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Chemical Engineering and Processing 49 (2010) 192–198

Contents lists available at ScienceDirect

Chemical Engineering and Processing:Process Intensification

journa l homepage: www.e lsev ier .com/ locate /cep

onic liquids for acetylene and ethylene separation: Material selection andolubility investigation

elliarko Palgunadi ∗, Hoon Sik Kim, Jung Min Lee, Srun Jungepartment of Chemistry, Kyung Hee University, 1-Hoegi-dong, Dongdaemoon-gu, Seoul 130-701, Republic of Korea

r t i c l e i n f o

rticle history:eceived 2 November 2009eceived in revised form7 December 2009ccepted 23 December 2009vailable online 6 January 2010

a b s t r a c t

Potential applications of ionic liquids (ILs) for the green separation process of acetylene in ethylene andfor the storage of acetylene were investigated. To deal with this proposal, the solubilities of the unsatu-rated hydrocarbons in various ionic liquids were evaluated. The solubility of ethylene shows a solubilityparameter-dependent behavior as indicated by the proportional relationship between the natural logvalue of Henry’s law constant and the inverse molar volume of ILs. This correlation suggests the mostimportant role of voids formation within IL to accommodate the solutes and the applicability of regular

eywords:cetylenethyleneas separation

onic liquidsolubility

solution theory to model the solubility behavior. Whereas, in addition to the free-volume contribution ofILs, the solubility of acetylene is largely controlled by a specific solute–solvent interaction as a result fromthe association of the acidic hydrogen character in acetylene and the relative basicity of the anion. Thosetwo different solubility behaviors result in a high absorption selectivity of acetylene over ethylene in thebasic ILs. 1H NMR experiment clearly demonstrated the presence of a substantial interaction betweenthe acetylene and the anion of IL. Interestingly, this solute–solvent interaction is reversible as indicated

tion

by the absorption–desorp

. Introduction

Ethylene obtained mainly from naphtha or from natural gasracking is the simplest olefin and is one of the very importanthemical feedstocks for many syntheses in petrochemical indus-ries such as plastics, solvents, packaging, lubricants, intermediatesnd so on [1,2]. Highly pure ethylene feed (>99.9%) is mostlyequired in industrial applications. For example, a ppm level ofcetylene (>5 ppm) in ethylene can poison Ziegler–Natta catalysturing ethylene polymerizations and can also lower the productuality of the resulting polymers [2]. However, the purification ofthylene involves a rather complicated and costly process becausef the inevitable co-generation of a small amount of highly reactivecetylene with a close boiling point [3].

The elimination of acetylene in ethylene is commerciallyccomplished through an intensive energy process namely par-ial hydrogenation of acetylene into ethylene over a noble metalatalyst such as a supported Pd catalyst [4,5]. However, ethylene

roduced from the hydrogenation of acetylene as well as ethyleneeed can undergo over-hydrogenation reaction producing paraffins,hus resulting in the loss of ethylene. Moreover, palladium catalystuffers from deactivation due to the formation of carbonaceous

∗ Corresponding author. Tel.: +82 2 961 0431; fax: +82 2 966 3701.E-mail address: jelliarko@gmail.com (J. Palgunadi).

255-2701/$ – see front matter © 2010 Elsevier B.V. All rights reserved.oi:10.1016/j.cep.2009.12.009

test of acetylene in [BMIM][Me2PO4].© 2010 Elsevier B.V. All rights reserved.

deposits. Within the green chemistry perspective, it is of greatinterest to develop a benign process for the efficient purificationof ethylene from acetylene.

Acetylene, a triple-bond hydrocarbon, though considered asan impurity in olefins, has its own utilities and highly pureacetylene is also in a high demand. Solvent extraction using anorganic solvent such as DMF (N,N-dimethylformamide) or NMP(N-methylpyrrolidinone) is known as a typically industrial tech-nique to separate acetylene from the gaseous mixture in a crackingprocess [6]. However, extractions using organic solvents are dis-advantageous in terms of the significant loss of the solvent aftermultiple operations and the low solubility selectivity of acetyleneover ethylene. Therefore, it is beneficial to search efficient solventsystems that selectively and reversibly interact with acetylene forthe acetylene/ethylene separation. In addition, acetylene is com-monly available as a compressed gas up to 15 psi and is dissolvedin acetone [7]. The use of volatile organic solvent for the gas stor-age reduces the purity of acetylene because it can exhaust togetherwith the gas stream. Therefore, it is also our motivation to searchhigh capacity, non-flammable, and less volatile solvents suitablefor the acetylene storage technology.

Recently, there has been increasing emphasis on the use of ionicliquids (ILs) as benign media in the separation and purification pro-cesses due to their unique and tunable physicochemical propertiesresulted from the variations or the modifications of the cations andthe anions such as temperature stability, negligible vapor pressure,

J. Palgunadi et al. / Chemical Engineering and Processing 49 (2010) 192–198 193

Table 1Ionic liquids used in this study, physicochemical properties, and ILs sources.

Absorbent CAS MW/g mol−1 �a/g cm−3, T/K = 298 �a/Pa s (T/K) Tdecd/K Assay (%) Source

[DMIM][MeHPO3] 1048377-03-2 192.15 1.24b 0.094b (303) n.a. ∼99 [16][BMIM][MeHPO3] n.a. 234.23 1.14b n.a. n.a. ∼98 [16][BMIM][BuHPO3] n.a. 276.32 1.08b 0.288b (303) n.a. ∼97 [16][DMIM][Me2PO4] 654058-04-5 222.18 1.26c 0.363c (293) n.a. ∼99 [19][BMIM][Me2PO4] 891772-94-4 264.26 1.18c 0.696c (293) n.a. ∼98 [19][BMIM][OAc] 284049-75-8 198.26 1.05 0.309 (303) 493e ≥95 Aldrich[BMIM][TFA] 174899-94-6 252.23 1.22 0.053 (303) 443e ≥98 [20][BMIM][MeSO4] 401788-98-5 250.32 1.21 0.161 (303) 633f ∼99 [21][BMIM][BF4] 174501-65-6 226.02 1.21 0.075 (303) 698g ≥97 Aldrich[BMIM][Tf2N] 174899-83-3 419.37 1.44 0.040 (303) 696g ≥98 Aldrich[BMPyrr][BuHPO3] n.a. 279.36 1.03b n.a. n.a. ∼97 [16][BMPyrr][OAc] n.a. 201.30 0.98 n.a. n.a. ∼98 Solvionic[BMPyrr][Tf2N] 223437-11-4 422.41 1.42 0.059 (303) n.a. ≥98 Aldrich

a Values were collected from http://ilthermo.boulder.nist.gov/ILThermo/pureprp.uix.do.b This work.c Literature [19].

h[ecsdteltavta

2

pafas

2

ap–wpb1maatmbrtos

accuracy 0.25% full scale). Volume of the EC (VEC = 100.0 ± 0.1 mL)was measured by filling it with distilled water at room tem-perature. Volumes of the GR and the rest of the system(VGR+rest = 87.0 ± 0.1 mL) were determined from nitrogen gas

d Onset temperature.e Literature [22].f Literature [23].g Literature [24].

ydrophylicity/hydrophobicity, and miscibility with hydrocarbons8,9]. ILs have been demonstrated as the potential agents for thextraction of aromatic hydrocarbons [9], for the extraction of S-ontaining compounds [10], and for the CO2 and hydrocarbonseparations [11–13]. Active metal ions (primarily silver and copper)issolved in ILs were also tested to recover olefins from paraffinshrough reversible complexation reactions [14,15]. Based on thosexamples, it is very reasonable to expect that a selective ethy-ene/acetylene separation could be successfully performed usinghose liquid salts. In this paper, solubility behaviors of acetylenend ethylene in imidazolium- and pyrrolidinium-based ILs bearingarious types of anion are evaluated and some crucial considera-ions for the green and effective separation of these hydrocarbonsre briefly discussed.

. Materials and methods

General procedure: chemical reagents for ILs synthesis withurity no less than 99% were purchased from Aldrich Chemicals Co.,nd were used as received. Ethylene and acetylene were obtainedrom Sin Yang Gas Co., Korea. 1H NMR spectra were recorded on

400 MHz Bruker NMR spectrometer using TMS as the internaltandard.

.1. Synthesis of ionic liquids

Dialkylimidazolium-based room-temperature ILs bearinglkylphosphonate (also known as alkylphosphite) and dialkylphos-hate anion ([RMIM][RHPO3] and [RMIM][R2PO4], where R = –CH3,C4H9) were prepared from the reaction of 1-methylimidazoleith the corresponding dialkyl phosphonate and trialkyl phos-hate, respectively. Similarly, 1-butyl-1-methylpyrrolidiniumutylphosphonate ([BMPyrr][BuHPO3]) was synthesized from-methylpyrrolidine and dibutyl phosphonate. 1-Butyl-3-ethylimidazolium or 1-butyl-1-methylpyrrolidinium containing

cetate, tetrafluoroborate, or trifluoroacetate were commerciallyvailable and were used as received. 1-Butyl-3-methylimidazoliumrifluoroacetate ([BMIM][TFA]) and 1-butyl-3-methylimidazolium

ethylsulfates ([BMIM][MeSO4]) were prepared by reacting 1-

utylimidazole with methyl trifluoroacetate and dimethyl sulfate,espectively. Purities of the synthesized ILs were found to be betterhan 97% as deduced from the proton NMR analysis and puritiesf the purchased ILs were no less than 95% as quoted from theuppliers. These ionic liquids were dried at 343 K under a reduced

pressure (∼10 mbar) at least for 4 h prior to use. All ILs used in thisstudy and the corresponding literatures for the preparation methodare tabulated in Table 1 (full names and molecular structures of ILsare provided in Table S-1 in the Supplementary Content, SC).

2.2. Hydrocarbons solubility measurement

Low-pressure, close to atmospheric individual hydrocarbon(HC) solubility measurements at an isothermal condition (313 K)were carried out based on the pressure-decay observation [16]using similar apparatus design as in the literature elsewhere [17].A stainless steel-made solubility test unit consisting of an equilib-rium cell (EC) and a gas reservoir (GR) was constructed and wasplaced in an isothermal oven as illustrated in Fig. 1.

The apparatus was equipped with K-type thermocouples andwith a pressure transducer for measuring pressures close to atmo-spheric pressure (OMEGA Engineering PX309-030AI, 0–207 kPa,

Fig. 1. Schematic diagram of the apparatus for HC solubility test: (EC) equilibriumcell, (GR) gas reservoir, (V1-5) valves, (T1-3) thermocouples, (P1) pressure trans-ducers and (S) magnetic stirrer.

1 ering and Processing 49 (2010) 192–198

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94 J. Palgunadi et al. / Chemical Engine

xpansion employing pVT relation. The solubility test unit waslaced in an isothermal oven (allowed operation temperature from13 to 333 K) and the inside temperature was carefully controlledith a K-type thermocouple with an accuracy of ±0.1 K.

In a typical experiment, a known quantity of an IL (6–9 mL) wasoaded into the EC and degassed under vacuum at 343 K at leastor 2 h prior to each test. At a specified oven temperature (313 K),he valve (V4) connecting EC and GR was closed to separate theseessels. HC was fed from the HC supply cylinder to the GR through2 and V3, and the amount of the gas at equilibrium was calcu-

ated from the pVT relation. HC was then brought into contact withhe absorbent in the EC by opening V4 and the absorbent was thentirred vigorously to facilitate the HC absorption. After the absorp-ion reached equilibrium, V4 was closed and more HC was addedo the GR. By repeating the absorption measurement, a new equi-ibrium was obtained. The amount of dissolved gas was calculatedrom the difference between the initial gas concentration in the GRnd the concentration in the gaseous phase after equilibration.

Gaseous solute concentrations were calculated using the gasirial equation of state truncated after the second term employ-ng the second virial coefficients taken from the compilation byymond et al. [18] (details of data reductions are available in

he SC). Based on the solubility data, Henry’s law coefficient wasbtained from the slope of a linear isotherm of fugacity versus moleraction.

.3. 1H NMR experiments

In a typical experiment, an IL dissolved in CDCl3 was loaded intothick walled NMR tubes (Wilmad-Lab Glass, O.D.: 5 mm equippedith a Teflon valve) and was pressurized with 140 kPa of ethylene

r acetylene after a brief evacuation. In every measurement, equiv-lent concentration of the IL was used. The proton chemical shiftsn ppm scale were recorded at room temperature.

.4. Recycle test of acetylene absorption

Recycle absorption–desorption tests of acetylene inBMIM][Me2PO4] were performed five times in an isothermalven using a 25 mL glass-made equilibrium cell equipped with aas inlet and a magnetic stirrer. The IL in the cell was degassednder vacuum and then was let in contact with a constant pres-ure of acetylene (∼1 atm) while stirred at 313 K until it reachedquilibrium. The absorbed gas was measured from the weightifference of the IL (digital Mettler AJ180 balance, accuracy of0−4 g) before and after the absorption taking into account themount of acetylene at the headspace. A complete regeneration ofhe IL was achieved by removing the dissolved acetylene under aeduced pressure (<1 mbar) for 30 min at the same temperatureondition as for the absorption test.

. Results and discussion

.1. Hydrocarbons solubility

Table 1 tabulates 13 room-temperature ILs used in thisnvestigation and also their important properties. Some ofhem particularly, imidazolium- or pyrrolidinium-based ILs withRHPO3] or [R2PO4] anion are still poorly characterized and receiveittle attention despite their “practically” feasible properties and

acile preparations [16,19]. Solubility evaluations of ethylene andcetylene in ILs were performed at pressures close to atmosphereanging from 10 to 170 kPa and at 313 K based on the pressure-ecay observation. Henry’s law constants representing the gasolubility were obtained from the slope of the linear isotherm of

Fig. 2. Isothermal absorption of ethylene and acetylene in [BMIM][Me2PO4] and in[BMIM][Tf2N] at 313 K.

pressure versus molar fraction. Fig. 2 shows acetylene and ethy-lene isotherms in [BMIM][Me2PO4] and in [BMIM][Tf2N] at 313 Kon a mol fraction (x2) basis (full solubility data on a mol fractionas well as on a volume basis are presented in the SC, Table S-1).The acetylene and ethylene solubilities virtually increase linearlywith the pressure rise for all the ILs tested. Such a solubility behav-ior could be an indication that both gases are physically absorbed,though some weak intermolecular interactions should not be ruledout [25].

The solubilization of a solute in a solvent is controlled by twodifferent thermodynamic factors. The first one is the formation ofcavities within the solvent to accommodate the solute molecules.The second one is the chemical interaction of the inserted solutewith the solvent molecules. The first factor requires energy tobreak the solvent–solvent interaction. Based on regular solutionmodel, this interaction can be described as Hildebrand solubilityparameter or cohesive parameter (ı). Accordingly, if the first ther-modynamic step is dominant in controlling the solubility of gaseoussolute, the solubility parameter for solvent (ı1) can be used to deter-mine the solubility behavior of a gas.

It is interesting to notice the application of modified regularsolution theory (RST) introduced by Camper et al. to demonstratethe importance of the first thermodynamic step in solubilizationand also to predict the solubility behavior of CO2 and some othergases in imidazolium-based ILs at low pressures [26]. Fundamentalaspects and approaches can be found in their recent publications[27–29]. Based on RST, the gas solubility in a physical solvent cansolely be determined from solubility parameters (ıi) as long asHenry’s law is valid, which means solubility at pressure close toatmosphere.

ln[H2.1] = a + b(ı1 − ı2)2 (1)

In Camper’s works, Kapustin-skii equation was used to estimatethe lattice energy density of ILs and later this value was used tocalculate the solubility parameter of ILs (ı1). This approach resultsin a specific relation of solubility parameter (ı1) as a function ofmolar volume of IL (Eq. (2)).

ı1 ∝(

1

V4/3m

)1/2

(2)

The combination of Eqs. (1) and (2) gives Eqs. (3) and (4) as the

simplified model of RST where a, b, and b* are constants that dependon the temperature and the type of gas.

ln[H2.1] = a + b(ı1)2 (3)

J. Palgunadi et al. / Chemical Engineering and Processing 49 (2010) 192–198 195

Fig. 3. Correlation of hydrocarbon solubility (in terms of natural log Henry’slaw constant) and inverse molar volume of ILs as suggested from Eq. (2). (�)[([(

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DMIM][MeHPO3], (♦) [BMPyrr][BuHPO3], (�) [DMIM][Me2PO4], (�) [BMPyrr][OAc],�) [BMIM][OAc], (�) [BMIM][Me2PO4], (�) [BMIM][MeSO4], (©) [BMIM][Tf2N], (�)BMIM][BuHPO3], (*) [BMPyrr][Tf2N], (+) [BMIM][TFA], (�) [BMIM][MeHPO3], and×) [BMIM][BF4].

n[H2.1] = a +(

b∗

V4/3m

)(4)

rom Eq. (4), a linear relationship must exist for RST to be validver the gas–IL-temperature combination. This relationship alsondicates that a solvent with a greater molar volume possesses amaller solubility parameter and hence a smaller cohesion energyeading to a higher solubility of solute molecules.

Table 2 lists the estimated values of the molar volume ofLs at room temperature, the solubility of acetylene, Henry’s lawonstants of ethylene and acetylene, and the selectivity of acety-ene. Data correlations of natural log Henry’s law constants versusnverse molar volumes for acetylene and ethylene in various ILsre given in Fig. 3. It is obvious that a relatively linear trend forthylene as suggested by Eq. (2) is existed implying that the solu-ility of ethylene is largely controlled by the first thermodynamicactor. In fact, Scovazzo et al. also demonstrated a similar sol-bility behavior for ethylene in imidazolium-, ammonium-, andhosphonium-based ILs and our experimental results with differ-nt types of anion reinforced that finding [30]. In contrast, suchlinear trend is not established for acetylene as shown by the

cattered data points in Fig. 3. It might indicate that the secondhermodynamic step corresponding to a specific molecular inter-ction or a solute–solvent complexation plays more important rolen determining the solubility of acetylene.

It is worthy to mention that one should be aware of the impor-ance of the unit selection as displayed in Table 2 in analyzing theolubility. For example, the mole fraction solubilities of acetylenen [BMIM][Me2PO4] and in [BMPyrr][BuHPO3] are equivalent butased on a volume unit, the acetylene solubility in the former IL isbviously much higher than that in the latter one. Depending on thebjective of the investigation, a molar basis (x) is more important innderstanding the solubility behavior at the molecular scale, whilevolume basis (c/mol L−1) is associated to the bulk solubility and

s relevant to a process application.Experimental and molecular modeling investigations pointed

ut that CO2 preferably interacts with the anion of simple ILs viaewis acid–base interaction [31]. Simple ILs here means that the

ation is not tethered with a functionality such as CO2-phylic group.lthough this interaction is weak and is not always a dominant

actor in determining the degree of CO2 solubility, recent worksndeed indicated that carboxylate-like or fluorinated anions haveome affinity to CO2 [32,33]. It is also worthy to notice that acetate, Ta

ble

2M

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196 J. Palgunadi et al. / Chemical Engineering and Processing 49 (2010) 192–198

Fig. 4. Plot of hydrocarbon solubility (in terms of inverse natural log Henry’s lawconstant) versus basic strength of ILs (in terms of inverse basicity constant). (�)[([(

osaclwcmcttp(saaiaposuihstciiamvsoot

3

aaes

Fig. 5. Plot of acetylene selectivity over ethylene (˛) versus solubility of acety-

DMIM][MeHPO3], (♦) [BMPyrr][BuHPO3], (�) [DMIM][Me2PO4], (�) [BMPyrr][OAc],�) [BMIM][OAc], (�) [BMIM][Me2PO4], (�) [BMIM][MeSO4], (©) [BMIM][Tf2N], (�)BMIM][BuHPO3], (*) [BMPyrr][Tf2N], (+) [BMIM][TFA], (�) [BMIM][MeHPO3], and×) [BMIM][BF4].

ther homologs of carboxylate and dicyanamide [DCA]− are clas-ified as basic anions and the anions such as [Tf2N]− and [BF4]−

re considered neutral [34]. Owing to the weakly (Brønsted) acidicharacter of hydrogen atom in acetylene (pKa = 25), it seems quiteogical to assume that this hydrogen atom will specifically interact

ith a basic anion of ILs. Interaction of acetylene with the cationan be ruled out since there is no specific site on the simple cationoiety for such association. As the anion basicity is assumed to also

ontrol degree of the acetylene solubility, pKbaq value (Table 2) of

he acid corresponding to the anion of ILs might be useful to expresshe relative basicity of the anion. It is expected that the lower theKb

aq of the acid, the stronger the basicity of the respective anionacid conjugate species). To show the basicity–solubility relation-hip, plots of inverse natural log Henry’s law constants of acetylenes well as of ethylene versus inverse pKb

aq values are constructednd displayed in Fig. 4. Although a regular trend does not appear,ndeed, these plots are still able to demonstrate the effect of thenion basicity in controlling the solubilization of acetylene com-ared with that of ethylene. Nevertheless, this time, the solubilityf acetylene is from the sum contributions of two thermodynamicteps as already mentioned previously. For example, acetylene sol-bilities in [BMIM][OAc] and in [BMPyrr][OAc] are very high and are

n a good agreement with the basicity of the anion but the latter oneas a slightly higher capacity due to its greater molar volume (ormaller cohesion energy). Similarly, this explanation is also appliedo [BMIM][BF4] and [BMIM][TFA] and the acetylene solubility is alsoontrolled by the moderate basicity of their anions. For a compar-son, relatively lower acetylene solubilities in [BMIM][Tf2N] andn [BMPyrr][Tf2N] can be associated with the low basicity of thenion. Smaller solvent–solvent interaction energies (the first ther-odynamic step of solubilization) as indicated by the greater molar

olumes of those two latter ILs, now, contribute dominantly to theolubility of acetylene. On the contrary, a linear trend of the plotf ethylene parallels with the x-axis as displayed in Fig. 4 obvi-usly indicating that the solubility of ethylene is independent onhe basicity of the anion.

.2. Hydrocarbons selectivity

So far, it is shown that practically all ILs under study exhibit highcetylene capacity. In particular, an IL with a greater molar volumend a more basic anion results in higher acetylene solubility. How-ver, for the practical application in industrial field, the absorptionelectivity might be the more important interest. To provide a clear

lene (c). (�) [DMIM][MeHPO3], (♦) [BMPyrr][BuHPO3], (�) [DMIM][Me2PO4], (�)[BMPyrr][OAc], (�) [BMIM][OAc], (�) [BMIM][Me2PO4], (�) [BMIM][MeSO4], (©)[BMIM][Tf2N], (�) [BMIM][BuHPO3], (*) [BMPyrr][Tf2N], (+) [BMIM][TFA], (�)[BMIM][MeHPO3], and (×) [BMIM][BF4].

guidance in the appropriate selection of IL for the process design ofthe ethylene/acetylene separation, ideal absorption selectivity (˛)values of acetylene over ethylene (given in Table 2 as the inverseratio of the Henry’s law constants) and solubility magnitudes (c)of acetylene in mol L−1 were plotted in Fig. 5. It can be seen thatthe highest acetylene solubility was found in [BMIM][Me2PO4] butthe selectivity of acetylene/ethylene in this IL is not that high. Toattain a higher selectivity for separations dominated by gas solubil-ity, a smaller molar volume of ILs is required [12,28]. This key pointis demonstrated by [DMIM][MeHPO3] and [DMIM][Me2PO4]. Theformer IL holds only the sixth-best regarding the acetylene capac-ity but possesses the greatest acetylene selectivity among others asthe consequence of sum contributions from the relatively smallermolar volume and the relatively high basicity strength of the anion.Meanwhile, the latter one occupies the second best place in termsof the capacity and the selectivity, again also as the result froma tradeoff between the size and the basicity making these twoILs as the promising candidates for the acetylene/ethylene sepa-ration. To illustrate the selectivity–size correlation concept, dataplot of C2H2/C2H4 solubility selectivity (˛) versus molar volume ofILs is shown in Fig. S-1 (SC) where it shows the better selectivity ofacetylene in [DMIM][MeHPO3] or in [DMIM][Me2PO4] than thosein other ILs.

3.3. 1H NMR experiments

Recalling the acidic character of hydrogen in acetylene, protonNMR spectroscopy could be a suitable method to demonstrate thepresence of a specific interaction between the hydrogen atom(s)of acetylene and the anion of ILs. The 1H NMR spectra of ethy-lene and acetylene dissolved in an arbitrarily selected basic IL,[DMIM][Me2PO4] and in a neutral IL, [BMIM][Tf2N] were presentedin Fig. S-2 (SC). As shown semi-quantitatively in the figure, the peakof proton(s) in free acetylene (1.92 ppm) is shifted to a more down-field region (2.15 ppm) and is downfield-shifted to a lesser extent(1.96 ppm) when this solute was in contact with [DMIM][Me2PO4]and with [BMIM][Tf2N], respectively. This downfield-shifting indi-cates the presence of a complex formation and the different shiftvalues reflect the different complexation strength between acety-lene and either two ILs. On the contrary, proton shift in free

ethylene (5.41 ppm) remains almost undisturbed (5.39 ppm in[DMIM][Me2PO4] as well as in [BMIM][Tf2N]). This result is notunexpected because the acidity of hydrogen on ethylene with apKa = 44 is much weaker than that on acetylene.

J. Palgunadi et al. / Chemical Engineering

Fa

3

wlasraa

4

oitotbe

wmacmuIsiamasiiCtm(tmaq

s

aq

ig. 6. Recycled solubility tests of acetylene in [BMIM][Me2PO4] under 0.1 MPa ofcetylene at 313 K.

.4. Recycle test of acetylene absorption

[BMIM][Me2PO4] which has the highest capacity for acetyleneas chosen for the recycle test. Fig. 6 shows that acety-

ene uptake in this IL remained unchanged after five times ofbsorption–desorption cycles. This recycle experiment demon-trated that the acetylene–IL binding is relatively weak andeversible. Regeneration without loss of IL was achieved completelyfter it was subject to a physical treatment such as evacuation underreduced pressure at a moderate temperature.

. Conclusions and outlook

The solubilization of ethylene in ILs was found to be dependentn the extent of solvent–solvent cohesion energy where the solubil-ty behavior can be explained by solubility parameter. Meanwhile,he acid–base interaction between the acidic hydrogen atom(s)n acetylene and the basic anion of ILs has a dominant role inhe degree of acetylene solubility. These two different solubilityehaviors result in the high absorption selectivity of acetylene overthylene in basic ILs.

A room-temperature IL absorbing a large amount of acetylenehile maintaining a very low solubilization of ethylene will be theost preferable solvent for the effective ethylene/acetylene sep-

ration. It might be achieved by combining two moieties, a smallation size and an acetylene-phylic anion such as basic anion toinimize the solubility of ethylene (molar volume-dependent sol-

bility) and to maximize the solubility of acetylene in the desiredL, respectively. However, synthesis of small-size ILs with suchpecifically physicochemical properties is not an easy task ands somewhat limited by the availability of the current knownnions. Fortunately, few anions as already mentioned here likeethylphosphate ([Me2PO4]−), methylphosphonate ([MeHPO3]−),

cetate ([OAc]−), and potentially dicyanamide ([DCA]−) meet theketched design when they are combined with small dialkylim-dazolium or dialkypyrrolidinium cations. One may also benterested in 1-ethyl-3-methylimidazolium acetate ([EMIM][OAc];AS 143314-17-4) for the selective absorption of acetylene becausehis room-temperature IL possesses reasonably high basicity, small

olar volume (Vm/cm3 mol−1 = 154.37) as well as low viscosity�/Pa s = 0.093 at room temperature) [35] which apparently meethe requirements. Alternatively, the cation of a neutral IL can be

odified by tethering a functionality to improve the affinity towardcetylene though the selectivity may be sacrificed as the conse-uence from the increase of the IL size.

From the molecular point of view, it is also exciting totudy the effect of various cation structures in controlling the

and Processing 49 (2010) 192–198 197

anion–acetylene interaction. Some studies indicated that the acidicproton at the C-2 position of the 1,3-disubtituted imidazoliumring associates with the anion’s moiety [36]. Intuitively, one mayexpect that due to this association, the affinity of the anion inimidazolium-based IL toward acetylene is less than that of theanion in ILs without acidic site such as pyrrolidinium or quaternaryammonium. A much lower absorption ability of acidic ILs (such asBrønsted acidic or Lewis acidic ILs) can also be expected becausethese ILs lack basic site for the acetylene complexation. Acety-lene and ethylene solubility measurements in various cation–anioncombinations as well as thermodynamic analysis are underwaynow. It is also of interest to investigate the utilization of ILs ascarriers in the supported liquid membrane system for the acety-lene/ethylene separation.

It is known that ILs have negligible vapor pressure making thesematerials superior than organic solvents in terms of reusability. Itwas demonstrated that ILs can reversibly and preferably absorbacetylene. Now, for the successful application in the industrialprocesses, other important requisites such as low viscosity, highthermal stability, low-production cost, and high purity also needto put into consideration. Particularly, the direct use of ILs for theextractive media is hampered by the relatively high viscosity. Inthe green chemistry context, perhaps the atom economy and thelow toxicity of ILs may also be included. ILs of [RMIM][X] (X = OAc,BF4, DCA, TFA, and Tf2N) and many other ILs available to dateare still widely prepared by the metathesis reaction from theirhalide salt with the alkali or silver salt of the desired anion. Thissynthesis technique though workable is probably undesirable in alarge-scale production process due to two reaction steps, use of pre-cious silver salts, generation of an equivalent mole of alkali/silverhalide waste, multi-steps purification, excessive use of solvent forwashing, and contamination of alkali/silver halide. The use of previ-ously prepared carbonate-based ionic liquids (CBILS©) introducedby proionic provides a more elegant synthetic method to producevarious kinds of ILs without using alkali salts.

Although more thorough investigations are definitely requiredbefore selecting a suitable IL, based on the physical properties ofILs and the present solubility/selectivity study so far, it seems that[DMIM][MeHPO3] and [DMIM][Me2PO4] mostly fulfill the criteriaabove. Moreover, these cost effective, hydrolytically stable, less vis-cous and halide-free ILs are easily prepared from versatile startingmaterials via one-pot and solvent-free reaction route.

Acknowledgements

We acknowledge the financial support by a grant (AC3-101 and2009K000694) from Carbon Dioxide Reduction & SequestrationResearch Center, one of the 21st Century Frontier Programs fundedby the Ministry of Science and Technology of Korean governmentand Kolon Industries Co.

Appendix A. Nomenclature and units

MW molecular weight (g mol−1)T absolute temperature (K)Tdec decomposition temperature (K)V volume (cm3)Vm molar volume (cm3 mol−1)pKb

aq basicity constant

pKa acidity constantp hydrocarbon pressure (kPa or MPa)H2.1 Henry’s law constant (atm)x mol fraction solubilityc molar solubility (mol L−1)

1 ering

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[34] D.R. MacFarlane, J.M. Pringle, K.M. Johansson, S.A. Forsyth, M. forsyth, Lewis

98 J. Palgunadi et al. / Chemical Engine

reek symbolsdensity (g cm−3)viscosity (Pa s)ideal absorption selectivityHildebrand solubility parameter

ppendix B. Supplementary data

Supplementary data associated with this article can be found, inhe online version, at doi:10.1016/j.cep.2009.12.009.

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