Effects of halide anions to absorb SO2 in ionic liquids

Article abstract of DOI:10.5012/bkcs.2010.31.7.1937

Ionic liquids with halide anions were prepared and the dependency of halide anions on the SO₂ solubility of ILs was investigated. The study shows that the SO₂ solubility of ionic liquids lies in the range 1.91 ～ 2.22 SO₂/ILs mol ratio. SO₂ solubility in ionic liquids with varying halide anions follows the order Br > Cl > I. Theoretical investigation was also conducted at the B3LYP level using the Gaussian 03 program. From the theoretical consideration of the interaction between SO₂ and [EMIm]X (where X = Cl, Br, and I), it has been proposed that primary interaction of halide occurs with C₂-H of the imidazolium and S of SO₂. Experimental results further shows that the absorption and desorption process of SO₂ in ILs was reversible by the three cycles of the absorption at 50 °C and desorption at 140 °C. The reversibility of SO₂ absorption was confirmed by FT-IR studies.

Full text of DOI:10.5012/bkcs.2010.31.7.1937

SO
2
Absorption in Ionic Liquids
Bull. Korean Chem. Soc. 2010, Vol. 31, No. 7 1937
DOI 10.5012/bkcs.2010.31.7.1937
Effects of Halide Anions to Absorb SO in Ionic Liquids
2
Ki Young Lee, Chang Soo Kim, Honggon Kim, Minserk Cheong, Deb Kumar Mukherjee,†,* and Kwang-Deog Jung*
†
Clean Energy Research Center, Korea Institute of Science and Technology, Seoul 136-791, Korea
*
E-mail: jkdcat@kist.re.kr
†
Department of Chemistry and Research Institute of Basic Sciences, Kyung Hee University, Seoul 130-701, Korea
Received January 19, 2010, Accepted May 10, 2010
Ionic liquids with halide anions were prepared and the dependency of halide anions on the SO
investigated. The study shows that the SO solubility of ionic liquids lies in the range 1.91 ~ 2.22 SO
SO solubility in ionic liquids with varying halide anions follows the order Br > Cl > I. Theoretical investigation was
also conducted at the B3LYPlevel using the Gaussian 03 program. From the theoretical consideration of the interac-
tion between SO and [EMIm]X (where X = Cl, Br, and I), it has been proposed that primary interaction of halide occurs
with C -H of the imidazolium and S of SO
cess of SO
of SO absorption was confirmed by FT-IR studies.
2
solubility of ILs was
2
2
/ILs mol ratio.
2
2
2
2
. Experimental results further shows that the absorption and desorption pro-
o
o
2
in ILs was reversible by the three cycles of the absorption at 50 C and desorption at 140 C. The reversibility
2
Key Words: Ionic liquids, SO
2
absorption, SO
2
solubility, Anionic effect
1
1
Introduction
in ILs, particularly for polar gases such as SO
2
.
In this report, ILs with halide anions were prepared and the
The absorption processes have been used to remove and reco-
ver SO in flue gases from the combustion of sulfur-containing
fossil fuels. The conventional SO
dependency of halide anions on the SO solubility of ILs was
2
2
investigated. High solubility of IL with Cl anion in our earlier
1
3
2
removal processes such as
reported work prompted us to investigate further the role of
lime and magnesia scrubbing result in lots of waste water, which
is not consistent with the sustainable principles. The solvent
absorption processes can be a promising technology by minimiz-
anions in IL-SO
ween SO
with the experimental SO
2
absorption system. The interaction energy bet-
2
and ILs with halide ions was calculated and correlated
solubility.
2
ing the additional pollutants by simple recovery of SO by pres-
2
sure swing or temperature swing. The selection of a good solvent
Experimental
is of critical importance for the SO
2
recovery in the absorption
2
process.
Halide anion effects of ILs were investigated using eight IL
samples: [BMIm]Cl (1-butyl-3-methylimidazolium chloride;
Aldrich 97%), [BMIm]Br (Aldrich, 97%), [BMIm]I (prepared),
[EMIm]Cl (1-ethyl-3-methylimidazolium chloride; c-tri Korea,
98%), [HMIm]Cl (1-hexyl-3-methylimidazolium chloride; c-tri
Korea, 98%), [OMIm]Cl (1-octyl-3-methylimidazolium chlo-
ride; c-tri Korea, 98%), [BMpyri]Cl (1-butyl-4-methylpyridini-
um chloride; prepared, 98%) and [BMpyrr]Cl (1-butyl-4-methyl-
pyrrolidinium chloride; prepared, 98%).
SO
2
absorption processes can also be applied to the SO
2
reco-
very in SO
2
-O streamfroma sulfuric acid decomposition unit in
2
a SI (sulfur-iodine) cycle to produce hydrogen using VHTR (very
high temperature reactor). Especially, the volatile organic sol-
vents are not proper for the SO recovery in the SI cycle because
2
3
the SI cycle is a closed-loop process.
Ionic liquids (ILs) have been considered as the absorbents
of SO , owing to their unique properties such as low volatility,
2
4-5
non-flammability and chemical stability. Especially, the non-
volatility of ILs makes them good absorbents to separate SO
in the SO -O stream from the sulfuric acid decomposition in
the SI cycle as well as to remove SO
It has already been reported that SO
1,1,3,3-tetramethylguanidine][lactate] through both physical
For the preparation of [BMIm]I, [BMIm]Cl (0.05 mol) was
dissolved in acetonitrile. NaI (0.05 mol) was added in the
[BMIm]Cl solution, which was stirred at room temperature for
48 h. Then, the solution was filtered to remove KCl, and dried
2
2
2
6
-8
2
from flue gases.
2
could be absorbed in
with MgSO . After evaporating acetonitrile, the resulting sample
4
[
was dried in a vacuum evaporator. The yield of [BMIm]I was
about 90%.
9
and chemical interactions, while it was demonstrated by IR and
NMR analyses that SO was reversibly absorbed in [1,1,3,3-
tetramethylguanidine][BF
was proposed that the cations of ILs were mostly responsible
for the SO absorption. Later, the investigation of the anion and
cation effects on SO absorption of ILs in a SO -O stream by our
group demonstrated that the SO solubility in ILs was mostly
solubility investigation
2
For the preparation of [BMpyrr]Cl, 1-methylpyrrolidine
(0.05 mol) was dissolved in toluene. Chlorobutane was added
in the 1-methylpyrrolidine solution, which was stirred at room
temperature for 10 h. After the purification, the resulting sample
was dried in a vacuum evaporator. The yield of [BMpyrr]Cl
was about 83%. [BMpyri]Cl was prepared following the same
procedure as [BMpyrr]Cl.
10
4
] through physical absorption. It
2
2
2
2
2
3
dependent on the anions in ILs. The SO
2
in ILs from first-principle calculation also suggested that the IL
anion-gasinteraction play a key role in deciding the gas solubility
The absorption and desorption measurements were conducted
by the following experimental apparatus and procedures. An

1
938 Bull. Korean Chem. Soc. 2010, Vol. 31, No. 7
Ki Young Lee et al.
accurate quantity of ionic liquids was loaded in the glass tube
which was set on an analytical balance, and its temperature was
cribed in experiments after neglecting the weight loss by the
pre-absorbed water in the ILs. The T values of the ILs are not
d
controlled by an electric heater. Then, SO
2
gas was introduced
so much different. It can be said that the ILs are stable at the
o
into the glass tube containing the ionic liquid at a constant flow
rate (20 cc/min) and the weight changes were monitored by
means of the analytical balance (Ohaus explore-4200) with an
accuracy of ± 0.01 g. It was confirmed that there was no weight
temperature lower than 200 C.
o
Table 1 also shows SO
2
solubility of the IL samples at 50 C.
The solubility data were taken after the weight remained un-
changed for 20 min. [BMIm]Cl and [EMIm]Cl showed similar
solubility in terms of the mole ratio. [HMIm]Cl showed the solu-
bility slightly higher than [BMIm] and [EMIm] because of a
more free volume within the molecular packing as similarly
found in the imidazolium-based homologs. It should be noted
that [HMIm]Cl, [OMIm]Cl, [BMPyri]Cl and [BMPyrr]Cl had
similar solubility curves, confirming that the cationic effects
variation in the blank test tube experiments where SO gas was
2
allowed to flow at the same rate.
The absorption spectra of ILs before and after gas flow were
measured by an IR spectroscopy (Perkin-Elmer 2000). The ab-
sorption spectra of SO -absorbed ionic liquids were therefore
2
compared with those of pure ionic liquids.
The thermal stability of the ILs was examined by TGA(Ther-
on the SO
2
solubility were not so important. It can also be said
solubility is not so important,
mo Gravimetric Analysis). The decomposition temperature (T
d
)
that halide effect on the SO
2
was expressed as the temperature at which the 5% weight loss
although the solubility differences were clear in the order of
Br > Cl > I (The errors in the solubility were within ± 0.02)
which cannot be explained by difference in free volume.
Figure 1 shows the absorption and desorption behaviors of
of the initial ILs was observed by TGA. The meaning of T can
d
be best described as the starting temperature of the ILs decom-
position stage.
The interactions between [EMIm]X (where X = Cl, Br, and I)
SO
the SO
o
2
in [BMIm]Cl. Three cycles were done without the loss of
and SO
2
were theoretically investigated using a Gaussian 03 pro-
2
absorption capability. The SO
2
absorption was conduct-
1
2
gram. The geometry optimizations and thermodynamic correc-
tions without including the bulk solvent effect were performed
with hybrid Becke 3-Lee-Yang-Parr (B3LYP) exchange-correla-
tion functional with the 6-31 + G* basis sets for C, H, N and O,
and LanL2DZ(ECP) basis sets for S, Cl, Br and I. In order to
obtain the most stable geometries, all kinds of possible interac-
tion patterns were optimized. No restrictions on symmetries
were imposed on the initial structures. All stationary points were
verified as minima by full calculation of the Hessian and a har-
monic frequency analysis.
ed at 50 C and atmospheric pressure (PSO = 1 atm). The solu-
2
bility of SO
2
in [BMIm]Cl was 2.06 SO -mol/ILs-mol. The de-
2
sorption rate was so slow that the total desorption could be
o
achieved by purging with N (20 mL/min) at 140 C, indicating
2
2.5
2.0
1.5
1.0
Results and Discussion
Table 1 shows T of the tested ILs from TGA experiments.
d
0
.5
.0
o
The weight of ILs decreased at about 100 C during TGA, which
can be due to the absorbed water loss in ILs. Especially, [BMIm]
Cl, [HMIm]Cl, [OMIm]Cl and [BMpyrr]Cl showed high weight
0
0
200 400 600 800 1000 1200 1400 1600 1800 2000 2200 2400
o
loss at about 100 C, indicating that those ILs absorbed much
Time (min)
water. Therefore, T values of those ILs were determined as des-
d
Figure 1. SO
absorption in SO
stream of 10 mL/min at 140 C
2
absorption and desorption cycles in [BMIm]Cl; SO
2
o
2
stream of 20 mL/min at 50 C and SO desorption in
2
o
N
2
Table 1. Thermal stability and SO solubility of ionic liquids by TGA
2
analysis
solubilityb
gasous SO2
[BMIm]Cl
SO
2
2
o
a
Ionic Liquids
T / C
d
SO
/ILs (mol/mol)
[
BMIm]Cl
BMIm]Br
BMIm]I
EMIm]Cl
237
252
240
233
225
224
218
213
2.06
2.11
1.91
2.03
2.19
2.19
2.22
2.22
[
[
BMIm]Cl+SO
2
[
[
2
SO Desorped
[
[
HMIm]Cl
OMIm]Cl
[
BMPyri]Cl
3
900 3600 3300 3000 1600 1400 1200 1000
800
600
[
BMPyrr]Cl
‒1
Wavenumber (cm )
a
T is the temperature at which decomposition begins and corresponds
d
b
to the 5% weight loss calculated from dynamic TGA in air. SO
2
solu-
Figure 2. FT-IR spectra of SO
2
gas, pure [BMIm]Cl, SO -[BMIm]Cl,
2
o
bility was measured at 50 C and atmospheric pressure
and SO desorbed [BMIm]Cl.
2

SO
2
Absorption in Ionic Liquids
Bull. Korean Chem. Soc. 2010, Vol. 31, No. 7 1939
(
(
(
a)
d)
g)
(b)
(e)
(h)
(c)
2
.69
Cl
2.06
2.84
2.98
3.10
2
.28
3
.19
Br
I
(f)
2.20
2
.22
2
.24
1.93
1.99
1
.96
2
.67
2.85
S
3.06
Cl
S
S
Br
I
(i)
2
.73
2
.30
2.51
2
.24
2
.30
2
.29
2
.71
2
.89
3.11
Figure 3. Optimized structures showing (a) [EMIM]Cl, (b) [EMIM]Br, (c) [EMIM]I and their interaction with SO
2
in mode I : (d) [EMIM]Cl +
= ‒11.1 kcal/mol) and in mode II : (g) [EMIM]
= ‒11.5 kcal/mol), (i) [EMIM]I + SO (∆G = ‒10.5 kcal/mol).
SO
2
(∆G
i
= ‒11.1 kcal/mol), (e) [EMIM]Br + SO
2
(∆G
i
= ‒10.7 kcal/mol), (f) [EMIM]I + SO
2
(∆G
i
Cl + SO
2
(∆G = ‒11.3 kcal/mol), (h) [EMIM]Br + SO
i
2
(∆G
i
2
i
‒
1 15
that the absorption and desorption process of SO
2
in [BMIm]
1138 cm . This shift to lower wavenumbers was explained
as a consequence of electrostatic and charge-transfer (CT) effects
Cl was reversible. However, high desorption temperature was
‒
16
required for the cycle and the desorption rate was exceedingly
between Br and SO
2
. On the other hand, no special changes
o
slow even at 140 C. It cannot be excluded that SO
2
is partially
in the bands of [BMIm]Cl were observed from the SO absorp-
2
o
chemisorbed on ILs from the absorption-desorption experiment.
The reversibility of SO absorption in [BMIm]Cl was also moni-
tored by FT-IR spectroscopy. Figure 2 shows the spectra of pure
tion. When, SO
2
was desorbed at 140 C, the band at about
‒1
2
3500 cm disappeared, indicating that water was removed with
SO during the desorption process. FT-IR studies clearly shows
2
[
BMIm]Cl, SO
2
absorbed [BMIm], and SO
2
desorbed [BMIm].
that the absorption and desorption process of SO
in [BMIm]Cl.
2
is reversible
‒1
The broad band at about 3500 cm indicates the stretching
peak of hydroxyl from water. Water was difficult to be removed
To explain the solubility behavior of SO in various [EMIm]
2
from [BMIm]Cl. When SO
2
was absorbed in [BMIm]Cl, The
X (where X = Cl, Br, and I), theoretical investigation was con-
ducted at the B3LYP level of the theory (6-31 + G* for C, H, N
and O and LanL2DZ(ECP) for S, Cl, Br and I) using the Gaussian
03 program. The optimized structures showing the interactions
‒1
band at about 3500 cm almost disappeared and new bands at
‒1
1
137 and 1308 cm appeared. The stretching bands of S=O for
‒1 13
SO
2
gas were observed at 1151 (ν
dissolved in hexane at 1148 (ν
1 14
1
) and 1362 (ν
3
) cm , while
those for SO
2
1
) and 1338 (ν
3
)
of [EMIm]X with SO are depicted in Figure 3. As can be de-
2
‒
cm . The shift of the asymmetric band indicates that the bond-
ing strength of S=O decreased by the interaction between SO
and [BMIm]Cl, when SO was absorbed in [BMIm]Cl. In Ra-
man spectra of [BMIm]Br-SO , it was observed that the sym-
metric stretching of SO (ν ) exhibited a very strong band at
duced from the experimental results, there exists a substantial
‒
2
interaction between the basic halide anion, X and acidic center
2
of SO
SO
where the oxygen is interacting with the imidazolium C
2
. There exist two modes of interaction. In the mode I,
2
2
is trapped between imidazolium cation and halide anion
-H
2
1
2

1
940 Bull. Korean Chem. Soc. 2010, Vol. 31, No. 7
Ki Young Lee et al.
‒
‒
10.6
10.7
bility lies in the range 1.91 ~ 2.22 SO
2
/ILs mol ratio. However,
it can be concluded that ILs with halide anions follow the SO
2
[
[ BB MM II mm ]] BB rr
solubility order Br > Cl > I. From theoretical points of view,
two models were examined. In the mode I, SO is trapped bet-
ween imidazolium cation and halide anion where the oxygen is
interacting with the imidazolium C -H and methyl part and
sulfur is interacting with the halide. In the mode II, primary
interaction of halide is with the imidazolium C -H and SO is
interacting with halide anion. The free energies of interaction
∆G ) with SO in mode I were calculated as ‒11.1, ‒10.7 and
11.1 kcal/mol for [EMIm]Cl, [EMIm]Br and [EMIm]I, res-
2
‒
‒
‒
‒
‒
10.8
10.9
11.0
11.1
11.2
2
2
2
[
B[ B MM I Imm ] C] Cll
(
i
2
[
BMIm]I
[
BMIm]I
‒
pectively in the order Cl ~ I > Br. Energy calculations in mode
II were ‒11.3, ‒11.5 and ‒10.5 kcal/mol for [EMIm]Cl, [EMIm]
Br and [EMIm]I, respectively in the order Br > Cl > I. Therefore,
it has been concluded that model II was more preferable to
model I which is consistent with the experimental results.
2
.00
2.05
2.10
2.15
2.20
2.25
SO
Figure 4. The relationship between the interaction energy of halide
anion-SO and SO solubility.
2
solubility (SO
2
gmol/IL gmol)
2
2
Acknowledgments. The authors wish to acknowledge the
financial support from MOST of Korea and KAERI. This work
has been carried out under the Nuclear Hydrogen Development
and Demonstration Project (NHDD).
and methyl part and sulfur is interacting with the halide. In the
mode II, primary interaction of halide is with the imidazolium
C
2
-H and SO
2
is interacting with the halide anion. The free ener-
gies of interaction (∆G
i
) with SO in mode I were calculated
2
as ‒11.1, ‒10.7 and ‒11.1 kcal/mol for [EMIm]Cl, [EMIm]Br
and [EMIm]I, respectively (see Figure 3-d, e and f) in the
order of Cl ~ I > Br. Similar energies in mode II were calculated
as ‒11.3, ‒11.5 and ‒10.5 kcal/mol for [EMIm]Cl, [EMIm]Br
and [EMIm]I, respectively (see Figure 3-g, h and i) in the order
of Br > Cl > I. The change in the magnitude of interaction de-
pending on the mode of interaction can be presumably ascribed
to the strength of interaction between the halide anion and the
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