusivityof CO2 in RTILs, as well as examine relationships betweenRTIL structure and these properties.

In this work, the solubility and diffusivity of CO2 in fivedifferent ionic liquids were determined. An experimental methodbased on the one-dimensional diffusion of solute gas into a thinionic liquid film was developed. A major advantage of this

method over other techniques is that both the solubility anddiffusivity of the target gas in a RTIL can be determined in asingle experiment that requires only a small sample of ionicliquid (∼250 µL). The effect of temperature on solubility anddiffusivity was examined by conducting measurements atdifferent temperatures, ranging from 283 K to 323 K. Becauseexperiments were conducted at low pressure (∼1-2 bar),measured diffusion coefficients can be considered to be infinitedilution values. An expression relating CO2 diffusivity to ionicliquid properties was developed. Ionic liquid density andviscosity were also measured for the ionic liquids whereliterature values were not available.

2. CO2 Uptake Model

The experimental method involves tracking the decrease inpressure that results following the introduction of CO2 into asmall closed chamber that contains a thin film of ionic liquid.The solubility and diffusivity were determined by fitting thepressure decay to a one-dimensional diffusion model for soluteuptake into the liquid. Before each experiment, vacuum wasapplied to the chamber. At timet ) 0, CO2 was introduced atpressureP0 and the system was sealed. The decay in pressureabove the liquid film resulting from CO2 absorption into the ILwas monitored as a function of time.

The system geometry is defined such thatz ) 0 correspondsto the base of the liquid film andz ) L corresponds to the gas/liquid interface. The assumptions of the model are as follows:(i) one-dimensional diffusion of CO2 in thez (vertical)-directiononly; (ii) no convective transport in the system; (iii) the ionicliquid has negligible vapor pressure; (iv) equilibrium is estab-lished at the gas/liquid interface, with the CO2 concentration inthe liquid phase described using Henry’s Law; (v) the thicknessof the ionic liquid film and the liquid viscosity are constantand spatially uniform; and (vi) the CO2 diffusion coefficient is

* To whom correspondence should be addressed. Tel.: 315-268-2368. Fax: 315-268-6654. E-mail address: [email protected].

8166 Ind. Eng. Chem. Res.2007,46, 8166-8175

10.1021/ie070501u CCC: $37.00 © 2007 American Chemical SocietyPublished on Web 10/18/2007

independent of CO2 concentration. Combining Fick’s first lawof diffusion with a mole balance written on the liquid-phaseyields

whereCCO2 is the molar concentration of CO2 in the ionic liquid.The initial and boundary conditions are given as follows.

Initial conditions:

Boundary conditions:

whereHCO2 is the Henry’s Law constant for CO2 in the ionicliquid, FIL the density, andMWIL the molecular weight of theionic liquid. The solution of eq 1, subject to the conditions ineq 2, is

A mole balance written on the gas phase yields

whereV is the volume of gas andVIL is the volume of the ionicliquid in the chamber. Substituting the derivative (∂CCO2/∂z)z)L

from eq 3 into eq 4 yields

Rearranging eq 5 and integrating fromt ) 0 to t yields

where

Using a nonlinear least-square method, the two unknownvariables HCO2 and DCO2 were simultaneously fit to theexperimentalP(t) data. From an examination of results obtainedusing a different number of terms in eq 6, it was found that 55terms in the summation in eq 6 were sufficient to determineHCO2 andDCO2 accurately.13 A comparison of theHCO2 andDCO2

values determined from data collected over different timeintervals shows that results obtained using the first∼8 h ofdata are essentially the same as results obtained using longertime intervals.13

A similar experimental approach was used by Camper et al.14

to measure the solubility and diffusivity of different gases,including CO2, in ethylmethylimidazolium bis(trifluoromethyl-sulfonyl)imide ([emim][Tf2N]). The experiments performed byCamper involved a larger volume of ionic liquid (∼6 mL).Pressure decay data were collected for the first 20 min of theexperiment, during which time the liquid was not stirred. Aftersufficient data were collected to determine the solute diffusioncoefficient, the ionic liquid was stirred to accelerate the approachto equilibrium. Gas solubility was determined from the differ-ence between initial and equilibrium pressures in the system.A model for diffusion into a semi-infinite volume was used todetermine the solute diffusion coefficient in the ionic liquid fromthe initial pressure data. In the model used by Camper tointerpret the initial transients, it was assumed that the penetrationdepth of CO2 into the ionic liquid is short and that CO2concentration in liquid far from the interface remains at zero.This differs from the model used in this study where CO2

penetrates into the entire liquid film. Later in this paper, wewill discuss the advantages and disadvantages of these, as wellas other experimental approaches, to determine the solubilityand diffusivity in ionic liquids.

3. Experimental Methods

3.1. Materials. 3.1.1. Ionic Liquids. Five different ionicliquids were studied in this work: 1-n-butyl-3-methylimidazo-lium bis(trifluoromethylsulfonyl)imide ([bmim][Tf2N]), 1,2-dimethyl-3-propylimidazolium bis(trifluoromethylsulfonyl)imide([pmmim][Tf2N]), 1-butyl-3-methylpyridinium bis(trifluoro-methylsulfonyl)imide ([bmpy][Tf2N]), 1-(3,4,5,6-perfluorohexyl)-3-methylimdazolium bis(trifluoromethylsulfonyl)imide ([perfluoro-hmim][Tf2N]), and 1-n-butyl-3-methylimidazolium tetrafluoro-borate ([bmim][BF4]). The ionic liquids [bmim][Tf2N], [pmmim]-[Tf2N], and [bmpy][Tf2N] were obtained from Covalent As-sociates (Woburn, MA) (electrochemical grade:>99.5% purity,<50 ppm H2O). The [bmim][BF4] was purchased from SigmaAldrich (St. Louis, MO) and had a minimum purity of 97%and water content of<0.05 wt %. The [perfluoro-hmim][Tf2N]sample (halide concentration [Br-], <10 ppm; [NH4

+],<20 ppm) was kindly supplied by the Brennecke group at theUniversity of Notre Dame.15

3.1.2. Gases.High-purity nitrogen (N2) and CO2 wereobtained from Merriam-Graves Co. (Charlestown, NH), withpurities of 99.998% and 99.995%, respectively. CO2 was usedas the absorption gas, and N2 was used as the desorption gas toregenerate the ionic liquid.

3.2. Apparatus and Procedures. 3.2.1. CO2 Uptake Mea-surements.A schematic of the experimental system used forCO2 uptake measurements is shown in Figure 1. The systemwas designed so that CO2 absorption measurements could berepeated several times with the same ionic liquid sample, withregeneration in a N2 atmosphere and under vacuum betweenCO2 uptake measurements. Uptake experiments were performedin a stainless steel sample chamber that was a four-way

∂CCO2

∂t) -DCO2

∂2CCO2

∂z2(1)

t ) 0, P ) P0, CCO2) 0 (2a)

t > 0, z ) 0,∂CCO2

∂z) 0 (2b)

z ) L, CCO2)

FIL

HCO2MWIL

P ) K‚P (2c)

CCO2

CCO2|z)L

) 1 - (4

π) ∑n)0

∞ (-1)n

(2n + 1)cos((2n + 1)πz

2L ) ×

exp[-(2n + 1)2π2DCO2

t

4L2 ] (3)

dPdt

) RTV

dndt

) RTV (VIL

L )(-DCO2

∂CCO2

∂z |z)L) (4)

dP

dt)

-RT

V (VIL

L )DCO2‚KP‚(2

L) ∑n)0

∞

exp[-(2n + 1)2π2DCO2

t

4L2 ] (5)

lnP

P0

)

( k

HCO2

) ∑n)0

∞ 1

(2n + 1)2{exp[-

(2n + 1)2π2DCO2t

4L2 ] - 1} (6)

k )8RTVILFIL

π2VMWIL

Ind. Eng. Chem. Res., Vol. 46, No. 24, 20078167

Swagelok cross with1/4 in. fittings (chamber volume) 8.7 mL).An ultrahigh accuracy pressure transducer (Sensotec Sensors,model FP-2000; accuracy of 0.10%, 0-50 psia) was installedin the top fitting and a plug was installed at the bottom. ValvesK4 and K5 were used to connect the inlet and exit lines. Theoutput from the pressure transducer was sent to a USB-baseddata acquisition board and was displayed and recorded on acomputer with TacerDAQ software. The chamber was immersedin a temperature-controlled water bath (Fisher Scientific ISO-temp Refrigerated Circulator, Model 901). The temperature onthe outside of the sample chamber was measured using a Fluke51 II digital thermometer with laboratory-grade accuracy (to0.05%). The chamber was evacuated to vacuum conditionsbefore each experiment, using a vacuum pump (Savant GelPump, model GP110) with an ultimate vacuum level of 7 Torr(9 mbar). Gas was introduced into the sample chamber from afeed tank that was a 150-mL stainless steel cylinder (HOPE,model DOT 3E1800). The feed tank pressure was controlledusing valve K3, with the tank pressure indicated with a digitalpressure gauge (Omega model DPG1000B, accuracy of 0.25%,-30 in. Hg (or∼100 psig)).

The experimental procedure involved the following steps.With the top port opened, a known quantity (∼0.25-0.28 mL)of ionic liquid was delivered to the sample chamber using amicropipette. The chamber was then sealed and the nitrogenline was connected to valve K3. With valve K2 closed, thesystem was evacuated using the vacuum pump until the chamberpressure reached a constant value. Valves K4 and K5 were thenclosed and nitrogen was introduced into the feed tank. ValveK3 was then closed and valve K4 was opened, introducingnitrogen into the sample chamber. Valve K4 was closed andthe pressure in the sample chamber was monitored by thepressure transducer for 12 h. If the pressure of the chamber didnot change over this period, it was confirmed that there was noleak and that N2 solubility in the ionic liquid was negligible.

After the system was leak-tested, the ionic liquid sample washeated to 70oC under vacuum for at least 8 h. The CO2 linewas then connected to valve K3. With valve K3 closed and valvesK4 and K5 open, the system was again evacuated, using thevacuum pump, until the entire line was under maximum vacuumconditions (approximately-29.9 in Hg). Valve K4 and K5 werethen closed, the vacuum pump was turned off and the waterbath temperature was set to the desired value. Valves K1 andK3 were opened, introducing CO2 into the feed tank until the

pressure of the feed tank reached the desired pressure (1-2 bar).Valves K3 and K1 were then closed and valve K4 was quicklyopened and closed, introducing CO2 into the sample chamber,with an initial pressure of∼2 bar. The pressure in the samplechamber decreased with time, because of CO2 absorption intothe RTIL. Pressure was recorded at 1 s intervals. The durationof each experiment was 12-14 h, depending on the ionic liquidand the system temperature.

Following each experiment, the ionic liquid was regeneratedby introducing N2 into the sample chamber, followed by systemevacuation at high temperature (70°C) to draw the absorbedCO2 out of the ionic liquid. The chamber was purged with N2

for a 30 min interval and then the system was evacuated for 1h. Experimental results showed that, after repeating this N2

introduction and evacuation sequence six times, the ionic liquidwas totally regenerated and a new experiment could beconducted.

For each ionic liquid sample, 2 or 3 experiments wereperformed at each temperature. With the exception of [perfluoro-hmim][Tf2N], experiments were repeated with each ionic liquidusing two different ionic liquid samples with different filmthickness (2-3 mm). Agreement between diffusion coefficientvalues obtained with films of different thicknesses confirms theassumptions used in the model.

3.2.2. Ionic Liquid Density Measurements.The density of[bmpy][Tf2N] was measured using a 1-mL Gay-Lussac pyc-nometer (Thomas Scientific, Swedesboro, NJ) at temperaturesof T ) 10, 15, 20, 25, 30, 35, 40, 45, and 50°C. The precisevolume of the pycnometer was determined using measurementswith water performed at room temperature. The validity of thisvalue was confirmed by additional measurements that wereperformed with water at other temperatures. Because of thelimited amount of [perfluoro-hmim][Tf2N], the density of thisionic liquid was determined by weighing a known volume ofliquid that was charged using a 10-µL syringe (Hamilton, Reno,NV). Measurements were repeated using different volumes ofthe ionic liquid. The density of the ionic liquid was determinedfrom the slope of a plot of sample mass versus ionic liquidvolume delivered. These measurements were performed at roomtemperature only (23°C).

Literature values for the densities of the other ionic liquidswere used to determine the Henry’s Law constant from pressuredecay data. The densities of [bmim][Tf2N] and [pmmim][Tf2N]

Figure 1. Schematic diagram of the experimental system used for the solubility and diffusivity measurements.

8168 Ind. Eng. Chem. Res., Vol. 46, No. 24, 2007

were obtained from Fredlake et al.,16 and the density of [bmim]-[BF4] was obtained from Jacquemin et al.17

3.2.3. Ionic Liquid Viscosity Measurements.The viscositiesof ionic liquids [bmim][Tf2N] and [pmmim][Tf2N] were mea-sured by a semi-micro extra-low-charge capillary viscometer(0.4 mL, Cannon Instrument Co., State College, PA). Kinematicviscosity was determined by measuring the efflux time requiredfor liquid to flow between marks on the viscometer. Theviscometer was calibrated by the manufacturer, who providedan expression relating the calibration constant to temperature.Dynamic viscosity was determined from the kinematic viscosityusing measured or literature values for ionic liquid density.Additional details can be found in the thesis of Y. Hou.13

Measurements were repeated at temperatures of 5-50 °C.Literature values for the viscosity of [bmpy][Tf2N]17 and of[bmim][BF4]15 were used when interpreting the diffusioncoefficient results.

4. Results and Discussion

4.1. Density.The density values determined for [bmpy][Tf2N]are plotted as a function of temperature in Figure 2. A linear fitof F vs T yields

where the temperatureT is given in degrees Celsius. Measureddensity values are provided in the Supporting Information. Thedensity of [perfluoro-hmim][Tf2N] measured at room temper-ature (23°C) was 1.589 g/cm3.

4.2. Viscosity.The viscosity values determined for [bmim]-[Tf2N] and for [pmmim][Tf2N] are presented as an Arrheniusplot in Figure 3. Viscosity values are provided in the SupportingInformation. The activation energy values for viscosity are listedin Table 1. Activation energies for viscosity of [bmpy][Tf2N]and [bmim][BF4] determined from values reported in theliterature are also included in Table 1. Because of the limitedsample available, it was not possible to determine the viscosityof [perfluoro-hmim][Tf2N].

4.3. CO2 Uptake Measurements.A representative plot ofpressure (P) versus time (t) for CO2 uptake into [perfluoro-hmim][Tf2N] is shown in Figure 4. Experimental results werefit to the model prediction (eq 6) using a MATLAB program,yielding values forHCO2 andDCO2. Comparison of experimentalpressure measurements to the model prediction shows excellentagreement over the entire time of the experiment, confirmingthe validity of the model for these systems.

4.4. CO2 Solubility. Values for the Henry’s Law constantfor CO2 in the tested ionic liquids are listed in Table 2. It ishelpful to note here that smallerHCO2 values correspond tohigher solubility at the same pressure. The Henry’s Lawconstants for CO2 in [bmim][Tf2N] and [bmim][BF4] measuredin this work are in good agreement with values that have beenreported by other researchers, using different experimentaltechniques,18-21 which provides validation for our experimentalmethod. A comparison of CO2 solubilities in [bmim][Tf2N] and[bmim][BF4] in Table 2 clearly shows that CO2 is considerablymore soluble in [bmim][Tf2N] than in [bmim][BF4]. This isconsistent with observations noted by others that CO2 solubilityis higher in ionic liquids with Tf2N- anions than in ionic liquids

Figure 2. Density (F) as a function of temperature (T) for [bmpy][Tf2N].The line is a linear regression fit of the data (eq 7).

Figure 3. Semilogarithmic plot of viscosity (η) versus inverse temperature(1/T).

F (g/cm3) ) 1.436- 8.995× 10-4T (7)

Figure 4. Comparison of the (×) experimental measurements and (s)model prediction of dimensionless pressure (P/P0) versus time (t) for CO2

absorption in [perfluoro-hmim][Tf2N] at 23 °C. The results from oneexperiment are shown. The system pressure was recorded each second;however, this plot only shows every 20th point for the first hour and every120th point for times of more than one hour.

Table 1. Activation Energy (Ea) Values of the Viscosity of IonicLiquids

ionic liquid E′a (kJ/mol)

[bmim][Tf2N] 32 ( 1[pmmim][Tf2N] 35 ( 2[bmpy][Tf2N] 34 ( 3a

[bmim][BF4] 29 ( 1b

a From ref 15.b From ref 17.


with BF4- anions.18,19,22,23The results in Table 2 also show that

ionic liquids with the Tf2N- anion have similar CO2 solubilities,implying that the ionic liquid anion has a stronger effect onCO2 solubility than does the cation, which has also beenobserved by other researchers.18

Although our results show that CO2 solubility is influencedmore strongly by the ionic liquid anion than the cation, onecan draw a number of conclusions about the effect of cationstructure from a comparison of theHCO2 values listed in Table2. The CO2 solubility results for ionic liquids with Tf2N- anionfollow the trend

The cations [pmmim] and [bmim] have the same molecularformula but exhibit different CO2 solubility. Here, the extramethyl group on [pmmim] and the shorter alkyl chain (C3 vsC4) reduces the free volume available for CO2. The combinationof these steric factors leads to a slightly higher CO2 solubilityin [bmim], compared to [pmmim]. The higher CO2 solubilityin [bmpy], compared to that in [bmim], can be explained byconsidering the fact that the larger pyridinium ring can distributeand stabilize the cation charge better than the imidazolium ring,resulting in weaker cation-anion interactions. These weakerinter-ion interactions can increase interactions between CO2 andthe anion, resulting in an increase in CO2 solubility. Of the ionicliquids investigated here, the highest CO2 solubility wasobserved with [perfluoro-hmim][Tf2N]. The reason for this isnot entirely clear. As has been reported from other studies, anincrease in CO2 solubility generally results when the length ofthe alkyl side chain on the cation is increased, because of theincreased free volume available for CO2, as well as a decreasein cation-anion interactions.22, 24 However, with [perfluoro-hmim] as the cation, the increase in free volume resulting fromthe long alkyl side chain is countered by the large F atoms onthat side chain. It is possible that the F atoms also impact thedistribution of charges on the cation, influencing the strengthof the cation-anion interactions.

A systematic comparison of the effect of the number, size,and position of alkyl groups on CO2 solubility in imidazolium-based ionic liquids is shown in Table 3. For this comparison,literature values forHCO2 for ionic liquids that were not studiedin this work have been included. A comparison ofHCO2 for[emim] to [emmim] and for [pmim] to [pmmim] shows that anadditional methyl group on the imidazolium ring reduces theCO2 solubility. The extra methyl group in [emmim] and[pmmim] should not significantly affect the distance and,therefore, the strength of the attraction between the cation andthe anion. However, the extra methyl group reduces the freevolume available for CO2, resulting in a decreased solubility.When one considers the large uncertainty in theHCO2 valuereported for [pmim][Tf2N], the results summarized in Table 3are generally consistent with the trend reported in previousstudies that CO2 solubility increases as the length of the alkylchain on [Cnmim][Tf2N] ionic liquids increases.22,24As discussed

previously, these observations can be explained by the increasein free volume and the decrease in cation-anion interactionsthat results as the length of the alkyl side chain is increased.

The effect of temperature on CO2 solubility can be relatedto the partial molar entropy and partial molar enthalpy ofabsorption:25

The values for∆sjCO2 and ∆hhCO2 were determined from ap-propriate plots ofHCO2 vs T, and the results are listed in Table4. The results show a generally inverse relationship between∆sjCO2, ∆hhCO2, and CO2 solubility. This trend follows one thatwas observed by Anthony et al.,18 where the∆sjCO2 and∆hhCO2

values for CO2 in ionic liquids with [BF4] and [PF6] anionswere larger than enthalpy and entropy values for CO2 in ionicliquids with [Tf2N] anions, whereas solubility was highest inthe [Tf2N] ionic liquids. The results with [pmmim][Tf2N] arean exception to this trend, with∆sjCO2 and∆hhCO2 being smallerthan those respective values for CO2 in [bmim][Tf2N], wherethe CO2 solubility is somewhat smaller. It seems that the extramethyl group on the imidazolium ring may impact the orderingof CO2 in the ionic liquid, which is not unexpected. The partialmolar enthalpies of absorption in [pmmim][Tf2N] and[bmim]-[Tf2N] are quite similar, which indicates similar interaction

Table 2. Henry’s Law Constants of CO2 in Ionic Liquids a

Henry’s Constant,HCO2 (bar)

ionic liquid at 10°C at 20°C at 25°C at 30°C at 40°C at 50°C

[bmim][ Tf2N] 28 ( 2 30.7( 0.3 34.3( 0.8 42( 2 45( 3 51( 2[pmmim][Tf2N] 29.6( 0.6 34( 3 38.5( 0.9 40.4( 0.6 46( 3 53( 2[bmpy][Tf2N] 26 ( 1 31.2( 0.1 33( 1 35( 2 41( 4 46( 1[perfluoro-hmim][Tf2N] 25.5( 0.2 29.2( 0.4 31( 2 32( 2 36( 4 42( 2[bmim][BF4] 41.9( 0.2 52( 2 56( 2 63( 2 73( 1 84( 4

a Uncertainty limits represent 95% confidence limits.

[pmmim] < [bmim] < [bmpy] < [perfluoro-hmim]

Table 3. Comparison of CO2 Solubility in Imidazolium IonicLiquids at 25 °C

a Values were measured in this work; others are literature values.

Table 4. Partial Molar Entropies and Enthalpies of CO2

Absorption, and Henry’s Law Constant (H), in Ionic Liquids

ionic liquid∆sjCO2

(J/(K mol))∆hhCO2

(kJ/mol)H @ 25°C

(bar)

[perfluoro-hmim][Tf2N] -29 ( 5 -9 ( 2 31( 2[bmpy][Tf2N] -34 ( 3 -10.4( 0.9 33( 1[bmim][Tf2N] -40 ( 13 -12 ( 4 34.3( 0.8[pmmim][Tf2N] -36 ( 5 -11 ( 1 38.5( 0.9[bmim][BF4] -44 ( 5 -14 ( 1 56( 2

∆sjCO2) -R(∂ ln HCO2

∂ ln T )P

(8)

∆hhCO2) R(∂ ln HCO2

∂(1/T) )P

(9)


forces between CO2 and these ionic liquids. However, theseexplanations can only be considered speculative, given theuncertainties in the enthalpy and entropy values.

4.5. CO2 Diffusivity. Diffusion coefficient values for CO2in the tested ionic liquids are presented as an Arrhenius plot inFigure 5. All measured values are listed in the SupportingInformation.

The validity of the results obtained with the transient thin-film method used in this study is dependent on severalassumptions that were made when developing the model usedto interpret the pressure decay data. It is assumed that the liquidfilm is of uniform thickness across the sample chamber. Visualinspection of the plug used to form the bottom of the samplechamber shows no obvious nonuniformity, except near thethreads. The fact that the diffusion coefficient values determinedfrom experiments performed with ionic liquid samples withdifferent film thicknesses were in good agreement with eachother (Figure 5) indicates that any nonuniformity in filmthickness was not significant. It is also assumed that gastransport in the ionic liquid is governed solely by moleculardiffusion, with no contributions from convection. Two observa-tions support this assumption. If natural convection due totemperature gradients was important, we would expect pooreragreement between experimental pressure values and modelpredictions for experiments run at temperatures farther fromroom temperature, when compared to experiments run at25 °C. There was no obvious difference in the ability of themodel to describe the pressure decay data for experiments runat different temperatures. If convective currents were createdwhen gas was initially introduced into the sample chamber, itis expected that (i) initial experimental pressure values wouldnot follow model predictions and (ii) better agreement wouldbe observed as the convective currents decay with time as thetransport mechanism changes to a purely diffusion controlledprocess. The good agreement between model and observationover the entire course of an experiment indicates that CO2

transport into the thin films was governed solely by moleculardiffusion. In summary, we are confident that the conditions usedfor our measurements are consistent with the assumptions usedin developing the model used to interpret the pressure decaydata.

A comparison of the diffusion coefficient values in thedifferent ionic liquids shows thatDCO2 is largest in [bmim][Tf2N]and smallest in [bmim][BF4]; the diffusion coefficients in[pmmim] [Tf2N] and [bmpy] [Tf2N] are comparable. A summary

of the activation energy (Ea) values for CO2 diffusion ispresented in Table 5. For comparison,Ea values that weredetermined from literature values for CO2 diffusion in thetraditional organic solvents methanol and isooctane are alsoincluded.26 Activation energy values for CO2 in ionic liquidswith [Tf2N] anion are slightly greater than theEa values forCO2 diffusion in traditional solvents, while the activation energyfor CO2 diffusion in [bmim][BF4] is smaller than that for CO2diffusion in methanol and isooctane.Ea values determined inthis study are in good agreement with the value of 16.9 kJ/molfor CO2 diffusion in [emim][Tf2N], which was reported byMorgan et al.,27 and are smaller than the value of 22.3 kJ/molthat was reported by this same group for CO2 diffusion in[bmim][C2F6(SO3)2N].

Using nuclear magnetic resonance (NMR), self-diffusioncoefficients of anions and cations in several ionic liquids havebeen measured at different temperatures.28,29Activation energiesfor self-diffusion were calculated from the results, yieldingEa

) 27.2 kJ/mol for [bmim] andEa ) 28.3 kJ/mol for [Tf2N] in[bmim][Tf2N], while Ea ) 38.4 kJ/mol for [bmim] andEa )40.5 kJ/mol for [PF6] in [bmim][PF6]. These values areconsiderably larger than theEa values for CO2 diffusion in ionicliquids and are similar to theEa values determined from viscositymeasurements (see Table 1). Similar molecular motion seemsto be involved with the anion and cation self-diffusion and ionicliquid viscosity, and these mechanisms are considerably differentthan the molecular motion required for CO2 diffusion in thesematerials.

There have been a few reports in the literature from othermeasurements of CO2 diffusivity in ionic liquids.14,27,30In Figure6, these literature values are compared to selected diffusivityvalues measured in this work. The ionic liquid [bmim][BF4] isthe only ionic liquid used in this work with CO2 diffusivity

Figure 5. Semilogarithmic plot ofDCO2 vs 1/T. Different symbols with the same ionic liquid represent results with different ionic liquid samples withdifferent film thicknesses.

Table 5. Activation Energy (Ea) Values of CO2 Diffusivity in IonicLiquids and Traditional Solvents

liquid Ea (kJ/mol)

[bmim][Tf2N] 10 ( 3[pmmim][Tf2N] 15 ( 4[bmpy][Tf2N] 12 ( 1[perfluoro-hmim][Tf2N] 15 ( 2[bmim][BF4] 6 ( 1methanola 9.8( 0.5isooctanea 8.9( 0.3

a Data taken from ref 26.


measured by other researchers. Shiflett and Yokoseki reporteddiffusion coefficient values for CO2 in [bmim][BF4] that weredetermined using a gravimetric microbalance.30 In these mea-surements, mass uptake of CO2 into an ionic liquid film wasmeasured and the diffusion coefficient was determined usingthe same thin-film model that was used in this study. To interpretmass changes in terms of solute uptake accurately, it is necessaryto make buoyancy corrections for several factors, and themagnitude of these correction factors can be significant. Thediffusion coefficients reported by Shiflett and Yokoseki30 areconsiderably smaller than the values measured in this study.Shiflett and Yokoseki30 also reported anEa value for CO2

diffusion in this ionic liquid which is larger (24.3( 4.2kJ/mol) than the values determined from our measurements. TheEa value reported in their study is similar to the values reportedfor anion and cation self-diffusion and for viscosity (see Table1). Note that the diffusion coefficient reported by Shiflett andYokoseki30 for CO2 diffusion in [bmim][PF6] is also consider-ably smaller than the value reported by Morgan et al.,27 whoused a membrane lag-time technique. There are many possiblereasons for these differences. It is recognized that the propertiesof ionic liquids with BF4

- and PF6- anions are more sensitiveto the presence of water than are ionic liquids with the Tf2N-

anion. It is possible that the experimental conditions for themeasurements performed by Shiflett and Tokoseki30 weresufficiently different than the conditions used for our measure-ments or those measurements that were used by Morgan et al.27

The measurements performed by Shiflett and Tokoseki30 wereconducted at pressures up to 2.0 MPa. The larger CO2 contentin the ionic liquids at higher pressures can affect the densityand viscosity of the ionic liquid. In the model that is used tointerpret mass changes, it is assumed that the physical propertiesof the liquid film are spatially uniform and are constant withtime, assumptions that are questionable at higher CO2 pressures.Therefore, theD values reported by Shiflett and Tokoseki30 canbe considered average, effective values with higher uncertaintiesthan the values determined from our low-pressure measurements.

Using a semi-infinite volume approach that was discussedearlier, Camper et al. measured the diffusivities of gases,including CO2, in several ionic liquids, all of which are differentthan those examined in this study.14 The results from the Camperet al.14 measurements for CO2 in [emim][Tf2N] are included inFigure 6. A comparison ofDCO2 for CO2 in [emim][Tf2N] toDCO2 values measured in this study with similar ionic liquidsshows excellent agreement.

4.5.1. Comparison of Different Experimental Methods.Ashas been noted in some of the previous discussion, severaldifferent methods have been developed for measuring gassolubilities and diffusivities in ionic liquids. These include anequilibrium pressure method,12,14,31-35 a gravimetric method,30,32,36

a membrane lag-time technique,27 and a semi-infinite volumeapproach.14 The equilibrium pressure method involves measur-ing gas pressure above an ionic liquid at the time when gas isinitially introduced and again after equilibrium has beenachieved. This method only provides solubility information, andthere can be considerable uncertainty in values because of thelimited amount of data that has been collected. The accuracycan be improved by repeating experiments with different initialpressures. However, this involves multiple experiments and canbe quite time-consuming.

In the transient thin-film method developed in this study, onecan determine both the Henry’s Law constant as well as targetgas diffusivity in a single experiment. It is also possible todetermine solubility without waiting for the system to achieveequilibrium. The results obtained using the first∼8 h of dataare essentially the same as results obtained using the entirepressure curve to equilibrium.13 The validity of the parametersdetermined from the pressure decay data is dependent onassumptions made in the model development, i.e., uniform filmthickness, constant physical properties, and no convectivetransport. With careful design of the experimental system, theseassumptions can be satisfied, as discussed previously in thispaper. The simplicity of the experimental apparatus used, andthe fact that only pressure data are needed, provide many otheradvantages for this method over other approaches. Other factors,such as the buoyancy correction in the gravimetric method ormembrane tortuosity in the lag-time technique do not come intoplay in this technique.

The transient thin-film method is similar to the semi-infinitevolume method used by Camper et al.,14 except that differentionic liquid sample sizes are used and a different one-dimensional model is used to interpret pressure decay data. Theprimary advantage of the method used in this study over thesemi-infinite volume approach is that a significantly smallersample size is required. In addition, the Henry’s law constantdetermined from a data fit to the transient thin-film diffusionmodel is more accurate than the solubility value determinedfrom only considering differences between initial and finalsystem pressures, because the entire pressure decay curve isused when determiningHCO2. Although the experiments that

Figure 6. Semilogarithmic plot ofDCO2 versus 1/T. Selected results from this study are compared to results reported by other researchers using differentexperimental techniques. Values from this study are average values from measurements with two different ionic liquid samples.


are performed with the ionic liquid film require significantlylonger time than those involving the larger volume, theexperiments do not require careful monitoring after they havebegun.

4.5.2. Diffusion Correlations.Many previously developeddiffusion correlations for solute diffusion in traditional liquidsrelate the diffusion coefficients to solute molar volume ordensity, solvent viscosity, and solvent molecular weight. Theseare well-summarized by Morgan et al.27 A commonly appliedempirical correlation for liquid solvents is the Wilke-Changequation, which is written here for the diffusion of CO2 in anionic liquid:

whereVCO2 is the molar volume of CO2 andæ is an associationfactor that characterizes the solvent-solvent interactions.37

Values foræ range from 1.0 for unassociated solvents to 2.6for water. The diffusion coefficients measured in this study wereused to determine a value foræ for each ionic liquid from aplot of DCO2 vs T/ηIL, usingVCO2 ) 34 cm3/mol.38 Calculatedassociation factors ranged from 17.5 for [bmpy][Tf2N] to 33.3for [pmmim][Tf2N]; these values are 10 times larger than thosefound to represent association in more-traditional solvents.13

However, even using the fittedæ values, there is poor agreementbetween the measured and predicted diffusion coefficients. Onepossible explanation for the poor agreement is thatæ may bedependent on temperature for these systems.13 Also note thatMorgan et al.27 reported a valueæ ) 0.15 from their measure-ments of gas diffusion in different imidazolium-based ionicliquids.

Hayduk and Cheng38 proposed a simple correlation relatingsolute diffusivity to solvent viscosity:

A log-log plot of DCO2 vs ηIL of the results for CO2 diffusionin [bmim][Tf2N], [pmmim][Tf2N], [bmpy][Tf2N], and [bmim]-[BF4] are shown in Figure 7. Diffusivity values for CO2 diffusionin [perfluoro-hmim][Tf2N] are not included, because viscositydata are not available for this ionic liquid. A regression fit of

all of the results presented in Figure 7 to eq 11 yieldsA ) 3.6× 10-5 andB ) -0.44. These values are very similar to thosereported by Hayduk and Cheng for CO2 diffusion in conven-tional liquids (A ) 3.5 × 10-5 and B ) -0.44).38 However,the overall fit of this expression to all of the data is not good.Clearly, there are differences in behavior between the [Tf2N]ionic liquids and the [BF4] ionic liquid that this simple modelis not capturing. Other established correlations were alsoconsidered, but none were satisfactory in explaining our results.13

The results from this study were used to develop a diffusioncorrelation equation for CO2 diffusion in ionic liquids. Althoughthe molar volume of solute is a parameter in many correlationequations, this parameter is a constant here, because CO2 is theonly solute that was investigated. Other important parametersinclude the viscosity, density, and molecular weight of the ionicliquid. Because our measurements were performed at differenttemperatures, temperature was included in our analysis. Usingall of the measured CO2 diffusivity values (except those for[perfluoro-hmim][Tf2N]), the following correlation was devel-oped:

where the uncertainty limits on the coefficients are the 95%confidence limits. Our analysis indicates that all of theparameters included in eq 12 are statistically significant inpredictingDCO2. A comparison of experimental and predicteddiffusion coefficients from this model is shown in Figure 8.All predicted diffusion coefficients are within 15% of theexperimental values, and most of the data are within 10% ofthe experimental values measured in this work. Other expres-sions were considered, with different collections of ionic liquidproperties. However, eq 12 is the correlation with the beststatistical fit to our measurements.

Scovazzo and co-workers27,39have developed several differentcorrelations for gas diffusivities in different ionic liquids, usingdata that were determined using a membrane lag-time technique.Data were collected with ionic liquids with both imidazolium-based and phosphonium-based cations and a variety of differentanions. These correlations were fit to isothermal data. Therefore,temperature was not included in their analyses and the depen-

Figure 7. Log-log plot of DCO2 versus ionic liquid. Each data point represents the average of the diffusion coefficient values obtained with two differentionic liquid samples. The power-law fit (eq 11) was developed using results for all four ionic liquids.

DCO2)

7.4× 10-8 xæMILT

ηILVCO2

0.6(10)

DCO2) AηIL

B (11)

DCO2) 6.7× 105ηIL

-0.66(0.12MWIL-0.89(0.57×

FIL4.8(2.0T-3.3(1.6 (12)


dence on ionic liquid viscosity and density only incorporateddifferences between ionic liquids, not changes in properties withtemperature. Because different solute gases were investigated,these correlations also incorporated solute size effects. Diffu-sivity was determined to be inversely proportional to theviscosity of the ionic liquid, with a power of 0.6 for imidazoliumionic liquids (essentially the same as the sensitivity determinedfrom our measurements). However, Ferguson and Scovazzo39

determined diffusivity to be less sensitive to viscosity for gasdiffusion in phosphonium-based ionic liquids, with a power of-0.35. The results reported by Morgan et al.27 (primarily withimidazolium-based ionic liquids) indicated essentially the samesensitivity to solvent molar volume and molecular weight,enabling them to collapse these two properties into a power-law dependence on density, with a power of 2. This is inreasonable agreement with our results, given the uncertainty inthe power onFIL in eq 12. The molecular weight of the ionicliquid was determined to be an insignificant predictor of gasdiffusivities in the phosphonium-based ionic liquids, with resultsshowing a 0.6 power dependence on the ionic liquid molarvolume.39

The inverse relationship between diffusivity and temperaturein eq 12 is counter-intuitive. It is important to note that ourmeasurements were performed over a limited temperature range.In addition, the fact that ionic liquid viscosity shows a strongsensitivity to temperature (Table 1) may be contributing to thisunusual relationship. More data must be collected over a broaderrange of temperatures and ionic liquid properties to develop amore-comprehensive correlation for the dependence of diffu-sivity on the properties of the ionic liquid.

5. Conclusions

In this study, a transient thin-film method was developed todetermine both the solubility and diffusivity of CO2 in ionicliquids simultaneously. Measurements were performed at lowCO2 pressures, providing us with parameter values under infinitedilution conditions. As reported previously by others, CO2

solubility was more strongly dependent on the ionic liquid anionthan the cation, and the solubility in ionic liquids with the [Tf2N]anion was observed to be higher than that in the ionic liquidswith the [BF4] anion. The entropies of CO2 absorption rangedfrom -29 J/(K mol) to-44 J/(K mol), and the enthalpies rangedfrom -9 kJ/mol to-14 kJ/mol, following an inverse relation-ship with solubility. The magnitude of the diffusion coefficientsof CO2 in the ionic liquids is 10-6 cm2/s, with the activation

energies being slightly higher than those for CO2 diffusion intraditional organic solvents. These activation energies areconsiderably smaller than the activation energies for viscosityand anion and cation self-diffusion. Diffusion coefficients wereobserved to be more sensitive to temperature than observed forsolubility. Measured diffusion coefficients are observed to bein excellent agreement with a correlation relating diffusivity toionic liquid viscosity, density and molecular weight, and thesystem temperature.

Nomenclature

A ) constant defined by eq 11B ) constant defined by eq 11CCO2 ) concentration of CO2 in an ionic liquid (mol/cm3)DCO2 ) diffusion coefficient of CO2 in an ionic liquid (cm2/s)E′a ) activation energy for viscosity (kJ/mol)Ea ) activation energy of the diffusion coefficient (kJ/mol)∆hhCO2 ) partial molar enthalpy of CO2 absorption (kJ/mol)HCO2 ) Henry’s law constant of CO2 in an ionic liquid (bar)L ) depth of ionic liquid (cm)MWIL ) molecular weight of ionic liquid (g/mol)n ) number of moles of CO2 (mol)P ) total pressure in the vapor phase (bar)P0 ) initial pressure of CO2 in the vapor phase (bar)R ) gas constant (8.31434 m3 Pa mol-1 K-1)∆sjCO2 ) partial molar entropy of CO2 absorption (J/(K mol))t ) time (s)T ) temperature (K)V ) volume of the gas (cm3)VIL ) volume of ionic liquid (cm3)VCO2 ) volume of CO2 (cm3)z ) Cartesian coordinate, direction (cm)

Greek Symbols

ηIL ) viscosity of ionic liquid (cP, mPa s)FIL) density of ionic liquid (g/cm3)æ ) association factor of solvent in the Wilke-Chang equation

Acknowledgment

We acknowledge the financial support from Clarkson Uni-versity and the National Science Foundation Grant No. CTS-0522589. Ionic liquid samples for this project were kindlysupplied by Covalent Associates ([bmim][ Tf2N], [pmmim]-[Tf2N] and [bmpy][Tf2N]) and Dr. Joan F. Brennecke at theUniversity of Notre Dame ([perfluoro-hmim][Tf2N]).

Supporting Information Available: Tables listing densityversus temperature data for [bmpy][Tf2N], viscosity versustemperature data for [pmmim][Tf2N] and [bmim][Tf2N], andCO2 diffusion coefficients versus temperature for all ionicliquids. (PDF files.) This material is available free of chargevia the Internet at http://pubs.acs.org.

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ReceiVed for reView April 9, 2007ReVised manuscript receiVed August 24, 2007

AcceptedAugust 28, 2007

IE070501U


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Experimental Measurement of the Solubility and Diffusivity of CO 2 in Room-Temperature Ionic Liquids Using a Transient Thin-Liquid-Film Method