# CO2 Capture in Ionic Liquids: A Review of Solubilities and Experimental Methods

Authors: Elena Torralba-Calleja; James Skinner; David Gutiérrez-Tauste
Journal: Journal of Chemistry
 (2013)

---

## Abstract

The growing concern of climate change and global warming has in turn given rise to a thriving research field dedicated to finding solutions. One particular area which has received considerable attention is the lowering of carbon dioxide emissions from large-scale sources, that is, fossil fuel power. This paper focuses on ionic liquids being used as novel media for CO2 capture. In particular, solubility data and experimental techniques are used at a laboratory scale. Cited CO2 absorption data for imidazolium-, pyrrolidinium-, pyridinium-, quaternary-ammonium-, and tetra-alkyl-phosphonium-based ionic liquids is reviewed, expressed as mole fractions (X) of CO2 to ionic liquid. The following experimental techniques are featured: gravimetric analysis, the pressure drop method, and the view-cell method.


---

## Body

## 1. Introduction

In recent years, increasing attention has been paid towards the worldwide climate change. Moreover, the exponential increase of carbon dioxide emissions into the atmosphere from the combustion of fossil fuels, making up the 86% of greenhouse gases [1], does not reflect a sustainable energy model. Entry into the Kyoto protocol has brought about the need to reduce anthropogenic emissions of CO2. Thus carbon capture and storage (CCS) proves to be one of the most important initiatives to mitigate this global warming effect.CCS is a concept based on the reduction of CO2 emissions into the atmosphere from industrial processes, such as ammonia production, natural gas processing, or cement manufacture, to name a few. This review however will focus on CO2 emissions from fossil fuel power plants, which is seen to be the main contributor to this effect [2]. It has been approximated that, if CCS is fully implemented, its potential by 2050 could be the total capture and storage of 236 billion tons of CO2 [3]. An approach to CCS that holds the greatest promise is the sequestration of captured carbon dioxide, in suitable deep sedimentary formations, for example, oil and depleted gas reservoirs, coal beds, and saline deposits [4–7]. The challenge is to develop a technology which will allow us to accomplish this task in an environmental, economic, and efficient way in the next years [8–10]. However the need to assess the environmental impact is great. The potential risks of geological storage to humans and ecosystems are abundant and need to be carefully monitored. Leakage of sequestered CO2 would be the main concern. This could happen along fault lines, ineffective confining layers, abandoned wells, and so forth. The pollution of groundwater and mineral deposits is also a problem and could have lethal effects on plant life and animals. A recent review by Manchao et al. [11] offers a detailed risk assessment of the CO2 injection process and storage in geological formations, with a main focus on abandoned coal mines and coal seams.An alternative to geological storage of CO2 would be the direct conversion of CO2 into a high-valued product after the initial capture; this is sometimes referred to as carbon capture and usage (CCU). CO2 is used in many industries such as the food industry (carbonation of beverages), electronics industry (surface cleaning and semiconductor manufacture), and the chemical industry (polymers, plastics, and fertilizers). CCU is yet to be a mainstream technology so that many process aspects and methods are being published and reviewed [12, 13].


## 2. State-of-the-Art CO2 Capture Technologies

The capture of CO2 is achieved through the use of specific materials that interact with the gas in one form or another. The materials that are used depend on the processes in which the flue gas is conditioned (Figure 1) [14].Figure 1
Possible techniques that can be used in conjunction with the processes of postcombustion, precombustion, and oxyfuel combustion.There are three processes, each of which conditions the CO2 for capture in different ways.Postcombustion. The separation of CO2 from the flue gas after the combustion of fuel. Air is typically used as the oxidant in this process; therefore the flue gas becomes largely diluted with nitrogen.Precombustion. The hydrocarbon fuel (in this case gasified coal) is converted into carbon monoxide (CO) and hydrogen (H2). This forms a synthesis gas. By using water shift conversion, CO is converted into CO2. Finally the CO2 is then separated from the H2.Oxyfuel CO
2
Combustion. It uses pure oxygen as the oxidant instead of air, creating a flue gas mainly consisting of high-concentrated CO2 and steam.Although CO2 capture and separation is a well-known technology, this technology is just applied in a small scale, so that right now it is not commercially available for being used in large power stations. The most challenging obstacle to overcome in CCS and CCU is finding an effective technique that satisfies environmental and economic factors. Some of the currently studied techniques for capturing CO2 from the three conditioning processes are as follows (Figure 1).Absorption occurs within the bulk of the material via a chemical or physical interaction. Chemical absorbents react with the CO2, forming covalent bonds between the molecules. The solvent can be habitually regenerated through heating and captured CO2 is released. This mechanism can also be made highly selective by the introduction of specific chemical complexes. Typical compounds used in this process are amines, or ammonia-based solutions. Physical absorbents obey Henry’s law, where gas solubility is directly proportional to the partial pressure of the said gas in equilibrium, at a constant temperature. Typically this is at high CO2 partial pressures and low temperatures. The interaction between CO2 and the solvent is by nonchemical surface forces, that is, Van der Waals interaction. Regeneration of the solvent is achieved by increasing the temperature and lowering the pressure of the system [15]. Selexol and Rectisol are examples of physical absorbents that have been used in natural gas sweetening and synthesis gas treatment.Adsorption, as opposed to absorption, takes place at the surface of the material. This interaction can also occur chemically (covalent bonding) or physically (Van der Waals). Typical adsorbers are solid materials with large surface areas, such as zeolites, activated carbons, metal oxides, silica gel, and ion-exchange resins. These can be used to capture CO2by separation, so that flue gas is put in contact with a bed of these adsorbers, allowing the CO2 capture from the other gases which pass through. When the bed is fully saturated with CO2, the flue gas is directed to a clean bed and the saturated bed is regenerated [16]. Three techniques can be employed to the adsorption mechanism: pressure swing adsorption (PSA) introduces the flue gas at high pressure until the concentration of CO2 reaches equilibrium, then the pressure is lowered to regenerate the adsorbent, temperature swing adsorption (TSA) increases the temperature to regenerate the adsorbent, and electric swing adsorption (ESA) is where a low-voltage electric current is passed through the sorbent to regenerate. Adsorption is not yet considered practical for large-scale applications as the CO2 selectivity in current sorbents is low. However, recently new sorbents are being investigated such as metal-organic frameworks and functionalised fibrous matrices that show some promise for the future of this particular technique.Membrane separation technology is based on the interaction of specific gases with the membrane material by a physical or chemical interaction. Through modifying the material, the rate at which the gases pass through can be controlled. There are wide varieties of membranes available for gas separation, including polymeric membranes, zeolites, and porous inorganic membranes, some of which are used in an industrial scale and have the possibility of being implemented into the process of CO2 capture. However achieving high degrees of CO2 separation in one single stage has so far proved to be difficult; therefore, having to rely on multiple stages has led to increasing energy consumption and cost. An alternative approach is to use porous membranes as platforms for absorption and stripping. Here a liquid (typically aqueous amine solutions) provides the selectivity towards the gases. As the flue gas moves through the membrane, the liquid selects and captures the CO2 [17].Cryogenic Separation is a technique based on cooling and condensation. This has the advantage of enabling the direct production of liquid CO2, benefiting transportation options. Although a major disadvantage of cryogenic technology in this respect is the high amounts of energy required to provide cooling for the process, this is especially prominent in low-concentration gas streams [18]. This technique is more suited to high-concentration and high-pressure gases, such as in oxyfuel combustion and precombustion.Within these techniques lie the materials with which research pathways aim to develop more effective CO2 capture mechanisms. Currently the postcombustion process is the most widely researched area for reducing CO2 emissions from power stations. This is mainly because it can be retrofitted to existing combustion systems without a great deal of modification, unlike the other two processes. The flue gas emitted, from the postcombustion of fossil fuels in power stations, has a total pressure of 1-2 bars with a CO2 concentration of approximately 15%. As this process creates low CO2 concentration and partial pressures, strong solvents have to be used to capture the CO2, resulting in a large energy input to regenerate the solvent for further use. This creates the technical challenge of finding an efficient, costeffective, and low-energy-demanding capture mechanism using novel materials.


### 2.1. Aqueous Amines Used in Postcombustion

The conventional technologies used in this postcombustion process are solvent-based chemical absorbers. The common chemical solvents used for separation are aqueous amines, which are ammonia derivatives, where one or more of the hydrogen atoms have been replaced by alkyl groups. Some common amines used in this process are (Table1) monoethanolamine (MEA) [19], methyldiethanolamine (MDEA) [20], and diethanolamine (DEA) [21]. Aqueous amines are stated as “conventional absorbers” because they are well-known solvents used in the oil and gas industries, dating back to the 1930s; for example, Gregory and Scharmann investigated the implementation of amine CO2 scrubbers in a hydrogenation plant of the Standard Oil Company of Louisiana in 1937. Today the aqueous amine absorption technology is still used in natural gas sweetening (removal of acidic gases, for example, hydrogen sulphide and carbon dioxide) and has also been applied to some small-scale fossil fuel power plants [22, 23], for example, Fundación Ciudad de la Energía (CIUDEN), Alstom power plant, and so forth.Table 1
Chemical structures of commonly used amines.
Amine
Acronym
Structure
Monoethanolamine
MEA
Methyldiethanolamine
MDEA
Diethanolamine
DEA
Piperazine
PIPABriefly, post-combustion capture with amines, seen in Figure2, involves the CO2 being removed by circulating a flue gas stream into a chamber containing an aqueous amine solution. In the case of primary amines like MEA, the CO2 is captured by a chemical absorption process in which the CO2 reacts with the amine in the form of a carbamate [24]. With secondary and tertiary amines, which do not possess a hydrogen atom attached to a nitrogen atom, they react with CO2 in the form of bicarbonate through hydrolysis. This is a reversible reaction, and at high temperatures the captured CO2 is released and the amine solution recycled. Piperazine (PIPA) is commonly used to improve reaction kinetics of secondary and tertiary amines in the form of an additive; this is because the heat of reaction to form a bicarbonate is low, causing more heat being needed for regeneration and thus higher costs [25].Figure 2
Conceptual scheme for CO2 absorption using amine-based chemical absorption.Amines are so effective for CO2 capture thanks to some of their properties such as high reactivity with CO2, high absorbing capacity (in terms of mass of CO2), relatively high thermal stability, and CO2 selectivity [26]. However there are inherent disadvantages linked with amines, which need to be addressed in order to make a valid and efficient process for CO2 capture. These disadvantages come in the form of high vapour pressure, corrosive nature, and high-energy input for regeneration. The high vapour pressure allows emission of amine gases into the air upon heating. These gases are unstable in nature thus giving them the possibility of producing dangerous toxins such as nitrosamines, nitramines, and amides. Nitrosamines are of the most concern as they are carcinogenic and toxic to humans even at low levels [27]. Amines are also corrosive, especially MEA. They take part in reactions in which waste forms and can eventually corrode the equipment, Kittel et al. [28] investigated the effects of MEA operating pilot plants and found that areas made of carbon steel had corrosion rates of 1 mm year−1; so besides environmental impacts, expense on a large industrial scale is another issue. The recycling/regeneration process leads to high-energy consumption in order to break the chemical bonds formed between the CO2 and amine [29]. This process also causes degradation of the amine which limits its CO2 capture rate, causing them to be replaced frequently.Much research has gone into developing new solvents with the foresight of being superior to amines. The factors that would allow new solvents to perform better than amines are lower cost, lower volatility, better thermal stability, less degradation, low corrosive nature, and low energy needed for regeneration and adaptability to an existing system. Although amines have high CO2 solubility and selectivity, environmental and economic effects are taken into consideration when selecting the criteria for the most suited CO2 capture mechanism. While continued research into improving the performance of these mature technologies is expected, research into novel materials and technologies could produce the significant breakthroughs required to minimise the environmental and energy penalties of capture.


### 2.2. Ionic Liquid Media for CO2 Capture

One of these advanced R&D pathways currently conveying great potential in the field of alternative technologies is ionic liquids (ILs). ILs are commonly defined as materials that are comprised of large organic cations and organic/inorganic anions, which demonstrate melting points below 100°C [30]. To date a wide range of ILs has been synthesised through different combinations of anions and cations. It has been stated that the theoretical number of potential ILs is to the order of 1018 [31]. An example of some of the common cations and anions used in IL synthesis can be seen in Table 2.Table 2
Structures of common IL cations and anions.
Cation
Structure
Anion
Structure
Imidazolium
Tetrafluoroborate
Pyrrolidinium
Hexafluoroborate
Pyridinium
Bis(trifluorophoshate)imide
Quaternary Ammonium
Nitrate
Tetra Alkyl Phosphonium
AcetateILs possess several unique and diverse characteristics such as high thermal and chemical stability, low vapour pressure, large electrochemical window, tuneable/designer nature, and excellent solvent properties for a range of polar and nonpolar compounds. It is due to these characteristics that research into developing and implementing ILs over the past decade has spanned into many sectors of industry [32], for applications such as electrolytes [33], solar cells [34], lubricants [35], electropolishing and electroplating [36], and biomass processing [37], to name a few. This has become possible due to the large number of ILs that can be synthesised in the lab [38–40] and purchased commercially. Companies like BASF, Merck, Sigma-Aldrich, Solvionic, Sachem, and IoLitec provide basic ILs and can also aid in the design and development of ILs for specific tasks. Therefore these compounds have created exciting new media for emerging technological applications already commercially available.


### 2.3. Ionic Liquids in the Scope of CO2 Capture

The aforementioned characteristics are particularly advantageous when applying ILs as solvents for CO2 capture in comparison to current aqueous amine technology:(i)
less energy is required when regenerating ILs to remove the captured CO2 [41] due to their physical absorption mechanism,(ii)
further efficiency is attained by their low vapour pressure, which allows them to be regenerated and reused with no appreciable losses into the gas stream [42, 43],(iii)
ILs have a high thermal and chemical stability; typically they degrade at temperatures >300°C [44] avoiding their reaction with impurities and causing corrosion to the equipment,(iv)
the tuneable and designer nature of ILs offers many options concerning the physicochemical properties (viscosities and densities [45–47], heat capacities [48], thermal decomposition temperatures [49], surface tension [50], toxicity and health issues [51, 52], and corrosion [53, 54]) in the sense that the anions and cations can be manipulated to create an IL for a specific task.This designer aspect can also be applied to the anion or cation in the sense that various chemical functionalities and structures can be attached, allowing properties such as absorption and viscosity to be controlled. These are commonly referred to as task-specific ionic liquids (TSILs). Generally ILs fulfil many of the major requirements stated in the green-chemistry principles stated by Anastas and Warner [55], in that they offer a new approach to industrial/chemical processes whereby steps are taken to eliminate hazardous waste in a system before a by-product is formed, thus neglecting the use of volatile organic solvent.Most development concerning ILs for CO2 capture is at present conducted at laboratory scale, while other technological applications are already in use as it was mentioned before. Conversely, their industrial application and implementation is being constantly investigated in areas of post-combustion [56]. For industrial-scale integration, it is necessary to achieve extensive knowledge of their physical and chemical properties. Therefore the need for experimental techniques and data is critical in enabling the ionic liquid to be the green, viable, and economic carbon capture technique of the future.The solubility of CO2 in ILs compared to other gases such as methane and nitrogen enables ILs to separate CO2 from the source, be it a power plants’ flue gas or natural gas. Even when there are low concentrations of CO2 in a mixed gas, the IL can be designed to incorporate a functional group, such as an amine, thus rendering it task specific. The capacity for CO2 solubility in ILs originates from the asymmetrical combination of the anion and cation, which results from short-range repulsive forces between their ionic shells. Therefore the more incompatible the ionic constituents are the greater the solubility is.


### 2.4. Conventional Ionic Liquids

Over the past decade, and at present, research has been built upon measuring the effects of variables such as pressure, temperature, and anion/cation choice. Results have shown high carbon dioxide solubility in what have become known as conventional ionic liquids. They are defined as ILs that do not possess an attached functional group and have been reported by many as portraying the typical behaviour of physical solvents [57–59]. This is evident when low-pressure CO2 (1-2 bars) is put in contact with the IL, resulting in low CO2 concentrations in the liquid phase. As the increment of pressure increases, typically to up to 100 bar, the concentration of absorbed CO2 increases. Thus displaying the general characteristics of a physical absorber. As a rule, the solubility of CO2 in ILs increases with increasing pressure and decreases with increasing temperature. The physical absorption mechanism is a result of the interaction between the CO2 molecules and the IL, in which the CO2 occupies the “free space” within the ILs structure through a large quadrupole moment and Van der Waals forces.


#### 2.4.1. Anion and Cation Effects

In order to create an optimal process for capturing CO2 in ILs, assessment of the essential building blocks, that is, cation/anion combinations, needs to be investigated. Synthesising ILs that encompass CO2-philic groups on the anion such as carbonyls or fluorines has proven to increase CO2 capture [60]. In the past decade studies have shown that the origin of high solubility is strongly dependent on the choice of anion [61]. Aki et al. [62] investigated the influence of the anion with seven ILs. They all contained the 1-butyl-3-methylimidazolium [Bmim] cation. The results are shown in Table 3.Table 3
Influence of anions in different ionic liquids.
Anion
Nomenclature
Classification
Solubility of CO2 in IL
Dicyanamide
[DCA]−
Nonfluorinated anions
Low
Nitrate
[NO3]−
Tetrafluoroborate
[BF4]−
Fluorinated anions
Relatively high
Hexafluorophosphate
[PF6]−
Trifuoromethanesulfonate
[TfO]−
Bis(trifluoromethylsulfonyl) imide
[Tf2N]−
Tris(trifluoromethylsulfonyl) methide
[methide]−Aki and coworkers also systematically investigated the effects of the cation on CO2 solubility; they found that, in general, the increase of the alkyl chain on the cation resulted in a slight increase in solubility, which became more apparent at higher pressures. The effect of increasing the alkyl chain results in the increased volume available for CO2 interaction. Muldoon et al. [60] concluded that adding partially fluorinated alkyl chains on the imidazolium cation does increase CO2 solubility. They compared [hmim][Tf2N] directly to [C6H4F9mim][Tf2N] and found that this increased solubility was due to fluorinating the last four carbons of the alkyl chain. Research on IL CO2 solubility, in general, has focused intensively on imidazolium-based structures. However some groups have focused on using different cations. Recently Carvalho et al. [63] reported CO2 solubilities in two phosphonium-based ILs, [THTDP][Tf2N] and [THTDP][Cl]. They found exceptionally high solubility measurements exceeding those of current imidazolium-based ILs; they go on to conclude that their study shows the highest recorded solubility observed without chemical interactions in the absorption process. Although imidazolium is the most stable and commercially available cation of choice, it is evident that there are further enhancements and possibilities that can be developed from other bases.To provide further insight into the interactions between CO2 and the constituent anions and cations of RTILs, researches using spectroscopic approaches and molecular simulations have been made. Of which has broadened our understanding of absorption mechanisms and structure-property relationships, Kazarian et al. [64] used ATR-FTIR spectroscopy to analyse the specific interactions of CO2 and ILs [Bmim][BF4] and [Bmim][PF6]. They saw evidence of chemical interactions between the anion [PF6]− and CO2. They concluded that they observed weak Lewis acid-base interactions, where the anion acts as a Lewis base. ILs by their nature have intrinsic acid-base properties. These properties can be enhanced with the addition of acidic functions like carbonic or halide acids; likewise, basic functions like amino and fluorine groups can be added. This has shown to create specific Lewis acid-base chemical interactions between CO2 and the IL.As it can be seen in Tables4 and 5, fluorination of the anion and in some cases the cation can improve CO2 solubility in RTILs. However the associated disadvantages are cost increase, poor degradability, and a negative environmental impact [65]. Therefore paths to develop ILs with enhanced CO2 solubility without fluorination are also being investigated.Table 4
CO2 solubility data for imidazolium-based ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
1-N-Octyl-3-methylimidazolium hexafluorophosphate
C8mim[PF6]
313
92.67
0.7550
Blanchard et al. 2001 [86]
1-N-Butyl-3-nethylimidazolium nitrate
Bmim[NO3]
323
92.62
0.5300
Blanchard et al. 2001 [86]
1-N-Octyl-3-methylimidazolium tetrafluoroborate
C8mim[BF4]
313
92.90
0.7080
Blanchard et al. 2001 [86]
1-Ethyl-3-methylimidazolium ethyl sulfate
Emim[EtSO4]
333
94.61
0.4570
Blanchard et al. 2001 [86]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
313
96.67
0.7290
Blanchard et al. 2001 [86]
1-Butyl-3-methylimidazolium acetate
C4mim[Ac]
333.3
12.75
0.2510
Carvalho et al. 2009 [87]
323.09
755.26
0.5990
Carvalho et al. 2009 [87]
1-Butyl-3-methylimidazolium trifluoroacetate
C4mim[TFA]
293.43
9.79
0.2250
Carvalho et al. 2009 [87]
293.59
436.25
0.6790
Carvalho et al. 2009 [87]
1-Butyl,3-methyl-imidazolium tetrafluoroborate
Bmim[BF4]
303
10
0.1461
Galan-Sanchez 2008 [88]
333
10
0.0895
Galan-Sanchez 2008 [88]
1-Octyl,3-methyl-imidazolium tetrafluoroborate
Omim[BF4]
303
10
0.1873
Galan-Sanchez 2008 [88]
333
10
0.1213
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium dicyanamide
Bmim[DCA]
303
10
0.1434
Galan-Sanchez 2008 [88]
333
10
0.0997
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
303
10
0.0978
Galan-Sanchez 2008 [88]
333
10
0.0664
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
303
10
0.1662
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
333
10
0.1012
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium methylsulfate
Bmim[MeSO4]
303
10
0.1190
Galan-Sanchez 2008 [88]
333
10
0.0733
Galan-Sanchez 2008 [88]
1-N-Ethyl-3-mehylimidazolium bis(trifluoromethylsulfonyl)Imide
Emim[NTf2]
303
10
0.2257
Galan-Sanchez 2008 [88]
333
10
0.1446
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
298.15
6.66
0.122
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium hexafluorophosphate
C6mim[PF6]
298.15
9.27
0.167
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium tetrafluoroborate
Emim[BF4]
298.15
8.75
0.106
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium tetrafluoroborate
C6mim[BF4]
298.15
8.99
0.163
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.15
9.03
0.209
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
C6mim[Tf2N]
298.15
8.59
0.236
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium trifluoromethane-sulfonate
C2mim[TfO]
303.85
149
0.6260
Shin and Lee 2008 [90]
303.85
15
0.2610
Shin and Lee 2008 [90]
1-Butyl-3-methylimidazolium trifluoromethane-sulfonate
C4mim[TfO]
303.85
160
0.6720
Shin and Lee 2008 [90]
303.85
11.5
0.2730
Shin and Lee 2008 [90]
1-Hexyl-3-methylimidazolium trifluoromethane-sulfonate
C6mim[TfO]
303.85
180
0.7170
Shin and Lee 2008 [90]
303.85
12.5
0.2880
Shin and Lee 2008 [90]
1-Octyl-3-methylimidazolium trifluoromethane-sulfonate
C8mim[TfO]
303.85
180
0.7410
Shin and Lee 2008 [90]
303.85
15.8
0.3440
Shin and Lee 2008 [90]
1,3-Dimethylimidazolium methylphosphonate
Dmim[MP]
313.35
95
0.4750
Revelli et al. 2010 [91]
313.45
34
0.1620
Revelli et al. 2010  [91]
1-Butyl,3-methyl-imidazolium tetrafluroborate
Bmim[BF4]
293.65
73
0.6100
Revelli et al. 2010 [91]
293.25
10.5
0.1410
Revelli et al. 2010 [91]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
313.65
99
0.4300
Revelli et al. 2010 [91]
292.35
10.5
0.1260
Revelli et al. 2010 [91]
1-Ethyl-3-methylimidazolium trifluoroacetate
Emim[TFA]
298.1
19.99
0.2820
Yokozeki et al.  2008 [67]
1-Ethyl-3-methylimidazolium acetate
Emim[Ac]
298.1
19.99
0.4280
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium trifluoroacetate
Bmim[TFA]
298.1
19.99
0.3010
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium acetate
Bmim[Ac]
298.1
19.99
0.4550
Yokozeki et al.  2008 [67]
1-Ethyl-3-methyl-imidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.1
19.99
0.3900
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium tris(pentafluoroethyl)trifluoro-phosphate
Hmim[FAP]
298.1
19.99
0.4930
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Hmim[Tf2N]
298.1
19.74
0.4330
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium 1,1,2,2-tetrafluoroethanesulfonate
Bmim[TFES]
298
19.9
0.2850
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium propionate
Bmim[PRO]
298.2
19.9
0.3900
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium isobutyrate
Bmim[ISB]
298.2
20
0.4030
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium trimethylacetate
Bmim[TMA]
298.1
19.9
0.4310
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium levulinate
Bmim[LEV]
298.1
19.9
0.4600
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium succinamate
Bmim[SUC]
298.1
19.9
0.2320
Yokozeki et al.  2008 [67]
Bis(1-butyl-3-methylimidazolium) iminodiacetate
Bmim2[IDA]
298.1
19.9
0.3950
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium iminoacetic acid acetate
Bmim[IAAc]
298.1
19.9
0.1910
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium  tris(pentafluoroethyl)trifluorophosphate
Hmim[FEP]
283.5
17.99
0.5170
Zhang et al. 2008 [92]Table 5
CO2 solubility for ammonium ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
Bis(2-hydroxyethyl)-ammonium acetate
(BHEAA)
298.15
15.15
0.1076
Kurnia et al. 2009 [78]
298.15
5.48
0.0391
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium acetate
(HHEMEA)
298.15
15.42
0.0761
Kurnia et al. 2009 [78]
298.15
6.15
0.0300
Kurnia et al. 2009 [78]
Bis(2-hydroxyethyl)-ammonium lactate
(BHEAL)
298.15
15.12
0.0835
Kurnia et al. 2009 [78]
298.15
3.46
0.0192
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium lactate
(HHEMEL)
298.15
15.23
0.0776
Kurnia et al. 2009 [78]
298.15
3.48
0.0179
Kurnia et al. 2009 [78]
2-Hydroxy ethyl ammonium formate
(HEF)
303
78.9
0.3083
Yuan et al. 2007 [93]
303
4.4
0.0340
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium acetate
(HEA)
303
90.1
0.4009
Yuan et al. 2007 [93]
303
8.9
0.0687
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium lactate
(HEL)
303
82
0.2422
Yuan et al. 2007 [93]
303
7.8
0.0410
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium acetate
(THEAA)
303
82.5
0.2561
Yuan et al. 2007 [93]
303
10.3
0.0534
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium lactate
(THEAL)
303
70.9
0.4617
Yuan et al. 2007 [93]
303
9.6
0.1006
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium formate
(HEAF)
303
72.8
0.1907
Yuan et al. 2007 [93]
303
6.6
0.0300
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium acetate
(HEAA)
303
65.7
0.4860
Yuan et al. 2007 [93]
303
7.6
0.0889
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
73.2
0.2640
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
12.4
0.0704
Yuan et al. 2007 [93]Due to certain limitations of conventional ionic liquid systems, where physical absorption takes place and high solubility is only seen at high pressures, numerous research groups have been developing the ILs designer character, by covalently tethering a functional group to either or both anion or cation. This resulting functionalized IL is capable of chemically binding to CO2, adding chemical absorption to the capture mechanism.[Bmim][Ac] has been found to be one of these RTILs in which a chemical complexion with CO2 occurs [66]. In 2008, Yokozeki et al. [67] completed CO2 solubility tests for 18 RTILs, eight of which showed chemical absorption mechanisms. They found that RTILs that show strong chemical absorption with CO2 all contain the anion [X-COO]−, that is, [Bmim][Ac], [Emim][Ac], [Bmim][PRO], [Bmim][IBS], [Bmim][TMA], and [Bmim][LEU]. Their results can be seen in Table 4. In general it is assumed that conventional RTILs with acidic or basic functionalities strongly influence the absorption of CO2.As discussed, RTILs sufficiently absorb CO2 especially those containing CO2-philic groups like fluorine. These are known as TSILs (task-specific ionic liquids). Widely researched TSILs are those with appended amine group, examples of which can be seen in Table 5. Bates and coworkers [68] synthesized the amine functionalized IL [pNH2Bim][Pf6] and found it to chemically react with the CO2. The CO2 reacts with the amine on the IL, this then reacts with another amine and forms an ammonium carbamate double salt. This form of capture results in one CO2 captured for every two ILs. This 1 : 2 capture mechanism is also observed on the molecular level with traditional aqueous amines. It is theoretically suggested that, when amines are tethered to the anion only, a 1 : 1 ratio can be met allowing a more efficient process.Evidence has shown that TSILs have the ability to absorb CO2 both chemically and physically. At low pressures (typically below 2 bars) chemical absorption takes place, in the same way as aqueous amines. After the majority of the chemical bonding have taken place, physical absorption dominates the capture mechanism; this is especially relevant at high pressures, whereas aqueous amines reach their absorption limits at low pressure. This shows how the absorption performances of TSILs with amine functionalities merge the characteristics of physical solvents with the attractive features of chemical solvents. In spite of TSILs showing greater CO2 solubility than conventional RTILs, they tend to exhibit high viscosity in comparison to other commercially available absorbents. This poses a large problem for their implementation into large-scale platforms, as the heat required for absorption and regeneration would be a lot larger and energy intensive. In order to reduce the viscosity, some groups have combined mixtures of TSIL and RTIL. Bara et al. [69] dissolved their TSIL in a common RTIL, [C6mim][Tf2N]. Although the solution was stable and capable of absorbing in a 1 : 2 molar ratio, the viscosity was still high. As a whole TSILs and TSILs + RTILs are robust and have a high absorption capacity; however, they are limited by the intensive synthesis that is required, high viscosity, and the fact that the TSIL serves as both the capture material and the dispersant.Instead of the direct incorporation of amino-functionalized anions and cations, some recent groups have reported using imidazolium-based RTILs with amines added in solution to act as the capture reagent. Camper and coworkers [70] first investigated this concept. They synthesized an [Rmim][Tf2N] RTIL solution containing 16% v/v of MEA and found that this is capable of rapid and reversible capture of one mole of CO2 per two moles of MEA at low CO2 partial pressures. An MEA-carbamate was found to precipitate from the RTIL solution; this helps to drive the capture reaction. They have currently seen that this MEA-carbamate seems to be a consequence of the [Tf2N] anion and does not occur in other [Cnmim][X] RTILs.


### 2.5. CO2 Solubility Results Reported by Various Experimental Groups

Tables4, 5, 6, and 7 aim to provide a range of experimental data cited by various experimental groups, for peak CO2 absorption values for different cation-based ILs. This can then be used to characterise an experimental system to ensure correct implementation and method. T (K) represents the system’s temperature when measurements were recorded. P (bar) is the corresponding pressure of CO2. XCO2 is the solubility of CO2 expressed as a mole fraction, that is, moles of CO2 to moles of IL. The tables also attempt to show the effects of temperature and pressure on ILs as well as different cation and anion combinations.Table 6
CO2 solubility for phosphonium, pyridinium and pyrrolidinium ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
N-Butylpyridinium tetrafluoroborate
N-BuPy[BF4]
323
92.35
0.5810
Blanchard et al. 2001 [86]
Trihexyltetradecylphosphonium chloride
THTDP[Cl]
302.55
149.95
0.8000
Carvalho et al. 2010 [94]
313.27
5.17
0.2000
Carvalho et al. 2010 [94]
Trihexyltetradecylphosphonium bis(trifluoromethylsulfonyl)imide
THTDP[NTf2]
296.58
721.85
0.8790
Carvalho et al. 2010 [94]
293.2
6.12
0.3080
Carvalho et al. 2010 [94]
N-Butyl-4-methylpyridinium tetrafluoroborate
MeBuPy[BF4]
303
10
0.1443
Galan-Sanchez 2008 [88]
333
10
0.0961
Galan-Sanchez 2008 [88]
N-Butyl-3-Methylpyridinium dicyanamide
MeBuPy[DCA]
303
10
0.1436
Galan-Sanchez 2008 [88]
333
10
0.0683
Galan-Sanchez 2008 [88]
N-Butyl-4-Methylpyridinium thiocyanate
MeBuPy[SCN]
303
10
0.0962
Galan-Sanchez 2008 [88]
333
10
0.0632
Galan-Sanchez 2008 [88]
1-Butyl-1-methylpyrrolidinium dicyanamide
MeBuPyrr[DCA]
303
10
0.1204
Galan-Sanchez 2008 [88]
333
10
0.0613
Galan-Sanchez 2008 [88]
1-Butyl-1-Methylpyrrolidinium thiocyanate
MeBuPyrr[SCN]
303
10
0.0971
Galan-Sanchez 2008 [88]
333
10
0.0608
Galan-Sanchez 2008 [88]
1-Butyl-1-methylpyrrolidinium trifluoroacetate
MeBuPyrr[TFA]
303
10
0.1674
Galan-Sanchez 2008 [88]
333
10
0.1030
Galan-Sanchez 2008 [88]
Tetrabutylphosphonium formate
TBP[FOR]
298.1
19.9
0.3480
Yokozeki et al. 2008 [67]
1-Butyl-1-methylpyrrolidinium tri(pentafluoroethyl)trifluorophosphate
BmPyrr[FEP]
283.5
18.00
0.4980
Zhang et al. 2008 [92]Table 7
CO2 Solubility data for functionalized ionic liquids (TSILs).
Acronym
Functionalization
Anion
Cation
T (K)
P (bar)
χCO2
References
APMim[NTf2]
NH2-cation
NTf2
Im
303
10.00
0.27
Galan-Sanchez 2008 [88]
343
10.00
0.18
APMim[DCA]
NH2-cation
DCA
Im
303
10.00
0.29
Galan-Sanchez 2008 [88]
APMim[BF4]
NH2-cation
BF4
Im
303
10.00
0.32
Galan-Sanchez 2008 [88]
343
10.00
0.36
AEMPyrr[BF4]
NH2-cation
BF4
Pyrr
303
10.00
0.28
Galan-Sanchez 2008 [88]
333
10.00
0.24
MeImNet2[BF4]
NR3-cation
BF4
Im
303
4.00
0.09
Galan-Sanchez 2008 [88]
Bmim[Tau]
NH2-Anion
Taureate
Im
333
10.00
0.43
Galan-Sanchez 2008 [88]
Bmim[Gly]
NH2-Anion
Glycinate
Im
333
10.00
0.39
Galan-Sanchez 2008 [88]


## 2.1. Aqueous Amines Used in Postcombustion

The conventional technologies used in this postcombustion process are solvent-based chemical absorbers. The common chemical solvents used for separation are aqueous amines, which are ammonia derivatives, where one or more of the hydrogen atoms have been replaced by alkyl groups. Some common amines used in this process are (Table1) monoethanolamine (MEA) [19], methyldiethanolamine (MDEA) [20], and diethanolamine (DEA) [21]. Aqueous amines are stated as “conventional absorbers” because they are well-known solvents used in the oil and gas industries, dating back to the 1930s; for example, Gregory and Scharmann investigated the implementation of amine CO2 scrubbers in a hydrogenation plant of the Standard Oil Company of Louisiana in 1937. Today the aqueous amine absorption technology is still used in natural gas sweetening (removal of acidic gases, for example, hydrogen sulphide and carbon dioxide) and has also been applied to some small-scale fossil fuel power plants [22, 23], for example, Fundación Ciudad de la Energía (CIUDEN), Alstom power plant, and so forth.Table 1
Chemical structures of commonly used amines.
Amine
Acronym
Structure
Monoethanolamine
MEA
Methyldiethanolamine
MDEA
Diethanolamine
DEA
Piperazine
PIPABriefly, post-combustion capture with amines, seen in Figure2, involves the CO2 being removed by circulating a flue gas stream into a chamber containing an aqueous amine solution. In the case of primary amines like MEA, the CO2 is captured by a chemical absorption process in which the CO2 reacts with the amine in the form of a carbamate [24]. With secondary and tertiary amines, which do not possess a hydrogen atom attached to a nitrogen atom, they react with CO2 in the form of bicarbonate through hydrolysis. This is a reversible reaction, and at high temperatures the captured CO2 is released and the amine solution recycled. Piperazine (PIPA) is commonly used to improve reaction kinetics of secondary and tertiary amines in the form of an additive; this is because the heat of reaction to form a bicarbonate is low, causing more heat being needed for regeneration and thus higher costs [25].Figure 2
Conceptual scheme for CO2 absorption using amine-based chemical absorption.Amines are so effective for CO2 capture thanks to some of their properties such as high reactivity with CO2, high absorbing capacity (in terms of mass of CO2), relatively high thermal stability, and CO2 selectivity [26]. However there are inherent disadvantages linked with amines, which need to be addressed in order to make a valid and efficient process for CO2 capture. These disadvantages come in the form of high vapour pressure, corrosive nature, and high-energy input for regeneration. The high vapour pressure allows emission of amine gases into the air upon heating. These gases are unstable in nature thus giving them the possibility of producing dangerous toxins such as nitrosamines, nitramines, and amides. Nitrosamines are of the most concern as they are carcinogenic and toxic to humans even at low levels [27]. Amines are also corrosive, especially MEA. They take part in reactions in which waste forms and can eventually corrode the equipment, Kittel et al. [28] investigated the effects of MEA operating pilot plants and found that areas made of carbon steel had corrosion rates of 1 mm year−1; so besides environmental impacts, expense on a large industrial scale is another issue. The recycling/regeneration process leads to high-energy consumption in order to break the chemical bonds formed between the CO2 and amine [29]. This process also causes degradation of the amine which limits its CO2 capture rate, causing them to be replaced frequently.Much research has gone into developing new solvents with the foresight of being superior to amines. The factors that would allow new solvents to perform better than amines are lower cost, lower volatility, better thermal stability, less degradation, low corrosive nature, and low energy needed for regeneration and adaptability to an existing system. Although amines have high CO2 solubility and selectivity, environmental and economic effects are taken into consideration when selecting the criteria for the most suited CO2 capture mechanism. While continued research into improving the performance of these mature technologies is expected, research into novel materials and technologies could produce the significant breakthroughs required to minimise the environmental and energy penalties of capture.


## 2.2. Ionic Liquid Media for CO2 Capture

One of these advanced R&D pathways currently conveying great potential in the field of alternative technologies is ionic liquids (ILs). ILs are commonly defined as materials that are comprised of large organic cations and organic/inorganic anions, which demonstrate melting points below 100°C [30]. To date a wide range of ILs has been synthesised through different combinations of anions and cations. It has been stated that the theoretical number of potential ILs is to the order of 1018 [31]. An example of some of the common cations and anions used in IL synthesis can be seen in Table 2.Table 2
Structures of common IL cations and anions.
Cation
Structure
Anion
Structure
Imidazolium
Tetrafluoroborate
Pyrrolidinium
Hexafluoroborate
Pyridinium
Bis(trifluorophoshate)imide
Quaternary Ammonium
Nitrate
Tetra Alkyl Phosphonium
AcetateILs possess several unique and diverse characteristics such as high thermal and chemical stability, low vapour pressure, large electrochemical window, tuneable/designer nature, and excellent solvent properties for a range of polar and nonpolar compounds. It is due to these characteristics that research into developing and implementing ILs over the past decade has spanned into many sectors of industry [32], for applications such as electrolytes [33], solar cells [34], lubricants [35], electropolishing and electroplating [36], and biomass processing [37], to name a few. This has become possible due to the large number of ILs that can be synthesised in the lab [38–40] and purchased commercially. Companies like BASF, Merck, Sigma-Aldrich, Solvionic, Sachem, and IoLitec provide basic ILs and can also aid in the design and development of ILs for specific tasks. Therefore these compounds have created exciting new media for emerging technological applications already commercially available.


## 2.3. Ionic Liquids in the Scope of CO2 Capture

The aforementioned characteristics are particularly advantageous when applying ILs as solvents for CO2 capture in comparison to current aqueous amine technology:(i)
less energy is required when regenerating ILs to remove the captured CO2 [41] due to their physical absorption mechanism,(ii)
further efficiency is attained by their low vapour pressure, which allows them to be regenerated and reused with no appreciable losses into the gas stream [42, 43],(iii)
ILs have a high thermal and chemical stability; typically they degrade at temperatures >300°C [44] avoiding their reaction with impurities and causing corrosion to the equipment,(iv)
the tuneable and designer nature of ILs offers many options concerning the physicochemical properties (viscosities and densities [45–47], heat capacities [48], thermal decomposition temperatures [49], surface tension [50], toxicity and health issues [51, 52], and corrosion [53, 54]) in the sense that the anions and cations can be manipulated to create an IL for a specific task.This designer aspect can also be applied to the anion or cation in the sense that various chemical functionalities and structures can be attached, allowing properties such as absorption and viscosity to be controlled. These are commonly referred to as task-specific ionic liquids (TSILs). Generally ILs fulfil many of the major requirements stated in the green-chemistry principles stated by Anastas and Warner [55], in that they offer a new approach to industrial/chemical processes whereby steps are taken to eliminate hazardous waste in a system before a by-product is formed, thus neglecting the use of volatile organic solvent.Most development concerning ILs for CO2 capture is at present conducted at laboratory scale, while other technological applications are already in use as it was mentioned before. Conversely, their industrial application and implementation is being constantly investigated in areas of post-combustion [56]. For industrial-scale integration, it is necessary to achieve extensive knowledge of their physical and chemical properties. Therefore the need for experimental techniques and data is critical in enabling the ionic liquid to be the green, viable, and economic carbon capture technique of the future.The solubility of CO2 in ILs compared to other gases such as methane and nitrogen enables ILs to separate CO2 from the source, be it a power plants’ flue gas or natural gas. Even when there are low concentrations of CO2 in a mixed gas, the IL can be designed to incorporate a functional group, such as an amine, thus rendering it task specific. The capacity for CO2 solubility in ILs originates from the asymmetrical combination of the anion and cation, which results from short-range repulsive forces between their ionic shells. Therefore the more incompatible the ionic constituents are the greater the solubility is.


## 2.4. Conventional Ionic Liquids

Over the past decade, and at present, research has been built upon measuring the effects of variables such as pressure, temperature, and anion/cation choice. Results have shown high carbon dioxide solubility in what have become known as conventional ionic liquids. They are defined as ILs that do not possess an attached functional group and have been reported by many as portraying the typical behaviour of physical solvents [57–59]. This is evident when low-pressure CO2 (1-2 bars) is put in contact with the IL, resulting in low CO2 concentrations in the liquid phase. As the increment of pressure increases, typically to up to 100 bar, the concentration of absorbed CO2 increases. Thus displaying the general characteristics of a physical absorber. As a rule, the solubility of CO2 in ILs increases with increasing pressure and decreases with increasing temperature. The physical absorption mechanism is a result of the interaction between the CO2 molecules and the IL, in which the CO2 occupies the “free space” within the ILs structure through a large quadrupole moment and Van der Waals forces.


### 2.4.1. Anion and Cation Effects

In order to create an optimal process for capturing CO2 in ILs, assessment of the essential building blocks, that is, cation/anion combinations, needs to be investigated. Synthesising ILs that encompass CO2-philic groups on the anion such as carbonyls or fluorines has proven to increase CO2 capture [60]. In the past decade studies have shown that the origin of high solubility is strongly dependent on the choice of anion [61]. Aki et al. [62] investigated the influence of the anion with seven ILs. They all contained the 1-butyl-3-methylimidazolium [Bmim] cation. The results are shown in Table 3.Table 3
Influence of anions in different ionic liquids.
Anion
Nomenclature
Classification
Solubility of CO2 in IL
Dicyanamide
[DCA]−
Nonfluorinated anions
Low
Nitrate
[NO3]−
Tetrafluoroborate
[BF4]−
Fluorinated anions
Relatively high
Hexafluorophosphate
[PF6]−
Trifuoromethanesulfonate
[TfO]−
Bis(trifluoromethylsulfonyl) imide
[Tf2N]−
Tris(trifluoromethylsulfonyl) methide
[methide]−Aki and coworkers also systematically investigated the effects of the cation on CO2 solubility; they found that, in general, the increase of the alkyl chain on the cation resulted in a slight increase in solubility, which became more apparent at higher pressures. The effect of increasing the alkyl chain results in the increased volume available for CO2 interaction. Muldoon et al. [60] concluded that adding partially fluorinated alkyl chains on the imidazolium cation does increase CO2 solubility. They compared [hmim][Tf2N] directly to [C6H4F9mim][Tf2N] and found that this increased solubility was due to fluorinating the last four carbons of the alkyl chain. Research on IL CO2 solubility, in general, has focused intensively on imidazolium-based structures. However some groups have focused on using different cations. Recently Carvalho et al. [63] reported CO2 solubilities in two phosphonium-based ILs, [THTDP][Tf2N] and [THTDP][Cl]. They found exceptionally high solubility measurements exceeding those of current imidazolium-based ILs; they go on to conclude that their study shows the highest recorded solubility observed without chemical interactions in the absorption process. Although imidazolium is the most stable and commercially available cation of choice, it is evident that there are further enhancements and possibilities that can be developed from other bases.To provide further insight into the interactions between CO2 and the constituent anions and cations of RTILs, researches using spectroscopic approaches and molecular simulations have been made. Of which has broadened our understanding of absorption mechanisms and structure-property relationships, Kazarian et al. [64] used ATR-FTIR spectroscopy to analyse the specific interactions of CO2 and ILs [Bmim][BF4] and [Bmim][PF6]. They saw evidence of chemical interactions between the anion [PF6]− and CO2. They concluded that they observed weak Lewis acid-base interactions, where the anion acts as a Lewis base. ILs by their nature have intrinsic acid-base properties. These properties can be enhanced with the addition of acidic functions like carbonic or halide acids; likewise, basic functions like amino and fluorine groups can be added. This has shown to create specific Lewis acid-base chemical interactions between CO2 and the IL.As it can be seen in Tables4 and 5, fluorination of the anion and in some cases the cation can improve CO2 solubility in RTILs. However the associated disadvantages are cost increase, poor degradability, and a negative environmental impact [65]. Therefore paths to develop ILs with enhanced CO2 solubility without fluorination are also being investigated.Table 4
CO2 solubility data for imidazolium-based ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
1-N-Octyl-3-methylimidazolium hexafluorophosphate
C8mim[PF6]
313
92.67
0.7550
Blanchard et al. 2001 [86]
1-N-Butyl-3-nethylimidazolium nitrate
Bmim[NO3]
323
92.62
0.5300
Blanchard et al. 2001 [86]
1-N-Octyl-3-methylimidazolium tetrafluoroborate
C8mim[BF4]
313
92.90
0.7080
Blanchard et al. 2001 [86]
1-Ethyl-3-methylimidazolium ethyl sulfate
Emim[EtSO4]
333
94.61
0.4570
Blanchard et al. 2001 [86]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
313
96.67
0.7290
Blanchard et al. 2001 [86]
1-Butyl-3-methylimidazolium acetate
C4mim[Ac]
333.3
12.75
0.2510
Carvalho et al. 2009 [87]
323.09
755.26
0.5990
Carvalho et al. 2009 [87]
1-Butyl-3-methylimidazolium trifluoroacetate
C4mim[TFA]
293.43
9.79
0.2250
Carvalho et al. 2009 [87]
293.59
436.25
0.6790
Carvalho et al. 2009 [87]
1-Butyl,3-methyl-imidazolium tetrafluoroborate
Bmim[BF4]
303
10
0.1461
Galan-Sanchez 2008 [88]
333
10
0.0895
Galan-Sanchez 2008 [88]
1-Octyl,3-methyl-imidazolium tetrafluoroborate
Omim[BF4]
303
10
0.1873
Galan-Sanchez 2008 [88]
333
10
0.1213
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium dicyanamide
Bmim[DCA]
303
10
0.1434
Galan-Sanchez 2008 [88]
333
10
0.0997
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
303
10
0.0978
Galan-Sanchez 2008 [88]
333
10
0.0664
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
303
10
0.1662
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
333
10
0.1012
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium methylsulfate
Bmim[MeSO4]
303
10
0.1190
Galan-Sanchez 2008 [88]
333
10
0.0733
Galan-Sanchez 2008 [88]
1-N-Ethyl-3-mehylimidazolium bis(trifluoromethylsulfonyl)Imide
Emim[NTf2]
303
10
0.2257
Galan-Sanchez 2008 [88]
333
10
0.1446
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
298.15
6.66
0.122
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium hexafluorophosphate
C6mim[PF6]
298.15
9.27
0.167
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium tetrafluoroborate
Emim[BF4]
298.15
8.75
0.106
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium tetrafluoroborate
C6mim[BF4]
298.15
8.99
0.163
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.15
9.03
0.209
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
C6mim[Tf2N]
298.15
8.59
0.236
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium trifluoromethane-sulfonate
C2mim[TfO]
303.85
149
0.6260
Shin and Lee 2008 [90]
303.85
15
0.2610
Shin and Lee 2008 [90]
1-Butyl-3-methylimidazolium trifluoromethane-sulfonate
C4mim[TfO]
303.85
160
0.6720
Shin and Lee 2008 [90]
303.85
11.5
0.2730
Shin and Lee 2008 [90]
1-Hexyl-3-methylimidazolium trifluoromethane-sulfonate
C6mim[TfO]
303.85
180
0.7170
Shin and Lee 2008 [90]
303.85
12.5
0.2880
Shin and Lee 2008 [90]
1-Octyl-3-methylimidazolium trifluoromethane-sulfonate
C8mim[TfO]
303.85
180
0.7410
Shin and Lee 2008 [90]
303.85
15.8
0.3440
Shin and Lee 2008 [90]
1,3-Dimethylimidazolium methylphosphonate
Dmim[MP]
313.35
95
0.4750
Revelli et al. 2010 [91]
313.45
34
0.1620
Revelli et al. 2010  [91]
1-Butyl,3-methyl-imidazolium tetrafluroborate
Bmim[BF4]
293.65
73
0.6100
Revelli et al. 2010 [91]
293.25
10.5
0.1410
Revelli et al. 2010 [91]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
313.65
99
0.4300
Revelli et al. 2010 [91]
292.35
10.5
0.1260
Revelli et al. 2010 [91]
1-Ethyl-3-methylimidazolium trifluoroacetate
Emim[TFA]
298.1
19.99
0.2820
Yokozeki et al.  2008 [67]
1-Ethyl-3-methylimidazolium acetate
Emim[Ac]
298.1
19.99
0.4280
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium trifluoroacetate
Bmim[TFA]
298.1
19.99
0.3010
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium acetate
Bmim[Ac]
298.1
19.99
0.4550
Yokozeki et al.  2008 [67]
1-Ethyl-3-methyl-imidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.1
19.99
0.3900
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium tris(pentafluoroethyl)trifluoro-phosphate
Hmim[FAP]
298.1
19.99
0.4930
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Hmim[Tf2N]
298.1
19.74
0.4330
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium 1,1,2,2-tetrafluoroethanesulfonate
Bmim[TFES]
298
19.9
0.2850
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium propionate
Bmim[PRO]
298.2
19.9
0.3900
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium isobutyrate
Bmim[ISB]
298.2
20
0.4030
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium trimethylacetate
Bmim[TMA]
298.1
19.9
0.4310
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium levulinate
Bmim[LEV]
298.1
19.9
0.4600
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium succinamate
Bmim[SUC]
298.1
19.9
0.2320
Yokozeki et al.  2008 [67]
Bis(1-butyl-3-methylimidazolium) iminodiacetate
Bmim2[IDA]
298.1
19.9
0.3950
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium iminoacetic acid acetate
Bmim[IAAc]
298.1
19.9
0.1910
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium  tris(pentafluoroethyl)trifluorophosphate
Hmim[FEP]
283.5
17.99
0.5170
Zhang et al. 2008 [92]Table 5
CO2 solubility for ammonium ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
Bis(2-hydroxyethyl)-ammonium acetate
(BHEAA)
298.15
15.15
0.1076
Kurnia et al. 2009 [78]
298.15
5.48
0.0391
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium acetate
(HHEMEA)
298.15
15.42
0.0761
Kurnia et al. 2009 [78]
298.15
6.15
0.0300
Kurnia et al. 2009 [78]
Bis(2-hydroxyethyl)-ammonium lactate
(BHEAL)
298.15
15.12
0.0835
Kurnia et al. 2009 [78]
298.15
3.46
0.0192
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium lactate
(HHEMEL)
298.15
15.23
0.0776
Kurnia et al. 2009 [78]
298.15
3.48
0.0179
Kurnia et al. 2009 [78]
2-Hydroxy ethyl ammonium formate
(HEF)
303
78.9
0.3083
Yuan et al. 2007 [93]
303
4.4
0.0340
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium acetate
(HEA)
303
90.1
0.4009
Yuan et al. 2007 [93]
303
8.9
0.0687
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium lactate
(HEL)
303
82
0.2422
Yuan et al. 2007 [93]
303
7.8
0.0410
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium acetate
(THEAA)
303
82.5
0.2561
Yuan et al. 2007 [93]
303
10.3
0.0534
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium lactate
(THEAL)
303
70.9
0.4617
Yuan et al. 2007 [93]
303
9.6
0.1006
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium formate
(HEAF)
303
72.8
0.1907
Yuan et al. 2007 [93]
303
6.6
0.0300
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium acetate
(HEAA)
303
65.7
0.4860
Yuan et al. 2007 [93]
303
7.6
0.0889
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
73.2
0.2640
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
12.4
0.0704
Yuan et al. 2007 [93]Due to certain limitations of conventional ionic liquid systems, where physical absorption takes place and high solubility is only seen at high pressures, numerous research groups have been developing the ILs designer character, by covalently tethering a functional group to either or both anion or cation. This resulting functionalized IL is capable of chemically binding to CO2, adding chemical absorption to the capture mechanism.[Bmim][Ac] has been found to be one of these RTILs in which a chemical complexion with CO2 occurs [66]. In 2008, Yokozeki et al. [67] completed CO2 solubility tests for 18 RTILs, eight of which showed chemical absorption mechanisms. They found that RTILs that show strong chemical absorption with CO2 all contain the anion [X-COO]−, that is, [Bmim][Ac], [Emim][Ac], [Bmim][PRO], [Bmim][IBS], [Bmim][TMA], and [Bmim][LEU]. Their results can be seen in Table 4. In general it is assumed that conventional RTILs with acidic or basic functionalities strongly influence the absorption of CO2.As discussed, RTILs sufficiently absorb CO2 especially those containing CO2-philic groups like fluorine. These are known as TSILs (task-specific ionic liquids). Widely researched TSILs are those with appended amine group, examples of which can be seen in Table 5. Bates and coworkers [68] synthesized the amine functionalized IL [pNH2Bim][Pf6] and found it to chemically react with the CO2. The CO2 reacts with the amine on the IL, this then reacts with another amine and forms an ammonium carbamate double salt. This form of capture results in one CO2 captured for every two ILs. This 1 : 2 capture mechanism is also observed on the molecular level with traditional aqueous amines. It is theoretically suggested that, when amines are tethered to the anion only, a 1 : 1 ratio can be met allowing a more efficient process.Evidence has shown that TSILs have the ability to absorb CO2 both chemically and physically. At low pressures (typically below 2 bars) chemical absorption takes place, in the same way as aqueous amines. After the majority of the chemical bonding have taken place, physical absorption dominates the capture mechanism; this is especially relevant at high pressures, whereas aqueous amines reach their absorption limits at low pressure. This shows how the absorption performances of TSILs with amine functionalities merge the characteristics of physical solvents with the attractive features of chemical solvents. In spite of TSILs showing greater CO2 solubility than conventional RTILs, they tend to exhibit high viscosity in comparison to other commercially available absorbents. This poses a large problem for their implementation into large-scale platforms, as the heat required for absorption and regeneration would be a lot larger and energy intensive. In order to reduce the viscosity, some groups have combined mixtures of TSIL and RTIL. Bara et al. [69] dissolved their TSIL in a common RTIL, [C6mim][Tf2N]. Although the solution was stable and capable of absorbing in a 1 : 2 molar ratio, the viscosity was still high. As a whole TSILs and TSILs + RTILs are robust and have a high absorption capacity; however, they are limited by the intensive synthesis that is required, high viscosity, and the fact that the TSIL serves as both the capture material and the dispersant.Instead of the direct incorporation of amino-functionalized anions and cations, some recent groups have reported using imidazolium-based RTILs with amines added in solution to act as the capture reagent. Camper and coworkers [70] first investigated this concept. They synthesized an [Rmim][Tf2N] RTIL solution containing 16% v/v of MEA and found that this is capable of rapid and reversible capture of one mole of CO2 per two moles of MEA at low CO2 partial pressures. An MEA-carbamate was found to precipitate from the RTIL solution; this helps to drive the capture reaction. They have currently seen that this MEA-carbamate seems to be a consequence of the [Tf2N] anion and does not occur in other [Cnmim][X] RTILs.


## 2.4.1. Anion and Cation Effects

In order to create an optimal process for capturing CO2 in ILs, assessment of the essential building blocks, that is, cation/anion combinations, needs to be investigated. Synthesising ILs that encompass CO2-philic groups on the anion such as carbonyls or fluorines has proven to increase CO2 capture [60]. In the past decade studies have shown that the origin of high solubility is strongly dependent on the choice of anion [61]. Aki et al. [62] investigated the influence of the anion with seven ILs. They all contained the 1-butyl-3-methylimidazolium [Bmim] cation. The results are shown in Table 3.Table 3
Influence of anions in different ionic liquids.
Anion
Nomenclature
Classification
Solubility of CO2 in IL
Dicyanamide
[DCA]−
Nonfluorinated anions
Low
Nitrate
[NO3]−
Tetrafluoroborate
[BF4]−
Fluorinated anions
Relatively high
Hexafluorophosphate
[PF6]−
Trifuoromethanesulfonate
[TfO]−
Bis(trifluoromethylsulfonyl) imide
[Tf2N]−
Tris(trifluoromethylsulfonyl) methide
[methide]−Aki and coworkers also systematically investigated the effects of the cation on CO2 solubility; they found that, in general, the increase of the alkyl chain on the cation resulted in a slight increase in solubility, which became more apparent at higher pressures. The effect of increasing the alkyl chain results in the increased volume available for CO2 interaction. Muldoon et al. [60] concluded that adding partially fluorinated alkyl chains on the imidazolium cation does increase CO2 solubility. They compared [hmim][Tf2N] directly to [C6H4F9mim][Tf2N] and found that this increased solubility was due to fluorinating the last four carbons of the alkyl chain. Research on IL CO2 solubility, in general, has focused intensively on imidazolium-based structures. However some groups have focused on using different cations. Recently Carvalho et al. [63] reported CO2 solubilities in two phosphonium-based ILs, [THTDP][Tf2N] and [THTDP][Cl]. They found exceptionally high solubility measurements exceeding those of current imidazolium-based ILs; they go on to conclude that their study shows the highest recorded solubility observed without chemical interactions in the absorption process. Although imidazolium is the most stable and commercially available cation of choice, it is evident that there are further enhancements and possibilities that can be developed from other bases.To provide further insight into the interactions between CO2 and the constituent anions and cations of RTILs, researches using spectroscopic approaches and molecular simulations have been made. Of which has broadened our understanding of absorption mechanisms and structure-property relationships, Kazarian et al. [64] used ATR-FTIR spectroscopy to analyse the specific interactions of CO2 and ILs [Bmim][BF4] and [Bmim][PF6]. They saw evidence of chemical interactions between the anion [PF6]− and CO2. They concluded that they observed weak Lewis acid-base interactions, where the anion acts as a Lewis base. ILs by their nature have intrinsic acid-base properties. These properties can be enhanced with the addition of acidic functions like carbonic or halide acids; likewise, basic functions like amino and fluorine groups can be added. This has shown to create specific Lewis acid-base chemical interactions between CO2 and the IL.As it can be seen in Tables4 and 5, fluorination of the anion and in some cases the cation can improve CO2 solubility in RTILs. However the associated disadvantages are cost increase, poor degradability, and a negative environmental impact [65]. Therefore paths to develop ILs with enhanced CO2 solubility without fluorination are also being investigated.Table 4
CO2 solubility data for imidazolium-based ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
1-N-Octyl-3-methylimidazolium hexafluorophosphate
C8mim[PF6]
313
92.67
0.7550
Blanchard et al. 2001 [86]
1-N-Butyl-3-nethylimidazolium nitrate
Bmim[NO3]
323
92.62
0.5300
Blanchard et al. 2001 [86]
1-N-Octyl-3-methylimidazolium tetrafluoroborate
C8mim[BF4]
313
92.90
0.7080
Blanchard et al. 2001 [86]
1-Ethyl-3-methylimidazolium ethyl sulfate
Emim[EtSO4]
333
94.61
0.4570
Blanchard et al. 2001 [86]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
313
96.67
0.7290
Blanchard et al. 2001 [86]
1-Butyl-3-methylimidazolium acetate
C4mim[Ac]
333.3
12.75
0.2510
Carvalho et al. 2009 [87]
323.09
755.26
0.5990
Carvalho et al. 2009 [87]
1-Butyl-3-methylimidazolium trifluoroacetate
C4mim[TFA]
293.43
9.79
0.2250
Carvalho et al. 2009 [87]
293.59
436.25
0.6790
Carvalho et al. 2009 [87]
1-Butyl,3-methyl-imidazolium tetrafluoroborate
Bmim[BF4]
303
10
0.1461
Galan-Sanchez 2008 [88]
333
10
0.0895
Galan-Sanchez 2008 [88]
1-Octyl,3-methyl-imidazolium tetrafluoroborate
Omim[BF4]
303
10
0.1873
Galan-Sanchez 2008 [88]
333
10
0.1213
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium dicyanamide
Bmim[DCA]
303
10
0.1434
Galan-Sanchez 2008 [88]
333
10
0.0997
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
303
10
0.0978
Galan-Sanchez 2008 [88]
333
10
0.0664
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
303
10
0.1662
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
333
10
0.1012
Galan-Sanchez 2008 [88]
1-Butyl-3-methylimidazolium methylsulfate
Bmim[MeSO4]
303
10
0.1190
Galan-Sanchez 2008 [88]
333
10
0.0733
Galan-Sanchez 2008 [88]
1-N-Ethyl-3-mehylimidazolium bis(trifluoromethylsulfonyl)Imide
Emim[NTf2]
303
10
0.2257
Galan-Sanchez 2008 [88]
333
10
0.1446
Galan-Sanchez 2008 [88]
1-Butyl,3-methyl-imidazolium hexafluorophosphate
Bmim[PF6]
298.15
6.66
0.122
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium hexafluorophosphate
C6mim[PF6]
298.15
9.27
0.167
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium tetrafluoroborate
Emim[BF4]
298.15
8.75
0.106
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium tetrafluoroborate
C6mim[BF4]
298.15
8.99
0.163
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.15
9.03
0.209
Kim et al. 2005 [89]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
C6mim[Tf2N]
298.15
8.59
0.236
Kim et al. 2005 [89]
1-Ethyl-3-methylimidazolium trifluoromethane-sulfonate
C2mim[TfO]
303.85
149
0.6260
Shin and Lee 2008 [90]
303.85
15
0.2610
Shin and Lee 2008 [90]
1-Butyl-3-methylimidazolium trifluoromethane-sulfonate
C4mim[TfO]
303.85
160
0.6720
Shin and Lee 2008 [90]
303.85
11.5
0.2730
Shin and Lee 2008 [90]
1-Hexyl-3-methylimidazolium trifluoromethane-sulfonate
C6mim[TfO]
303.85
180
0.7170
Shin and Lee 2008 [90]
303.85
12.5
0.2880
Shin and Lee 2008 [90]
1-Octyl-3-methylimidazolium trifluoromethane-sulfonate
C8mim[TfO]
303.85
180
0.7410
Shin and Lee 2008 [90]
303.85
15.8
0.3440
Shin and Lee 2008 [90]
1,3-Dimethylimidazolium methylphosphonate
Dmim[MP]
313.35
95
0.4750
Revelli et al. 2010 [91]
313.45
34
0.1620
Revelli et al. 2010  [91]
1-Butyl,3-methyl-imidazolium tetrafluroborate
Bmim[BF4]
293.65
73
0.6100
Revelli et al. 2010 [91]
293.25
10.5
0.1410
Revelli et al. 2010 [91]
1-Butyl-3-methylimidazolium thiocyanate
Bmim[SCN]
313.65
99
0.4300
Revelli et al. 2010 [91]
292.35
10.5
0.1260
Revelli et al. 2010 [91]
1-Ethyl-3-methylimidazolium trifluoroacetate
Emim[TFA]
298.1
19.99
0.2820
Yokozeki et al.  2008 [67]
1-Ethyl-3-methylimidazolium acetate
Emim[Ac]
298.1
19.99
0.4280
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium trifluoroacetate
Bmim[TFA]
298.1
19.99
0.3010
Yokozeki et al. 2008 [67]
1-Butyl-3-methylimidazolium acetate
Bmim[Ac]
298.1
19.99
0.4550
Yokozeki et al.  2008 [67]
1-Ethyl-3-methyl-imidazolium bis(trifluoromethylsulfonyl)imide
Emim[Tf2N]
298.1
19.99
0.3900
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium tris(pentafluoroethyl)trifluoro-phosphate
Hmim[FAP]
298.1
19.99
0.4930
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide
Hmim[Tf2N]
298.1
19.74
0.4330
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium 1,1,2,2-tetrafluoroethanesulfonate
Bmim[TFES]
298
19.9
0.2850
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium propionate
Bmim[PRO]
298.2
19.9
0.3900
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium isobutyrate
Bmim[ISB]
298.2
20
0.4030
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium trimethylacetate
Bmim[TMA]
298.1
19.9
0.4310
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium levulinate
Bmim[LEV]
298.1
19.9
0.4600
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium succinamate
Bmim[SUC]
298.1
19.9
0.2320
Yokozeki et al.  2008 [67]
Bis(1-butyl-3-methylimidazolium) iminodiacetate
Bmim2[IDA]
298.1
19.9
0.3950
Yokozeki et al.  2008 [67]
1-Butyl-3-methylimidazolium iminoacetic acid acetate
Bmim[IAAc]
298.1
19.9
0.1910
Yokozeki et al.  2008 [67]
1-Hexyl-3-methylimidazolium  tris(pentafluoroethyl)trifluorophosphate
Hmim[FEP]
283.5
17.99
0.5170
Zhang et al. 2008 [92]Table 5
CO2 solubility for ammonium ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
Bis(2-hydroxyethyl)-ammonium acetate
(BHEAA)
298.15
15.15
0.1076
Kurnia et al. 2009 [78]
298.15
5.48
0.0391
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium acetate
(HHEMEA)
298.15
15.42
0.0761
Kurnia et al. 2009 [78]
298.15
6.15
0.0300
Kurnia et al. 2009 [78]
Bis(2-hydroxyethyl)-ammonium lactate
(BHEAL)
298.15
15.12
0.0835
Kurnia et al. 2009 [78]
298.15
3.46
0.0192
Kurnia et al. 2009 [78]
2-Hydroxy-N-(2-hydroxyethyl)-N-methylethanaminium lactate
(HHEMEL)
298.15
15.23
0.0776
Kurnia et al. 2009 [78]
298.15
3.48
0.0179
Kurnia et al. 2009 [78]
2-Hydroxy ethyl ammonium formate
(HEF)
303
78.9
0.3083
Yuan et al. 2007 [93]
303
4.4
0.0340
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium acetate
(HEA)
303
90.1
0.4009
Yuan et al. 2007 [93]
303
8.9
0.0687
Yuan et al. 2007 [93]
2-Hydroxy ethyl ammonium lactate
(HEL)
303
82
0.2422
Yuan et al. 2007 [93]
303
7.8
0.0410
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium acetate
(THEAA)
303
82.5
0.2561
Yuan et al. 2007 [93]
303
10.3
0.0534
Yuan et al. 2007 [93]
Tri-(2-hydroxyethyl)-ammonium lactate
(THEAL)
303
70.9
0.4617
Yuan et al. 2007 [93]
303
9.6
0.1006
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium formate
(HEAF)
303
72.8
0.1907
Yuan et al. 2007 [93]
303
6.6
0.0300
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium acetate
(HEAA)
303
65.7
0.4860
Yuan et al. 2007 [93]
303
7.6
0.0889
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
73.2
0.2640
Yuan et al. 2007 [93]
2-(2-Hydroxyethoxy)-ammonium lactate
(HEAL)
303
12.4
0.0704
Yuan et al. 2007 [93]Due to certain limitations of conventional ionic liquid systems, where physical absorption takes place and high solubility is only seen at high pressures, numerous research groups have been developing the ILs designer character, by covalently tethering a functional group to either or both anion or cation. This resulting functionalized IL is capable of chemically binding to CO2, adding chemical absorption to the capture mechanism.[Bmim][Ac] has been found to be one of these RTILs in which a chemical complexion with CO2 occurs [66]. In 2008, Yokozeki et al. [67] completed CO2 solubility tests for 18 RTILs, eight of which showed chemical absorption mechanisms. They found that RTILs that show strong chemical absorption with CO2 all contain the anion [X-COO]−, that is, [Bmim][Ac], [Emim][Ac], [Bmim][PRO], [Bmim][IBS], [Bmim][TMA], and [Bmim][LEU]. Their results can be seen in Table 4. In general it is assumed that conventional RTILs with acidic or basic functionalities strongly influence the absorption of CO2.As discussed, RTILs sufficiently absorb CO2 especially those containing CO2-philic groups like fluorine. These are known as TSILs (task-specific ionic liquids). Widely researched TSILs are those with appended amine group, examples of which can be seen in Table 5. Bates and coworkers [68] synthesized the amine functionalized IL [pNH2Bim][Pf6] and found it to chemically react with the CO2. The CO2 reacts with the amine on the IL, this then reacts with another amine and forms an ammonium carbamate double salt. This form of capture results in one CO2 captured for every two ILs. This 1 : 2 capture mechanism is also observed on the molecular level with traditional aqueous amines. It is theoretically suggested that, when amines are tethered to the anion only, a 1 : 1 ratio can be met allowing a more efficient process.Evidence has shown that TSILs have the ability to absorb CO2 both chemically and physically. At low pressures (typically below 2 bars) chemical absorption takes place, in the same way as aqueous amines. After the majority of the chemical bonding have taken place, physical absorption dominates the capture mechanism; this is especially relevant at high pressures, whereas aqueous amines reach their absorption limits at low pressure. This shows how the absorption performances of TSILs with amine functionalities merge the characteristics of physical solvents with the attractive features of chemical solvents. In spite of TSILs showing greater CO2 solubility than conventional RTILs, they tend to exhibit high viscosity in comparison to other commercially available absorbents. This poses a large problem for their implementation into large-scale platforms, as the heat required for absorption and regeneration would be a lot larger and energy intensive. In order to reduce the viscosity, some groups have combined mixtures of TSIL and RTIL. Bara et al. [69] dissolved their TSIL in a common RTIL, [C6mim][Tf2N]. Although the solution was stable and capable of absorbing in a 1 : 2 molar ratio, the viscosity was still high. As a whole TSILs and TSILs + RTILs are robust and have a high absorption capacity; however, they are limited by the intensive synthesis that is required, high viscosity, and the fact that the TSIL serves as both the capture material and the dispersant.Instead of the direct incorporation of amino-functionalized anions and cations, some recent groups have reported using imidazolium-based RTILs with amines added in solution to act as the capture reagent. Camper and coworkers [70] first investigated this concept. They synthesized an [Rmim][Tf2N] RTIL solution containing 16% v/v of MEA and found that this is capable of rapid and reversible capture of one mole of CO2 per two moles of MEA at low CO2 partial pressures. An MEA-carbamate was found to precipitate from the RTIL solution; this helps to drive the capture reaction. They have currently seen that this MEA-carbamate seems to be a consequence of the [Tf2N] anion and does not occur in other [Cnmim][X] RTILs.


## 2.5. CO2 Solubility Results Reported by Various Experimental Groups

Tables4, 5, 6, and 7 aim to provide a range of experimental data cited by various experimental groups, for peak CO2 absorption values for different cation-based ILs. This can then be used to characterise an experimental system to ensure correct implementation and method. T (K) represents the system’s temperature when measurements were recorded. P (bar) is the corresponding pressure of CO2. XCO2 is the solubility of CO2 expressed as a mole fraction, that is, moles of CO2 to moles of IL. The tables also attempt to show the effects of temperature and pressure on ILs as well as different cation and anion combinations.Table 6
CO2 solubility for phosphonium, pyridinium and pyrrolidinium ionic liquids.
Ionic liquid
Acronym
T (K)
P (bar)
χCO2
References
N-Butylpyridinium tetrafluoroborate
N-BuPy[BF4]
323
92.35
0.5810
Blanchard et al. 2001 [86]
Trihexyltetradecylphosphonium chloride
THTDP[Cl]
302.55
149.95
0.8000
Carvalho et al. 2010 [94]
313.27
5.17
0.2000
Carvalho et al. 2010 [94]
Trihexyltetradecylphosphonium bis(trifluoromethylsulfonyl)imide
THTDP[NTf2]
296.58
721.85
0.8790
Carvalho et al. 2010 [94]
293.2
6.12
0.3080
Carvalho et al. 2010 [94]
N-Butyl-4-methylpyridinium tetrafluoroborate
MeBuPy[BF4]
303
10
0.1443
Galan-Sanchez 2008 [88]
333
10
0.0961
Galan-Sanchez 2008 [88]
N-Butyl-3-Methylpyridinium dicyanamide
MeBuPy[DCA]
303
10
0.1436
Galan-Sanchez 2008 [88]
333
10
0.0683
Galan-Sanchez 2008 [88]
N-Butyl-4-Methylpyridinium thiocyanate
MeBuPy[SCN]
303
10
0.0962
Galan-Sanchez 2008 [88]
333
10
0.0632
Galan-Sanchez 2008 [88]
1-Butyl-1-methylpyrrolidinium dicyanamide
MeBuPyrr[DCA]
303
10
0.1204
Galan-Sanchez 2008 [88]
333
10
0.0613
Galan-Sanchez 2008 [88]
1-Butyl-1-Methylpyrrolidinium thiocyanate
MeBuPyrr[SCN]
303
10
0.0971
Galan-Sanchez 2008 [88]
333
10
0.0608
Galan-Sanchez 2008 [88]
1-Butyl-1-methylpyrrolidinium trifluoroacetate
MeBuPyrr[TFA]
303
10
0.1674
Galan-Sanchez 2008 [88]
333
10
0.1030
Galan-Sanchez 2008 [88]
Tetrabutylphosphonium formate
TBP[FOR]
298.1
19.9
0.3480
Yokozeki et al. 2008 [67]
1-Butyl-1-methylpyrrolidinium tri(pentafluoroethyl)trifluorophosphate
BmPyrr[FEP]
283.5
18.00
0.4980
Zhang et al. 2008 [92]Table 7
CO2 Solubility data for functionalized ionic liquids (TSILs).
Acronym
Functionalization
Anion
Cation
T (K)
P (bar)
χCO2
References
APMim[NTf2]
NH2-cation
NTf2
Im
303
10.00
0.27
Galan-Sanchez 2008 [88]
343
10.00
0.18
APMim[DCA]
NH2-cation
DCA
Im
303
10.00
0.29
Galan-Sanchez 2008 [88]
APMim[BF4]
NH2-cation
BF4
Im
303
10.00
0.32
Galan-Sanchez 2008 [88]
343
10.00
0.36
AEMPyrr[BF4]
NH2-cation
BF4
Pyrr
303
10.00
0.28
Galan-Sanchez 2008 [88]
333
10.00
0.24
MeImNet2[BF4]
NR3-cation
BF4
Im
303
4.00
0.09
Galan-Sanchez 2008 [88]
Bmim[Tau]
NH2-Anion
Taureate
Im
333
10.00
0.43
Galan-Sanchez 2008 [88]
Bmim[Gly]
NH2-Anion
Glycinate
Im
333
10.00
0.39
Galan-Sanchez 2008 [88]


## 3. Experimental and Measurement Techniques

In order to integrate CO2 separation techniques into large industrial systems, one needs to experimentally determine the ILs gas solubility in order to characterise the carrying capacity and selectiveness. These measurements can be accumulated via a number of experimental techniques, in which factors such as pressure and temperature can be controlled. The variety of techniques used for measuring solubility for high- and low-pressure phase equilibrium is vast and the naming of these techniques tends to vary from author to another. However all the techniques fall into two categories both of which are dependent on the equilibrium phases and mixture composition. If these two factors are unknown, measurements can be carried out analytically (analytical method); if the mixture is prepared with a precisely known composition, the synthetic method can be used. The experimental and measurement techniques reviewed here are gravimetric analysis, pressure drop method, and view-cell method and gas chromatography. All of which are being specifically applied to pure CO2 solubility in ionic liquids. It is important to remember that impurities can occur in the gas and liquid, affecting the accuracy and precision of the results. Therefore degassing the liquid fully before analysis allows an accurate determination of the true solubility of the gas. This also relies on allowing true equilibrium conditions to be met between the gas and the liquid.


### 3.1. Gravimetric Analysis

Gravimetric analysis is an analytical method which describes the quantitative determination of, in this case, gas solubility by measuring the overall weight change of a sample during absorption. The gravimetric method is most commonly applied when the analyte is converted into a solid; however, as ILs are nonvolatile in nature and exhibit properties such as low vapour pressure, they can be used to a great effect with this method of analysis. Gravimetric gas analysers are used in laboratories conducting both fundamental studies into the physical properties of ILs and applications where the ability to measure gas solubility is of interest. The basic components of a typical gravimetric instrument can be seen in Figure3.Figure 3
Basic components of a gravimetric system.High-precision gravimetric instruments are commercially available, in the form of thermogravimetric microbalances. These allow in situ measurements of gas absorption that record the mass gain of a sample with a high-precision electrobalance, which is capable of taking readings at high temperature and pressure. Also available are analysers which use magnetic suspension balances rather than an electrobalance [71]. The main difference between these two weighing systems is that, in magnetic suspension, the sample is weighed from the outside. Therefore the balance is not in physical contact with the high temperature and pressures subjected to the sample. This particular system is helpful when working with samples under extreme conditions. Petermann et al. [72] show the use and experimental setup of a magnetic suspension balance in conjunction with a volumetric determination method. The advantages of using magnetic suspension balances are also discussed by Dreisbach and Lösch [73].Gravimetric analysis systems often measure gas solubility by recording isotherms, isobars, and kinetic sorption data, which can be output through a computer from which the system can be controlled. Hence when a sample is loaded, the operation of the instrument can be fully automated and programmed to carry out isothermal absorption and desorption measurements.Due to gravimetric balances undergoing constant changes in temperature and pressure during measurements and the high sensitivity in which they operate, readings must be corrected for the changes in buoyant forces on the sample. In some apparatus, a counterweight side, which is symmetrical to the sample side, is used to minimise these effects. However they still need to be considered. Liu et al. [74] show a concise approach to calculate this.A detailed experimental procedure using the gravimetric balance can be seen in [67]. Also measurements of CO2 solubility for two imidazolium-based ILs using a thermogravimetric microbalance can be found in [75].


### 3.2. The Pressure Drop Method

The pressure drop method is a synthetic technique that is widely used in this scientific community and is also known as the isochoric method. In this instance the volume of the system is held constant, as well as the temperature, and the pressure difference is recorded during gas absorption into the sample. This method for working out gas solubility is practically suited for ILs as they have negligible vapour pressure, therefore ensuring that the gas phase remains pure, and therefore the assumption can be made that changes in pressure are due to gas sorption. From an initial measurement of pressure, temperature, and volume, and a final measurement of these variables at equilibrium, the amount of gas absorbed by the IL can be calculated. This calculation can be performed using an equation of state to convert all three variables into moles of gas.The basic principles of this method are as follows: CO2 gas is transferred into a reservoir of known volume and brought to a constant system temperature. An initial reading of pressure is measured. By using a PVT relation, the moles of CO2 in the reservoir are calculated. The IL is loaded into an equilibrium cell/stainless steel reactor and equalized to system temperature. The CO2 is then introduced to the ionic liquid and the pressure drop is recorded when the cell’s pressure remains stable; this is the equilibrium point. From the pressure drop measured, the number of moles of CO2 left in the gas phase can be calculated. The difference between CO2 mole values corresponds to the amount of gas absorbed in the IL. A typical setup for the pressure drop method can be seen in Figure 4.Figure 4
The pressure drop apparatus, whereP and T correspond to pressure and temperature sensors.The moles of dissolved CO2 in the ionic liquid can be calculated by (1).Number of CO2 moles dissolved in the ionic liquid
(1)nCO2=PinitialVGRZCO2(Pinitial,Tinitial)RTinitial-Peq(Vtot-VIL)ZCO2(Peq,Teq)RTeq,Pinitial and Tinitial are the initial pressure and temperature in the gas reservoir. Peq and Teq are the pressure and temperature at equilibrium in the equilibrium cell. Vtot is the total volume of the entire apparatus. VIL is the volume of the ionic liquid, assumed to be constant. R is the ideal gas constant. ZCO2 is the compressibility factor for CO2; this modifies the ideal gas to account for real gas behaviour. A detailed experimental procedure and full calculations for CO2 solubility measurements using the pressure drop method can be seen in [76, 77].Further investigations that utilize this pressure drop method to derive gas solubility can be found where alternative experimental setups are shown [78–80].


### 3.3. View-Cell Methods

These involve the preparation of a mixture with a precisely known composition and then the observation of phase behavior inside an equilibrium cell, where measurements are recorded in the equilibrium state, that is, temperature and pressure. Synthetic methods consist of two main techniques, one being with a phase transition, and the other without. In synthetic methods with a phase transition a known amount of gas and IL is loaded into the equilibrium cell. The pressure is then varied at a constant temperature (or vice versa) until a second phase is formed, where the gas dissolves in the ionic liquid causing the vapor phase to diminish, whereby using different gas pressures, solubility can be worked out at various pressure, and temperatures.In synthetic methods without a phase transition, equilibrium properties like temperature, pressure, density, cell volume, and gas/liquid phase volumes are measured, and the composition of the phase mixtures can be calculated in terms of moles or by a mass balance equation.As can be seen in Figure5, a pump releases CO2 at a constant selected pressure and monitors the volume of CO2 flowing into the system. The CO2 is also heated to a constant temperature. By monitoring the volume, a known amount of CO2 is then introduced to the high-pressure view cell, which contains a known amount of IL. In the case of non-phase transition, the amount of CO2 absorbed is calculated by the difference in the amount of gas delivered to the cell and the amount of gas in the vapor phase. The amount of gas in the vapor phase can be calculated using a mass balance, shown in (2), coupled with an equation of state.Equation to calculate amount of gas in the vapor phase:(2)mg=mpump-mlines-mheadspace+mlines0+mheadspace0,where mg is the mass of CO2 in the liquid phase, mpump is the mass of CO2 injected into the system, mlines is the mass of CO2 in the gas lines, connecting the pump to the equilibrium cell, mheadspace is the mass of the gas in the headspace of the cell, mlines0 is the mass of the gas in the lines after venting the system, and mheadspace0 is the mass of gas in the headspace, initially in the system after venting.Figure 5
Scheme of a synthetic method setup.A full experimental procedure using a synthetic method without phase transition and demonstrating the use of mass balancing to determine gas solubility is explained in the literature [30, 81, 82].


### 3.4. Further Techniques

Gas chromatography is an analytical method that boosts high precision and accuracy. When applied to measuring gas solubility in absorption media, the gas chromatograph is usually coupled with a high-pressure reactor cell in which a synthetic or pressure drop method is applied, and at equilibrium, a sample is taken and analyzed [83]. Solubility data from gas chromatography can be achieved by using an extractive technique; here the solvent (IL) is saturated with the solute (CO2) and then coated on a column. Nitrogen, or any other nonabsorbing carrier gas, is directed on to the column in order to extract the CO2 from the IL. The nitrogen is then analyzed in the gas chromatograph. This determines the amount of CO2 removed (per amount of coating). A detailed method for applying gas chromatography can be found in the thesis by Wilbanks [84].Other analytical techniques can be used in some cases to determine the solubility of specific gases; this may be in the form of a titration; this was demonstrated by Shen and Li [85] with aqueous amine solutions. However this has so far not been applied to ILs. Inline gas sensors also have the potential to be used. A possible scenario could involve linking an electrochemical sensor to measure the difference in CO2 concentration of the vapor phase before and after equilibrium conditions.Many advantages come from using gravimetric microbalances for solubility measurements. The ability to measure mass change to a high precision is helpful for a variety of reasons. When initially degassing the ionic liquid sample, being able to measure mass decrease allows the experimentalist to see when a constant mass value has been reached, thus allowing the assumption that full degassing has occurred. Also, the mass reading is very important to ensure equilibrium conditions once the CO2 has been introduced. Equilibrium is reached when the mass change is zero. In situ gravimetric balances, that is, when the balance is enclosed in the measuring gas, are limited to lower pressures and temperatures. Disadvantages of gravimetric systems are mainly due to their high retail price, making them impractical for small projects where the funding is restricted.In comparison to the gravimetric analysis, the pressure drop and synthetic methods are much simpler in design. As samples of any size can be investigated with these methods, a high sensitivity can generally be achieved, however, not a high accuracy. The most significant errors in the pressure drop and synthetic methods are the error calculations of dead space; for gravimetric methods, it is the determination of buoyancy forces. In pressure drop and synthetic methods, the two variables, pressure and gas absorbed, are determined by the pressure sensors and calibrated volumes; this can result in measuring error which is added on each step of the absorption isotherm. With the gravimetric method, all of the variables, temperature, gas pressure, and absorbed gas, are measured independently and the absorption pressure is monitored at each step of the isotherm.


## 3.1. Gravimetric Analysis

Gravimetric analysis is an analytical method which describes the quantitative determination of, in this case, gas solubility by measuring the overall weight change of a sample during absorption. The gravimetric method is most commonly applied when the analyte is converted into a solid; however, as ILs are nonvolatile in nature and exhibit properties such as low vapour pressure, they can be used to a great effect with this method of analysis. Gravimetric gas analysers are used in laboratories conducting both fundamental studies into the physical properties of ILs and applications where the ability to measure gas solubility is of interest. The basic components of a typical gravimetric instrument can be seen in Figure3.Figure 3
Basic components of a gravimetric system.High-precision gravimetric instruments are commercially available, in the form of thermogravimetric microbalances. These allow in situ measurements of gas absorption that record the mass gain of a sample with a high-precision electrobalance, which is capable of taking readings at high temperature and pressure. Also available are analysers which use magnetic suspension balances rather than an electrobalance [71]. The main difference between these two weighing systems is that, in magnetic suspension, the sample is weighed from the outside. Therefore the balance is not in physical contact with the high temperature and pressures subjected to the sample. This particular system is helpful when working with samples under extreme conditions. Petermann et al. [72] show the use and experimental setup of a magnetic suspension balance in conjunction with a volumetric determination method. The advantages of using magnetic suspension balances are also discussed by Dreisbach and Lösch [73].Gravimetric analysis systems often measure gas solubility by recording isotherms, isobars, and kinetic sorption data, which can be output through a computer from which the system can be controlled. Hence when a sample is loaded, the operation of the instrument can be fully automated and programmed to carry out isothermal absorption and desorption measurements.Due to gravimetric balances undergoing constant changes in temperature and pressure during measurements and the high sensitivity in which they operate, readings must be corrected for the changes in buoyant forces on the sample. In some apparatus, a counterweight side, which is symmetrical to the sample side, is used to minimise these effects. However they still need to be considered. Liu et al. [74] show a concise approach to calculate this.A detailed experimental procedure using the gravimetric balance can be seen in [67]. Also measurements of CO2 solubility for two imidazolium-based ILs using a thermogravimetric microbalance can be found in [75].


## 3.2. The Pressure Drop Method

The pressure drop method is a synthetic technique that is widely used in this scientific community and is also known as the isochoric method. In this instance the volume of the system is held constant, as well as the temperature, and the pressure difference is recorded during gas absorption into the sample. This method for working out gas solubility is practically suited for ILs as they have negligible vapour pressure, therefore ensuring that the gas phase remains pure, and therefore the assumption can be made that changes in pressure are due to gas sorption. From an initial measurement of pressure, temperature, and volume, and a final measurement of these variables at equilibrium, the amount of gas absorbed by the IL can be calculated. This calculation can be performed using an equation of state to convert all three variables into moles of gas.The basic principles of this method are as follows: CO2 gas is transferred into a reservoir of known volume and brought to a constant system temperature. An initial reading of pressure is measured. By using a PVT relation, the moles of CO2 in the reservoir are calculated. The IL is loaded into an equilibrium cell/stainless steel reactor and equalized to system temperature. The CO2 is then introduced to the ionic liquid and the pressure drop is recorded when the cell’s pressure remains stable; this is the equilibrium point. From the pressure drop measured, the number of moles of CO2 left in the gas phase can be calculated. The difference between CO2 mole values corresponds to the amount of gas absorbed in the IL. A typical setup for the pressure drop method can be seen in Figure 4.Figure 4
The pressure drop apparatus, whereP and T correspond to pressure and temperature sensors.The moles of dissolved CO2 in the ionic liquid can be calculated by (1).Number of CO2 moles dissolved in the ionic liquid
(1)nCO2=PinitialVGRZCO2(Pinitial,Tinitial)RTinitial-Peq(Vtot-VIL)ZCO2(Peq,Teq)RTeq,Pinitial and Tinitial are the initial pressure and temperature in the gas reservoir. Peq and Teq are the pressure and temperature at equilibrium in the equilibrium cell. Vtot is the total volume of the entire apparatus. VIL is the volume of the ionic liquid, assumed to be constant. R is the ideal gas constant. ZCO2 is the compressibility factor for CO2; this modifies the ideal gas to account for real gas behaviour. A detailed experimental procedure and full calculations for CO2 solubility measurements using the pressure drop method can be seen in [76, 77].Further investigations that utilize this pressure drop method to derive gas solubility can be found where alternative experimental setups are shown [78–80].


## 3.3. View-Cell Methods

These involve the preparation of a mixture with a precisely known composition and then the observation of phase behavior inside an equilibrium cell, where measurements are recorded in the equilibrium state, that is, temperature and pressure. Synthetic methods consist of two main techniques, one being with a phase transition, and the other without. In synthetic methods with a phase transition a known amount of gas and IL is loaded into the equilibrium cell. The pressure is then varied at a constant temperature (or vice versa) until a second phase is formed, where the gas dissolves in the ionic liquid causing the vapor phase to diminish, whereby using different gas pressures, solubility can be worked out at various pressure, and temperatures.In synthetic methods without a phase transition, equilibrium properties like temperature, pressure, density, cell volume, and gas/liquid phase volumes are measured, and the composition of the phase mixtures can be calculated in terms of moles or by a mass balance equation.As can be seen in Figure5, a pump releases CO2 at a constant selected pressure and monitors the volume of CO2 flowing into the system. The CO2 is also heated to a constant temperature. By monitoring the volume, a known amount of CO2 is then introduced to the high-pressure view cell, which contains a known amount of IL. In the case of non-phase transition, the amount of CO2 absorbed is calculated by the difference in the amount of gas delivered to the cell and the amount of gas in the vapor phase. The amount of gas in the vapor phase can be calculated using a mass balance, shown in (2), coupled with an equation of state.Equation to calculate amount of gas in the vapor phase:(2)mg=mpump-mlines-mheadspace+mlines0+mheadspace0,where mg is the mass of CO2 in the liquid phase, mpump is the mass of CO2 injected into the system, mlines is the mass of CO2 in the gas lines, connecting the pump to the equilibrium cell, mheadspace is the mass of the gas in the headspace of the cell, mlines0 is the mass of the gas in the lines after venting the system, and mheadspace0 is the mass of gas in the headspace, initially in the system after venting.Figure 5
Scheme of a synthetic method setup.A full experimental procedure using a synthetic method without phase transition and demonstrating the use of mass balancing to determine gas solubility is explained in the literature [30, 81, 82].


## 3.4. Further Techniques

Gas chromatography is an analytical method that boosts high precision and accuracy. When applied to measuring gas solubility in absorption media, the gas chromatograph is usually coupled with a high-pressure reactor cell in which a synthetic or pressure drop method is applied, and at equilibrium, a sample is taken and analyzed [83]. Solubility data from gas chromatography can be achieved by using an extractive technique; here the solvent (IL) is saturated with the solute (CO2) and then coated on a column. Nitrogen, or any other nonabsorbing carrier gas, is directed on to the column in order to extract the CO2 from the IL. The nitrogen is then analyzed in the gas chromatograph. This determines the amount of CO2 removed (per amount of coating). A detailed method for applying gas chromatography can be found in the thesis by Wilbanks [84].Other analytical techniques can be used in some cases to determine the solubility of specific gases; this may be in the form of a titration; this was demonstrated by Shen and Li [85] with aqueous amine solutions. However this has so far not been applied to ILs. Inline gas sensors also have the potential to be used. A possible scenario could involve linking an electrochemical sensor to measure the difference in CO2 concentration of the vapor phase before and after equilibrium conditions.Many advantages come from using gravimetric microbalances for solubility measurements. The ability to measure mass change to a high precision is helpful for a variety of reasons. When initially degassing the ionic liquid sample, being able to measure mass decrease allows the experimentalist to see when a constant mass value has been reached, thus allowing the assumption that full degassing has occurred. Also, the mass reading is very important to ensure equilibrium conditions once the CO2 has been introduced. Equilibrium is reached when the mass change is zero. In situ gravimetric balances, that is, when the balance is enclosed in the measuring gas, are limited to lower pressures and temperatures. Disadvantages of gravimetric systems are mainly due to their high retail price, making them impractical for small projects where the funding is restricted.In comparison to the gravimetric analysis, the pressure drop and synthetic methods are much simpler in design. As samples of any size can be investigated with these methods, a high sensitivity can generally be achieved, however, not a high accuracy. The most significant errors in the pressure drop and synthetic methods are the error calculations of dead space; for gravimetric methods, it is the determination of buoyancy forces. In pressure drop and synthetic methods, the two variables, pressure and gas absorbed, are determined by the pressure sensors and calibrated volumes; this can result in measuring error which is added on each step of the absorption isotherm. With the gravimetric method, all of the variables, temperature, gas pressure, and absorbed gas, are measured independently and the absorption pressure is monitored at each step of the isotherm.


## 4. Conclusion

The versatility and inherent advantages of ionic liquids in the process of CO2 capture are giving rise to a promising and expansive field. Their potential as physical absorbents is highly attractive, although at present their capture rate is not to the same scale as current aqueous amine technologies; the fact that amines for CO2 capture have been developed through many years and that ILs are a new research field leaves room for further research and improvement.Solubility data of CO2 in different imidazolium-based ionic liquids are the most often found in the literature. This is especially the case for bmim[BF4] and bmim[PF6], because these ionic liquids were among the first ones commercially available. Therefore an abundant amount of previous data is available and allows the validation of subsequent experimental procedures. Although commercially available, the price of these ionic liquids remains high. Quaternary ammonium and tetra alkylphosphonium bases provide a cheaper alternative. In comparison the synthesis process of these ionic liquids is simpler and the raw materials are accessible. However the lack of experimental data with these solvents means that they are constantly overshadowed.Although experimental data on CO2 solubility in ionic liquids is available in the literature, more is needed for process design. Here several different methods have been presented in order to obtain this data. These include gravimetric analysis, pressure drop, and synthetic methods, all of which are particularly well suited for the measurement of gases in nonvolatile liquids. In terms of solubility data measurements, gravimetric balances offer the simplest and most precise route; however, their general high prices make them impractical for small research groups conducting initial experiments with ionic liquids. Pressure drop and synthetic methods provide a cheaper alternative and do not need sampling. However these methods depend on the models used to calculate the thermodynamic properties and phase equilibrium. It is important to observe that for some thermodynamic properties, such as excess molar enthalpy, research groups use a test system to check their equipment and methods accuracy. In the case of gas-liquid solubility, however, there is no test system, especially at elevated temperatures and pressures.The main challenges affecting ionic liquids as green solvent for CO2 capture are availability, cost, purity, and compatibility. These challenges are faced at present on a laboratory scale and must have solutions before expanding to industry. At present the advantages and disadvantages of ionic liquids and amines seem to be equally balanced. The main criteria for ideal CO2 capture mechanisms are high CO2 solubility, low energy input for regeneration, low cost, long-term reusability, and being environmentally benign. At the moment amines have the advantage of having high CO2 solubility and being of low cost. However due to the vast number of ionic liquids that can be developed and different ways in which they can be synthesized, the potential is there. Moreover through increasing research and commercialization of ionic liquids in other areas of industry, the cost is set to decrease.



---

*Source: 473584-2013-06-15.xml*