Efficient and reversible absorption of ammonia by cobalt ionic liquids through Lewis acid-base and cooperative hydrogen bond interactions

Article abstract of DOI:10.1039/c8gc00215k

Ammonia (NH₃) emissions have caused a wide range of environmental problems and serious harm to human health. However, efficiently separating NH₃ and simultaneously recovering high purity NH₃ easily remains a great challeng

Full text of DOI:10.1039/c8gc00215k

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Efficient and Reversible Absorption of Ammonia by Cobalt
Ionic Liquids through Lewis Acid-Base and Cooperative
Hydrogen Bond Interactions
Shaojuan Zeng a, Lei Liu b, Dawei Shang a, c, Jiangpeng Feng a,c, Haifeng Donga, Qiuxia Xu a,
Xiangping Zhang a,c,∗, Suojiang Zhang a
a
Key Laboratory of Green Process and Engineering, State Key Laboratory of Multiphase
Complex System, Beijing Key Laboratory of Ionic Liquids Clean Process, Institute of Process
Engineering, Chinese Academy of Sciences, Beijing, 100190, China
b
Mulliken Center for Theoretical Chemistry, Institute for Physical and Theoretical Chemistry,
University of Bonn Beringstr. 4, 53115 Bonn, Germany
c
College of Chemical and Engineering, University of Chinese Academy of Sciences, Beijing
100049, China
∗Corresponding author. Tel./fax: +86-010-82544875
E-mail address: xpzhang@ipe.ac.cn
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ABSTRACT
Ammonia (NH3) emissions have caused a wide range of environmental problems
and serious harm to human health. However, efficiently separating NH3 and
simultaneously recovering high purity of NH3 easily remains a great challenge. A new
strategy to design transition metal ionic liquids (MILs) by combing specific metal
centers and ligands with ILs was proposed for efficient and reversible absorption of
NH3 . Not only exceptional NH3 absorption capacity and high NH3/CO2 selectivity,
but also excellent recyclability were achieved by the cobalt ILs [Cnmim]2[Co(NCS)4].
The maximal capacity of NH3 is up to 6.09 molNH3·molIL-1 at 30ꢀ and 0.10 MPa,
which is much higher than all reported ILs to date, and is over 30 times higher than
the conventional ILs [Cnmim][SCN]. The superior NH3 capacity and desorption
performance is originated from the moderate Lewis acid-base and cooperative
hydrogen bond interactions between the metal centers-ligands and NH3.
Keywords: Cobalt ionic liquids, Ammonia absorption, Lewis acid-base interaction,
Hydrogen bond
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1. Introduction
Emission of toxic gases and greenhouse gases is a serious problem in the
establishment of a sustainable society and is driving the implementation of green
chemistry concepts. Therefore, gas separation and purification as an everlasting
research topic has already attracted extensive attentions worldwide1, 2. As one of the
typical alkaline and poisonous pollutants, ammonia (NH3) is widely considered to be
a primary factor for the formation of fog and haze, which has been seriously
hazarding environment and human health for a long time3. On the other hand, NH3 is
also an important raw material and applied in various fields, such as fertilizer,
freezing medium4. Therefore, a few technologies for NH3 separation, like water
scrubbing and acid scrubbing, have been applied in industries. However, some
inherent drawbacks come with yet. For example, water scrubbing consumes a large
amount of water to reach ultra-low NH3 emission requirements (lower than 10
mg·m-3), and finally obtains low concentration of aqueous ammonia (less than 3 wt%),
which incurs great difficulties for NH3 desorption and recycle. Acid scrubbing
produces huge amounts of irreversible ammonium salts as low value-added products
and wastewater that doesn’t follow the principles of green chemistry. Hence, the
development of green techniques for efficiently separating and recovering NH3 is of
critical significance.
Ionic liquids (ILs) have caught considerable attentions as competitive alternatives
in SO2, CO2 and NH3 separation owing to their peculiar properties, like extremely low
vapour pressure, high thermal stability and tuneable properties5-12. It is reported that
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NH3 has optimistic solubility in most conventional ILs, and the cation plays a greater
role in NH3 absorption than the anion13-18, which differs from the case for acid gases
19, 20
such as SO2
and CO221-25. Comparing with other cations such as pyrrolodinium,
quinolinium, and phosphonium, imidazolium-based ILs show higher NH3 capacity26.
Furthermore, the introduction of hydroxyl groups into cations can substantially
improve NH3 capacity due to hydrogen bonding between the hydroxyl group and
NH38, 27. For example, NH3 solubilities in 1-2(-hydroxyethyl)-3-methylimidazolium
tetrafluoroborate ([EtOHmim][BF4]) and choline bis(trifluoromethylsulfonyl)imide
([Choline][Tf2N]) were 1.70 and 1.86 molNH3·molIL-1 at 20ꢀ and 0.10 MPa,
respectively, whereas that in [Emim][BF4] was only 0.28 molNH3·molIL-1. Recently,
Shang et al.18 reported that
a
novel protic IL 1-butylimidazolium
bis(trifluoromethylsulfonyl)imide ([Bim][Tf2N]) absorbs 2.69 molNH3·molIL-1 at 40ꢀ
and 0.10 MPa through the strong hydrogen bonding between NH3 and cations. Xing et
al.28 also designed the hybrid deep eutectic solvents with high NH3 capacity of 0.13
gNH3·gabsorbent-1 at 40ꢀ and 0.10 MPa through a flexible hydrogen-bonded
supramolecular network. However, these physisorption ILs are still unable to compete
with the traditional absorbents. For this reason, an alternative approach has been
proposed, that is chemisorption, which is normally accomplished by the incorporation
of metal centers. Metal chlorides as chemical adsorbents show higher capacity of NH3
adsorption than that in water, but metal chloride-NH3 systems have poor heat transfer
performance and attenuation of adsorptive capacity due to their irreversibility29.
Kohler et al.30 prepared novel spherical activated carbon support materials containing
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copper(II) salts, and found that the copper(II)-based materials show better NH3
adsorption performance by combining chemical and physical interaction than that of
the metal-free materials. However, the common drawback for these materials is the
formation of strong chemical interactions between metal centers and gas molecules,
which makes it impossible to regenerate NH3 and reuse the adsorption materials.
Therefore, an appropriate balance and tradeoff between the absorption capacity and
desorption ability must be considered to design novel absorbents for highly efficient
and reversible absorption of NH3.
In this work, we proposed a new strategy to achieve high performance on both
NH3 absorption and desorption using transition metal ILs (MILs) through the Lewis
acid-base and hydrogen bonding interaction between metal centers-ligands and NH3.
The key of desiging MILs is to screen the suitable metal centers and ligands, which
can react with NH3 forming high coordination structures (e.g. octahedron) with
moderate stability constants (Table S1). Besides, the MILs should keep in a liquid
state before and after NH3 absorption under the temperature (≤ 50ꢀ), and have good
flowability, which is favorable for the transport and kinetic aspects of NH3 absorption.
Thus, the cobalt ILs [Cnmim]2[Co(NCS)4] (n=2, 4, 6) were designed and synthesized
for efficient and reversible absorption of NH3. The physicochemical properties of ILs
(density, viscosity, thermal decomposition temperatures and glass transition
temperatures), the effect of temperature, pressure and water contents on NH3
absorption performance as well as NH3 selectivity and the regeneration and
recyclability of ILs were systematically investigated. Moreover, the mechanism of
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NH3 absorption in [Cnmim]2[Co(NCS)4] was well explained through both
experimental characterizations and Quantum Chemical calculations.
2. Experiment
2.1. Materials
NH3 (99.999%), CO2 (99.999%) and N2 (99.99%) were supplied by Beijing
Beiwen Gas Factory. [Emim]Cl, [Bmim]Cl and [Hmim]Cl (> 99.0%) were supplied
by Linzhou Keneng Material Technology Co. Ltd. Cobalt (II) chloride, anhydrous
(CoCl2, 97.0%) was obtained from Sinopharm Chemical Reagent Co., Ltd. Other
chemical reagents of analytical grade, including acetone, dichloromethane, kalium
thiocyanate (KSCN) were purchased from Beijing Chemical Company.
2.2. Synthesis and characterizations of ILs
Two kinds of ILs including three conventional ILs [Cnmim][SCN] and three
MILs [Cnmim]2[Co(NCS)4] (n=2, [Emim], 1-ethyl-3-methylimidazolium; n=4,
[Bmim], 1-butyl-3-methylimidazolium; n=6, [Hmim], 1-hexyl-3-methylimidazolium.
[SCN], thiocyanate, and [Co(NCS)4]2-, tetraisothiocyanatocobaltate(II)) were
synthesized by a one-step method. The conventional ILs [Cnmim][SCN] including
1-ethyl-3-methylimidazolium
1-butyl-3-methylimidazolium
thiocyanate
thiocyanate
([Emim][SCN]),
([Bmim][SCN]),
1-hexyl-3-methylimidazolium thiocyanate ([Hmim][SCN]), were synthesized
following the reported methods5, 6. Taking [Bmim][SCN] as an example, the anion
exchange reaction was performed by mixing [Bmim]Cl and equimolar KSCN in
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acetone. After stirring for 48 h at room temperature, the mixture was filtered and
washed with acetone. The organic layer was collected and concentrated by rotary
evaporation to remove the solvent. Then the raw product with dichloromethane was
frozen for 3~5 h under -5ꢀ. After filtration and removal of solvents, the final product
[Bmim][SCN] was obtained as the dark orange liquid under vacuum at 50ꢀ for 48 h.
[Emim][SCN] and [Hmim][SCN] were synthesized in the same method expect the
[Bmim]Cl was substituted by [Emim]Cl and [Bmim]Cl, respectively.
The cobalt ILs [Cnmim]2[Co(NCS)4] were synthesized by mixing CoCl2 and
KSCN in an acetone solution of [Bmim]Cl through the anion-exchange method. In the
typical experiment, 8.000 g CoCl2 (0.062 mol), 23.951 g KSCN (0.246 mol) and
21.526 g [Bmim]Cl (0.123 mol) were added into 100 ml acetone and stirred 48 h at
room temperature. After filtration and removal of the solvents by rotary evaporation,
the resulting product was washed with dichloromethane, and dried under vacuum at
50ꢀ for 48 h to afford [Bmim]2[Co(NCS)4] as the blue liquid. [Emim]2[Co(NCS)4]
and [Hmim]2[Co(NCS)4] were synthesized in an analogous manner, using [Emim]Cl
and [Bmim]Cl instead of [Bmim]Cl, respectively. Elemental analysis (%) calculated
for [Emim]2[Co(NCS)4]: C 37.42, H 4.32, N 21.82; Found: C 37.69, H 4.29, N 21.74;
For [Bmim]2[Co(NCS)4]: C 42.17, H 5.31, N 19.56; Found: C 42.05, H 5.40, N 19.67;
For [Hmim]2[Co(NCS)4]: C 46.09, H 6.09, N 17.92; Found: C 46.26, H 6.27, N 17.46.
All the ILs were dried under vacuum at 50ꢀ for 48 h before use. The FT-IR
spectra were recorded by a Thermo Nicolet 380 Spectrometer to confirm the
structures of the studied ILs. The structures of three conventional ILs ([Emim][SCN],
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1
[Bmim][SCN] and [Hmim][SCN]) were confirmed by H NMR and 13C NMR
spectroscopy (see the supporting information) with Bruker 600 spectrometer. UV/vis
spectra in the range of 200-900 nm were recorded using UV-Vis spectrophotometer
(UV-2550). Measurements were done on the solutions of [Cnmim]2[Co(NCS)4] in
acetone, and the absorption maximum of [Cnmim]2[Co(NCS)4] in acetone were 623
and 587 nm. The water contents in the ILs after drying were determined with
volumetric Karl Fischer Titration (Metrohm, 787 KF Titrino) and found to be less
than 0.01 wt%. The residual halide contents in the ILs were measured using
PXSJ-226 Series Ion meter (INESA Scientific Instrument Co. Ltd), and found to be
less than 0.02 wt%. The density and viscosity of ILs were measured from 20 to 70ꢀ
with a density meter (Anton Paar DMA 5000) and an automated microviscometer
(Anton Paar AMVn), respectively. Decomposition temperatures were determined by
thermogravimetric analysis (TGA Q5000) by heating from room temperature to 600ꢀ
under nitrogen atmosphere with heating rate of 10 ꢀ·min-1. The glass transition
temperature of ILs was measured by differential scanning calorimeter (Mettler Toledo
DSC1) from -150 to 50ꢀ under nitrogen atmosphere with heating rate of 10 ꢀ·min-1.
2.3. NH3 absorption and desorption
NH3 absorption and desorption experiments followed the same procedures of our
previous work5, 6, 31. In a typical absorption, NH3 gas of atmosphere pressure is
bubbled through about 5 g ILs in the glass container with an inner diameter of 3.0 cm
at a flow rate of 140 ml·min-1. The glass container was partly immersed in a circulated
water bath at the desired temperature within ± 0.10ꢀ. The amount of NH3 absorbed

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in the IL was measured at different times by an electronic balance with an accuracy of
± 0.10 mg, and was finally determined according to the increase of the weight. For
NH3 absorption under reduced pressure, the NH3 partial pressure was adjusted by
changing the flow of NH3 and N2. For NH3 desorption, NH3 was desorbed by heating
and bubbling N2 through IL solutions with saturated NH3. In this study, pure N2 gas
with the same flow rate was bubbled through the IL with absorbed NH3 in the glass
container at 80ꢀ, and the amount of NH3 desorption was also determined at regular
intervals by an electronic balance until the weight kept constant. The reproducibility
of the absorption and desorption experiments was less than ± 3%, and it was
estimated that the experimental error involved in the method was at the level of ± 5%.
2.4. Computational methods
All density functional theory (DFT) calculations were performed by employing
Turbomole 7.0 software32. The structures were optimized at the TPSS level of
theory33, with the BJ-damped variant of the DFT-D3 dispersion correction34 in
conjunction with the def2-TZVP basis set35, 36. Harmonic vibrational frequency
calculations were conducted at the same level to characterize the nature of the
stationary points along the reaction coordinates: no imaginary frequencies were found
for the local minima, and one and only one imaginary frequency was found for the
transition state. The thermostatistical contributions to the Gibbs energy in the gas
phase were obtained from a modified rigid-rotor-harmonic oscillator approximation37
at temperatures of 25ꢀ and 1 atm pressure. The density-fitting RI-J38 approach for
the Coulomb integrals was used to accelerate the geometry optimization and
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frequencies calculations. Accurate electronic energies were obtained from single point
calculations upon the optimized structures at both the hybrid metal PBE0-D3
functional39 with the basis set of def2-TZVPP. The COSMO-RS
(Conductor-like Screening Model for Real Solvents) solvation model40, 41 was used to
compute the solvation Gibbs energies by employing the gas–phase optimized
structures, and with ILs as the solvent. These calculations were done with the
COSMOtherm program42.
3. Results and discussion
3.1. Physicochemical properties of ILs
Density and viscosity of ILs are two fundamental parameters for screening and
evaluation of ILs for gas separation. Therefore, Densities and viscosities of the
conventional ILs [Cnmim][SCN] and the cobalt ILs [Cnmim]2[Co(NCS)4] in the
temperature range from 20 to 70ꢀ were studied as shown in Figure 1. Comparing
with [Cnmim][SCN], [Cnmim]2[Co(NCS)4] with doubly charged anions have much
higher densities and viscosities. For instance, the density and viscosity of
[Bmim][SCN] were 1.08 g·cm-3 and 83.71 mPa·s at 20ꢀ, whereas those of
[Bmim]2[Co(NCS)4] increased to 1.22 g·cm-3 and 480.41 mPa·s, respectively.
However, the viscosities of [Cnmim]2[Co(NCS)4] are much lower than those of other
reported MILs. Most MILs with transition metal halide anions are solid at room
temperature due to the extensive hydrogen-bonding network43. Only a few of MILs,
like [FeCl4]- -based ILs, and phosphonium-based ILs containing different metal halide
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anions at room temperature are liquid, and their viscosities are extremely high, with
the maximum of 123500 mPa·s at 20ꢀ44. The lower viscosities of
[Cnmim]2[Co(NCS)4] in this work may be due to the weaker interactions between soft
pseudo-halide anions and the hard hydrogen atoms of the imidazolium cations than
the halogen groups45.
1.35
(a)
1.30
1.25
1.20
1.15
1.10
1.05
1.00
0.95
20
30
40
50
60
70
Temperature / oC
700
600
500
400
300
200
100
0
(b)
20
30
40
50
60
70
Temperature / oC
Figure 1. Variation in the densities (a) and viscosities (b) of ILs with temperature: ■,
[Emim]2[Co(NCS)4]; ●, [Bmim]2[Co(NCS)4]; ▲, [Hmim]2[Co(NCS)4]; □, [Emim][SCN]; ○,
[Bmim][SCN]; ∆, [Hmim][SCN].
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In addition, the thermal properties including thermal decomposition temperatures
(Td) and glass transition temperatures (Tg) of [Cnmim][SCN] and [Cnmim]2[Co(NCS)4]
were also studied. As shown in Table 1, the presence of anionic transition-metal
thiocynate complexes plays an important role in thermal properties of ILs. The
thermal decomposition temperatures of the cobalt ILs are all above 280ꢀ, and much
higher than those of [Cnmim][SCN], indicating that [Cnmim]2[Co(NCS)4] exhibit
more excellent thermal stability. Moreover, no melting points were detected for
[Cnmim]2[Co(NCS)4], but glass transition temperatures was around -63ꢀ.
Table 1. Thermal decomposition temperatures and glass transition temperatures of ILs
ILs
aTd /ꢀ
302.8
300.0
288.4
230.4
229.3
228.9
Tg /ꢀ
-63.22
-62.91
-62.57
-101.13
-92.70
-91.28
[Emim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Hmim]2[Co(NCS)4]
[Emim][SCN]
[Bmim][SCN]
[Hmim][SCN]
aTd is defined as the temperature at which the IL loses 5% of its initial mass while heating from
room temperature to 600ꢀ with a heating rate of 10ꢀ·min-1 under N2 atmosphere.
3.2. NH3 absorption by cobalt ILs
̄
In order to study the effect of the anionic transition-metal complex [Co(NCS)4]2
with thiocynate ligands on NH3 absorption, NH3 capacity by the cobalt ILs
[Cnmim]2[Co(NCS)4] and the conventional ILs [Cnmim][SCN] as a function of time
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was studied and shown in Figure 2. It was found that all the cobalt ILs
[Cnmim]2[Co(NCS)4] exhibited very high NH3 capacity, which can absorb 5.99~6.09
molNH3·molIL-1 at 30ꢀ and 0.10 MPa. However, the maximal NH3 capacity by the
conventional ILs [Cnmim][SCN] is only 0.20 molNH3·molIL-1 under the same
conditions. This means that NH3 capacity by [Cnmim]2[Co(NCS)4] is over 30 times
than that of in the corresponding metal-free ILs [Cnmim][SCN]. When the cation
varied from [Emim]2 to [Bmim]2 and [Hmim]2, NH3 capacity of the cobalt ILs
increased slightly from 5.99 to 6.03 and 6.09 molNH3·molIL-1, respectively. The
similar results are obtained using [Cnmim][SCN], implying the minor influence of the
alkyl chain length of the cation on NH3 capacity. Therefore, significant improvement
in NH3 absorption by [Cnmim]2[Co(NCS)4] is mostly attributed to the strong
interaction between the anion and NH3, which is different from the mechanism of
NH3 physisorption in ILs8, 16, 18. Meanwhile, the molar ratio between NH3 and the
amount of cobalt(II) ion also indicates the formation of [Co(NH3)6]2+ complexes
through chemical complexation.
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7
6
5
4
3
2
1
[Emim][SCN]-NH3
[Bmim][SCN]-NH3
[Hmim][SCN]-NH3
[Emim]2[Co(NCS)4]-NH3
[Bmim]2[Co(NCS)4]-NH3
[Hmim]2[Co(NCS)4]-NH3
[Bmim]2[Co(NCS)4]-CO2
[Bmim][SCN]-CO2
0.2
0.1
0.0
0
20
40
60
80
100
Time / min
Figure 2. NH3 and CO2 absorption behaviors of [Cnmim]2[Co(NCS)4] and [Cnmim][SCN] at 30ꢀ
and 0.10 MPa
3.3. NH3 selectivity in cobalt ILs
There are always other gases accompanied with NH3 in real industries, such as
low concentration of CO2 in the tail gases of urea or ammonium bicarbonate plants.
Therefore, selectively separating NH3 from gas mixtures becomes particularly
important. Table 2 listed the NH3 and CO2 capacity and NH3/CO2 selectivity in
[Cnmim]2[Co(NCS)4] and [Cnmim][SCN], respectively. The results showed that
[Cnmim]2[Co(NCS)4] exhibit much higher NH3/CO2 selectivity in the range from 39.0
to 52.5, whereas the NH3/CO2 selectivity are lower than 7 for the corresponding
[Cnmim][SCN], indicating that [Cnmim]2[Co(NCS)4] can effectively and selectively
separate NH3 from mixture gas containing CO2.
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Table 2. NH3/CO2 selectivity in [Cnmim]2[Co(NCS)4] and [Cnmim][SCN]
Molgas·molIL-1 a
NH3/CO2
ILs
selectivity
NH3
CO2
[Emim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Hmim]2[Co(NCS)4]
[Emim][SCN]
5.99
6.03
6.09
0.18
0.19
0.20
0.114
0.129
0.156
0.028
0.042
0.058
52.5
46.7
39.0
6.4
[Bmim][SCN]
4.5
[Hmim][SCN]
3.4
aThe absorption experiment was carried out at 30ꢀ and 0.10 MPa.
3.4. Effect of water on NH3 absorption
Besides, water is also inevitably present in industrial gases. Hence, saturated
water contents of ILs by mass fraction at 30ꢀ were measured, and the effect of water
in ILs on NH3 absorption at 30ꢀ and 0.10 MPa was also investigated. It was found
that [Cnmim]2[Co(NCS)4] has low saturated water contents in the range of 0.97 to
1.77 wt%, and become gradually hydrophobic with the increasing alkyl chain length
of the imidazodium cation. The addition of the small amounts of water (less than 2.0
wt%) into [Cnmim]2[Co(NCS)4] has the slight negative impact in NH3 capacity. For
example, when the water content in [Bmim]2[Co(NCS)4] increases to 1.22 wt%, the
NH3 capacity decreases to 0.172 gNH3·gIL-1, which corresponds to 4.44% reduction
in NH3 capacity. As the water content in [Bmim]2[Co(NCS)4] is about 0.66 wt%, the
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NH3 capacity is nearly close to that of the pure IL. However, the presence of water
could reduce the viscosity of the cobalt ILs obviously, which is helpful for improving
the gas-absorption kinetics in these ILs46.
Table 3. Saturated water contents and NH3 capacity in [Cnmim]2[Co(NCS)4] with water
ILs
Water contents
ag NH3·g IL-1
bg NH3·g IL-1
[Emim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Hmim]2[Co(NCS)4]
1.77%c
1.22%c
1.06%
0.66%
0.97%c
0.192
0.172
0.175
0.179
0.155
0.198
0.180
0.166
aCobalt ILs with water. bPure cobalt ILs. cSaturated water contents of cobalt ILs.
3.5. Effect of temperature and pressure on NH3 absorption
Moreover, [Bmim]2[Co(NCS)4] was selected as an example to further investigate
the effect of temperature and pressure on NH3 absorption. As shown in Figure 3, both
the temperature and the pressure play a significant role in NH3 absorption. The mass
ratio of NH3 to IL for [Bmim]2[Co(NCS)4] decreased continuously from 0.18 to 0.063
with an increase of temperature from 30 to 80ꢀ, and increased from 0.052 to 0.18
with the increasing partial pressure of NH3 from 0.01 to 0.10 MPa. When the partial
pressure was 0.005 MPa, the mass ratio of NH3 to [Bmim]2[Co(NCS)4] was 0.018,
which is even higher than that in several conventional ILs at 0.10 MPa14, 16, indicating
that the cobalt ILs is suitable for low concentration of NH3 absorption. The NH3
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capacity in [Bmim]2[Co(NCS)4] increases with the increasing pressure and the
decreasing temperature. Thus, the absorbed NH3 in the cobalt ILs can be easily
released by heating or/and reducing the partial pressure of NH3.
0.20
(a)
0.16
0.12
0.08
0.04
30
40
50
60
70
80
Temperature / oC
0.20
0.16
0.12
0.08
0.04
0.00
(b)
0.00
0.02
0.04
0.06
0.08
0.10
NH3 partial pressure / MPa
Figure 3. Effect of temperature (a) and pressure (b) on NH3 absorption by [Bmim]2[Co(NCS)4]
3.6. Regeneration and recycle of cobalt ILs
In order to examine the regeneration and recyclability of the cobalt ILs
[Cnmim]2[Co(NCS)4], five consecutive absorption-desorption experiments by
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[Bmim]2[Co(NCS)4] were further investigated. As seen in Figure 4, all the NH3
absorbed can be easily released as the IL-NH3 complexes are exposed to reduced
pressure and relatively high temperature. Importantly, no changes have been observed
on the structures of the cobalt ILs after NH3 desorption, and the IR spectra are almost
identical to that of the fresh cobalt ILs (Figure S10). Furthermore, [Bmim]2[Co(NCS)4]
can be recycled and exhibit no significant loss in NH3 capacity after five
complexation-evolution cycles, implying that the process is completely reversible.
Meanwhile, the multiple absorption and desorption of NH3 by [Bmim]2[Co(NCS)4]
with saturated water (1.22 wt%) was also investigated (Figure S9). Similarly, the
absorption performance of NH3 by [Bmim]2[Co(NCS)4] with water after five cycles
also remains stable. However, the absorbed NH3 is more easily and faster desorbed
from [Bmim]2[Co(NCS)4] with water than the pure IL under the same conditions.
0.20
0.16
0.12
0.08
0.04
0.00
0
100
200
300
400
500
600
700
800
900
Time / min
Figure 4. Five consecutive cycles of NH3 absorption and desorption by [Bmim]2[Co(NCS)4]. NH3
absorption was carried out at 30ꢀ and 0.10 MPa, and NH3 desorption at 80ꢀ under N2.
Absorption ■, desorption □
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3.7. The mechanism for NH3 absorption
All the results discussed above demonstrated that the [Cnmim]2[Co(NCS)4] not
only have higher capacity of NH3 absorption, but also exhibit excellent recyclability.
Why do the cobalt ILs [Cnmim]2[Co(NCS)4] exhibit such high capacity and excellent
regeneration performance? In order to explain the NH3 absorption and desorption
mechanism with [Cnmim]2[Co(NCS)4], the optimized structures of NH3 and CO2 with
the SCN− and Co(NCS)42− anions by DFT calculations were depicted in Figure 5. The
shortest distance between the gases (CO2 and NH3) and SCN− are 2.6 Å and 2.7 Å,
respectively (Figure 5(a) and 5(b)). The computed reaction Gibbs free energies ∆G are
-0.1 kcal·mol−1 and 0.8 kcal·mol−1 for NH3 and CO2 absorption, respectively, which
are in the range of typical binding affinity of physisorption. The calculated NH3/CO2
selectivity is 4.6, which is in good agreement with the experimental values (Table 2).
For the Co(NCS)42− anion, six NH3 molecules replace four SCN− anions and
coordinate to the cobalt center, forming a [Co(NH3)6]2+ complex (Figure 5(c)).
Subsequently, the NH3 molecules interact with four SCN− anions by hydrogen bonds
between H atoms of NH3 and N atom of SCN−, finally forming an octahedral
coordinated cobalt complex [Co(NH3)6(NCS)4]2−. The computed bond orders for four
Co-N bonds are about 1.0, and the computed ∆G is −0.6 kcal·mol‒1 (Supporting
Information, Equation S6), indicating the chemical bonding characteristics of the
Co−N bonds. This slightly negative value means that NH3 desorption is possible when
increasing the temperature of the IL-NH3 system, which is consistent with our
experiments, i.e., NH3 starts to be released from the cobalt ILs at 80ꢀ. For
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comparison, we also calculated the ∆G for NH3 adsorption by a neutral compound
Co(SCN)2 (Supporting Information, Equation S7), which shows a lower capacity of
-1
3.83 molNH3·molCo(SCN)2 than the cobalt ILs. The overall reaction is strongly
exothermic with a computed ∆G of −16 kcal·mol−1 per NH3 (Supporting Information,
Figure S8), implying that it is impossible for NH3 desorption from the metal salt
Co(SCN)2. In addition, an optimized structure for [Co(NCS)4(CO2)4]2− was obtained
as shown in Figure 5(d). The shortest distance between CO2 and Co(NCS)42- is about
3.3 Å, and no chemical bonds are found between CO2 and Co(NCS)42- based on WBO
analysis. The cobalt ILs have about 3 to 4 times higher CO2 capacity than that of the
conventional ILs due to four absorption sites (i.e. SCN−) in the cobalt ILs while only
one in the conventional ILs. This is in consistence with the experimental results.
Figure 5. Optimized structures of (a) SCN− with CO2, (b) SCN− with NH3, (c)[Co(NH3)6(NCS)4]2−,
and (d)[Co(NCS)4(CO2)4]2−. The selected distances are given in angstrom, Å. Color legend: S
yellow, N blue, C black, O red and H white
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The experimental and simulated IR spectra were also further compared to
examine NH3 absorption mechanism by the cobalt ILs (Figure 6). As shown in Figure
6(a), four newly appeared bands in the experimental IR spectrum of
[Bmim]2[Co(NCS)4] after NH3 absorption were observed, i.e. bands A1 (~1200 cm−1),
A2 (~1600 cm−1), A3 (~3300 to 3500 cm−1) and B (~2800 cm−1). Among them, the
peaks A1, A2 and A3 are essentially corresponding to vibrational modes of NH3 as
the comparison of the IR spectra of Co(NH3)62+ (Figure 6(c)) with the IR spectra of a
free NH3 molecule (Figure 6(f)) , indicating that chemical absorption of NH3 occurs in
[Bmim]2[Co(NCS)4]. For example, the peaks around 1200 cm−1 (A1) and 1600 cm−1
(A2) are the N−H wagging mode and the H−N−H scissoring mode, respectively. The
third one, A3 (~3300 to 3500 cm−1), is the N−H stretching mode, including symmetric
and asymmetric stretching. For free NH3 molecules, the intensities of both modes are
relative low, i.e. 2.4 and 0.009 km·mol‒1, for symmetric and asymmetric stretching,
respectively. When coordinating with cobalt atom, these two modes become stronger
with computed intensities increasing to 20 (~3380 cm‒1) and 50 (~3467 cm‒1)
km·mol‒1, respectively. The fourth peak, B (~2800 cm‒1), is the C−H stretching mode
of the CH group in the five‒membered ring of Bmim+ cation if we analyze the IR
2‒
spectra of [Bmim][SCN]. After NH3 absorption, SCN− in Co(SCN)4 complex is
replaced by NH3 and interacts with NH3 only though hydrogen bonds (Figure 6(a)).
As a result, SCN− becomes a nearly free anion. This allows the N atom of SCN− to
interact with CH group from the Bmim+ cation forming N…H−C hydrogen bonds,
resulting in a vibrational mode around 2800 cm‒1. We indeed observed a band of
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about 2800 cm‒1 in the experimental IR spectra of the conventional ILs (Figure S11).
The band C, which has the highest intensity, is the C‒N stretching mode of the SCN‒
anion. The last bands D (i.e. D1, D2 and D3) belong to the Bmim+ cationÈwhich
include the stretching mode of the C-H group in CH3‒, CH2‒ and five‒membered ring.
Interestingly, we found that the C-H stretching mode in Bmim+ cation appears around
3200 cm‒1, and it shifts to 2800 cm‒1 when forming a hydrogen bond with SCN‒ anion
(Figure 6(a)).
Figure 6. Experimental IR spectra of [Bmim]2[Co(NCS)4] after (a) and before (b) absorption of
NH3. The simulated IR spectra of [Co(NH3)6]2+ (c), [Bmim][SCN] (d), free NH3 (e), [SCN]− anion
(f), and [Bmim]+ cation (g)
3.8. Comparison with other ILs
In order to evaluate the potential of ILs for NH3 absorption, NH3 absorption
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capacity in the cobalt ILs [Cnmim]2[Co(NCS)4] were compared with other ILs
reported in literature in Table 4. It was clearly seen that [Cnmim]2[Co(NCS)4] not only
have much higher gravimetric and molar capacity of NH3 absorption than those of the
reported ILs, but also exhibit excellent recyclability, which implies that the cobalt ILs
[Cnmim]2[Co(NCS)4] show great potentials for efficient and reversible removal of
NH3 from NH3-containing exhausted gas.
Table 4. Comparison of NH3 capacity with other reported ILs
Temperature Pressure
NH3 capacity
ILs
References
ꢀ
MPa
0.10
molNH3·molIL-1 gNH3·gIL-1
[Emim]2[Co(NCS)4]
[Bmim]2[Co(NCS)4]
[Hmim]2[Co(NCS)4]
[Emim][SCN]
[Bmim][SCN]
[Hmim][SCN]
[Emim][BF4]
30
5.99
6.03
6.09
0.18
0.19
0.20
0.18
0.45
0.56
0.39
0.15
0.54
1.70
0.198
0.180
0.166
0.018
0.016
0.015
0.016
0.034
0.038
0.024
0.007
0.032
0.135
This work
This work
This work
This work
This work
30
30
30
30
30
25
25
25
25
25
25
20
0.10
0.10
0.10
0.10
0.10
0.11
0.22
0.22
0.12
0.14
0.17
0.10
This work
16
16
16
16
14
14
26
[Bmim][BF4]
[Hmim][BF4]
[Omim][BF4]
[Emim][Tf2N]
[Bmim][PF6]
[EOHmim][BF4]
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26
47
18
18
18
[Choline][Tf2N]
[TMG][BF4]
20
20
40
40
40
0.10
0.10
0.10
0.10
0.10
1.86
1.08
2.69
0.28
1.54
0.082
0.046
0.113
0.011
0.058
[Bim][Tf2N]
[Bmim][Tf2N]
[HOOC(CH2)3mim][NTf2]
4. Conclusions
In summary, a new strategy to design MILs with highly efficient and reversible
absorption of NH3 through the cooperative interactions between the metal
coordination ions and NH3 has been proposed. By screening the suitable metal centers
and ligands, the cobalt ILs [Cnmim]2[Co(NCS)4] were synthesized for NH3 absorption.
The results demonstrated that [Cnmim]2[Co(NCS)4] not only exhibit superior
recyclability, but also show very high NH3 capacity up to 6.09 molNH3·molIL-1,
which is much higher than all reported ILs to date, and is over 30 times higher than
the conventional ILs [Cnmim][SCN]. Spectroscopic investigations and simulation
calculations confirmed that significant improvement in NH3 capacity and regeneration
performance of the cobalt ILs is attributed to the moderate Lewis acid-base and
cooperative hydrogen bond interactions between the metal centers-ligands and NH3.
We believe that the cobalt ILs will be potential candidates for efficient purification
and recyle of NH3 from industrial NH3-containing gases.
Acknowledgements
This work was supported by the National Natural Science Foundation of China
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(21425625, U1662122, 21506219 and 21776277), Beijing Municipal Natural Science
Foundation (2182071), and Beijing Hundreds of Leading Talents Training Project of
Science and Technology (Z171100001117154).
Notes and references
‡ The optimized structures of ILs with NH3 and CO2, the recycle of cobalt ILs with
saturated water, the experimental IR and NMR spectra, the simulated IR spectra as
well as Cartesian coordinates were shown in the Supporting information.
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Table of Contents
Cobalt ionic liquids with highly efficient and reversible absorption of NH3 through Lewis
acid-base and hydrogen bond interactions were designed.