Natural Gas Sweetening with Ionic Liquids A Selectivity ...

Natural Gas Sweetening with Ionic LiquidsA Selectivity Analysis

MSc Thesis by A. AmplianitisFebruary 2014

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Natural Gas Sweetening with Ionic Liquids

A Selectivity Analysis

by

A. Amplianitis

In partial fulfillment of the requirements for the degree of

Master of Science in Sustainable Energy Technology (SET)

at the Delft University of Technology, to be defended publicly on Thursday February 13, 2014 at 14:00.

Supervisor: Prof.dr.ir. Thijs Vlugt

Thesis committee:

Dr.ir. Theo de Loos

Dr.ir. Klaas Besseling

Ir. Mahinder Ramdin

Report number: 2607

An electronic version of this thesis is available at http://repository.tudelft.nl/.

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Abstract

Global natural gas consumption has grown significantly over the years rendering natural gas one of the

most important energy sources of the future. Although the known natural gas resources are significant, half of

the known gas fields are estimated to contain more than 2% CO2 rendering them sub-quality reserves. Apart

from carbon dioxide, natural gas also contains other sour gasses like hydrogen sulfide (H2S). These impurities

need to be removed from natural gas because in the presence of water, acids that can corrode pipelines and

other processing equipment are formed. Furthermore, as CO2 provides no heating value it has to be removed

in order to meet gas quality specifications before distribution to customers. Gas processing to remove acid

gasses is referred to as natural gas sweetening.

Among the materials investigated as new potential solvents for CO2 absorption processes, ionic liquids

are one category of solvents that may in the future offer an alternative to amines and low capacity physical

solvents. CO2 solubility and selectivity data in ILs are essential if we are to judge the separation performance

of these solvents.

The goal of the present thesis was to experimentally determine the solubility of methane in two

phosphonium based ionic liquids for which CO2 solubility data were already available. In that way the ideal

selectivity of CO2/CH4 in those ILs was determined. Furthermore, mixed gas solubilities were experimentally

determined in ternary systems of IL+CO2+CH4. As the acquired experimental data are not enough in order to

calculate real selectivities, the present work focuses on comparing the mixed gas bubble-point pressures to the

sum of the single gas bubble-point pressures acquired from measurements in binary systems. The ultimate

goal is to determine whether or not there is significant interaction between the two gases that can eventually

result in real selectivity deviating from ideal selectivity. Finally, solubility of CO2 in an imidazolium-based IL

is experimentally determined.

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Acknowledgements

First of all I would like to thank Prof.Thijs Vlugt as well as the entire Engineering Thermodynamics group

for giving me the opportunity to be part of a highly skilled team and work in the forefront of chemical

engineering research.

I would like to give my special thanks to Prof. Theo de Loos for his indispensable guidance and help week

after week on all the theoretical and technical issues that arose during my time in the thermolab.

I would like to express my outmost gratitude to Mahinder Ramdin. Without his guidance and uninterrupted

help, this dissertation would not have been possible.

My experimental work in the Engineering Thermodynamics laboratory would not have been possible

without the valuable help and continuous guidance of Eugene Straver. Furthermore I want to thank Mariette

de Groen and María Teresa Mota Martínez for their support and presence in the thermolab.

I would like to thank my dear friends Georgiana, Myrto, Lina, Despoina, Spyros, Christina and

Constantinos for their continuous moral support.

Last and most importantly, I would like to thank my parents who enabled my studies abroad by supporting

me both morally and financially.

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Table of contents

ABSTRACT IV!

ACKNOWLEDGEMENTS VI!

1. INTRODUCTION 1!

2.1 NATURAL GAS SWEETENING 4!

2.1.1 LIQUID-PHASE PROCESSES 4!

2.1.2 DRY-BED PROCESSES 7!

2.1.3 MEMBRANES 8!

2.1.4 CRYOGENIC FRACTIONATION 8!

2.2 IONIC LIQUIDS 9!

2.3 CO2 /CH4 SELECTIVITY 10!

3. EXPERIMENTAL PROCEDURE 12!

3.1 PHASE BEHAVIOR THEORY 12!

3.2 EXPERIMENTAL METHOD 13!

3.3 MEASUREMENTS AND CORRECTIONS 16!

3.2 IONIC LIQUIDS USED 17!

4. MODELING THE PHASE BEHAVIOR OF BINARY SYSTEMS 18!

4.1 PENG-ROBINSON EQUATION OF STATE 18!

4.2 VAPOR-LIQUID EQUILIBRIA (VLE) 20!

4.3 MATLAB 22!

5. RESULTS 23!

5.1 BINARY SYSTEMS 23!

5.1.1 THTDP-DICYANAMIDE 23!

5.1.2 THTDP-PHOSPHINATE 27!

5.1.3 EMIM-PHOSPHATE 31!

5.2 TERNARY SYSTEMS 33!

TERNARY SYSTEMS SUMMARY 39!

6. CONCLUSION AND FUTURE DIRECTION 40!

6.1 CONCLUSION 40!

6.2 FUTURE DIRECTIONS 42!

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BIBLIOGRAPHY 43!

APPENDIX A 46!

CALCULATION OF THE SECOND VIRIAL COEFFICIENTS OF PURE GASSES AND GAS MIXTURES. 46!

APPENDIX B 48!

EXPERIMENTAL MEASUREMENTS 48!

1. Introduction

Global natural gas consumption has grown significantly over the years rendering natural gas as

one of the most important energy sources of the future.1 It is thought to be the most environmentally

friendly fossil fuel, since burning it leads to negligible SO2 emissions, low NOx levels and less than

half of the CO2 emitted when burning coal or oil.1 In 2010, natural gas supplied 23.81% of the world’s

energy demand and the volume of the consumed natural gas increased by 7.4% over 2009 levels.3 The

volumes of natural gas traded as pipeline and liquefied natural gas (LNG) in 2012 is shown in Table1.

Although the known natural gas resources are significant, half of the known gas fields are

estimated to contain more than 2% CO2 rendering them sub-quality reserves.2,4 The increased demand

for natural gas has inevitably led to a re-evaluation of the development potential of those sub-quality

gas reserves that had been previously considered economically unviable. Apart from carbon dioxide,

natural gas also contains other sour gasses like hydrogen sulfide (H2S). These impurities need to be

removed from natural gas because in the presence of water, acids are formed that can corrode

pipelines and other processing equipment. 2 Furthermore, as CO2 provides no heating value it has to

be removed in order to meet gas quality specifications before distribution to customers. Those

specifications usually stipulate a maximum level of 3% by volume CO2 in natural gas transmitted to

customers by pipeline.2,5

As most of the significant natural gas resources are located far from the large and established

natural gas markets, significant volumes of natural gas need to be transported as LNG.2 The

specifications of CO2 removal from natural gas to be processed in a cryogenic plant to produce LNG

are far more stringent than those for typical gas pipelines. Carbon dioxide concentration needs to be

less than 50 ppm before the sweetened gas enters the cryogenic processes within the plant in order to

avoid the formation of dry ice.2,5

Table 1. Volumes of natural gas traded as pipeline and LNG (billions of cubic meters) by top natural gas exporting and importing countries in 2012.2,3

Top natural gas exporters Top natural gas importers

Pipeline LNG Total Pipeline LNG Total

1 2 3 4 5 6 7 8 9 10

Russian Federation Norway Qatar Canada Algeria Netherlands Indonesia Malaysia U.S. Australia

186.5 95.9 19.2 92.4 36.5 53.3 9.9 1.5 30.3 0.00

13.4 4.7 75.8 0.0 19.3 0.0 31.4 30.5 1.6 25.4

199.9 100.6 94.9 92.4 55.8 53.3 41.3 32.0 32.0 25.4

U.S. Japan Germany Italy U.K. France South Korea Turkey Spain Ukraine

93.3 0.0 92.8 66.3 35.0 35.0 0.0 28.8 8.9 33.0

12.2 93.5 0.0 9.1 18.7 13.9 44.4 7.9 27.5 0.0

105.5 93.5 92.8 75.3 53.6 48.9 44.4 36.7 36.4 33.0

2

The development therefore of the sub-quality natural gas reserves together with the increased

LNG production and commercialization have presented new challenges to gas processing that require

more efficient approaches to the conventional absorption and separation technologies that are most

commonly used for CO2 removal from natural gas. Various technologies have been practiced both on

laboratory and industrial scale for natural gas sweetening. These processes involve

chemisorption/physisorption, membrane separation or molecular sieves, amine physical absorption,

carbamation, amine dry scrubbing and mineral carbonation.6 Currently, the most vastly used methods

for natural gas sweetening purposes are amine-based processes. Alkanolamines such as

monoethanolamine (MEA), diethanolamine (DEA), and methyldiethanolamine (MDEA) based

methods are vastly used and account for over 95% of all gas sweetening in the United States.6,7 There

are however certain disadvantages for the commercial use of these solutions. Main issues are the loss

of amine reagents, transfer of water into the gas stream during the desorption stage, degradation of the

amine reagent to form corrosive byproducts and high-energy consumption during regeneration.6 The

regeneration step may account for the 70% of the total operating costs of the capture process.6 Issues

like that have to be overcome with new, more efficient and less costly methods.

Among the materials investigated as new potential solvents for CO2 absorption processes, ionic

liquids (ILs) are one category of solvents that may in the future offer an alternative to amines and low

capacity physical solvents. ILs are commonly defined as organic salts with melting temperatures of

less than 373K2. Properties of ILs such as extremely low vapor pressure, high chemical/thermal

stability, no flammability and in certain cases zero toxicity may render them useful replacements for

volatile organic solvents.2,6 Furthermore the physical and chemical properties of ILs can be enhanced

and modified by altering both their cationic and anionic moieties.7 In that way, IL properties can be

tailor-designed to meet the specific needs for natural gas sweetening processes: high solubility of

CO2, increased CO2 selectivity, high thermal stability and low toxicity.

The fact that CO2 is inherently soluble in many ionic liquids has prompted many researchers to

explore the vast synthetic landscape provided by ILs. Initially most of the studies focused on ILs with

imidazolium-based cations because of their observed affinity towards CO2.6,7 Several of those ILs

have been studied with regard to their CO2 solubility and in order to understand phase behavior of

CO2-IL pairs.6,8 Combinations of different anions and cations however, give a huge number of

possible ILs that need to be studied. Furthermore, as separation processes involve mixtures of two or

more components that need to be separated, apart from solubility data, selectivity data are also

essential if we are to judge the separation performance of a solvent.

Although there is great availability of CO2 solubility data in the literature, selectivity data of

CO2 in ILs are scarse.7 Furthermore these selectivity data mainly refer to ideal selectivities (i.e. the

ratio of pure gas solubilities) and not real selectivities. Real or actual selectivities cannot always be

accurately determined from pure gas solubilities with the assumption of ideal mixing.7 However since

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measuring mixed-gas solubilities is significantly more difficult there are almost no mixture data

available in the literature. The selectivities that are mainly relevant for CO2 capture from natural gas

are CO2/CH4 and CO2/H2S.6,7

With all the above in mind, the goal of the present thesis was to experimentally determine the

solubility of methane in two phosphonium based ILs for which CO2 solubility data were already

available. In that way the ideal selectivity of CO2/CH4 in those ILs was determined. Furthermore

mixed gas solubilities were experimentally determined in ternary systems of IL+CO2+CH4. As there

is no single straightforward way to calculate real selectivities, the present work focused on comparing

the mixed gas bubble-point pressures to the sum of the single gas bubble point pressures acquired

from measurements in binary systems. The ultimate goal was to determine whether or not there is

significant interaction between the two gases that would eventually result in real selectivity deviating

from ideal selectivity. Finally solubility of CO2 in an imidazolium-based IL was experimentally

determined.

The experimental work involved in the present thesis can be summarized in the next points:

• Experimental study of the gas-liquid equilibrium of ionic liquid

1-Ethyl-3-methylimidazolium diethyl phosphate ([emim]-[phosphate]) with CO2.

• Experimental study of the gas-liquid equilibrium of the phosphonium-based ionic liquids

trihexyl-tetradecyl-phosphonium-dicyanamide([thtdp][dca]) and trihexyl-tetradecyl-

phosphonium-bis(2,4,4-trimethyl-pentyl)-phosphinate ([thtdp][phosphinate]) with

methane and subsequent modeling of the two binary systems with the Peng-Robinson

EoS.

• Experimental study of the gas-liquid equilibrium of the ternary systems of the two

phosphonium-based ionic liquids, mentioned above, with gas mixtures of CO2 and CH4.

The content of this thesis is divided in five chapters. After this brief introduction, a detailed

literature review of the various gas separation technologies currently employed is presented in

Chapter 2. In Chapter 3 the experimental procedure and set-up is presented. Chapter 4 is dedicated to

binary systems modeling using Equations of State. Chapter 5 summarizes all the experimental and

modeling results for binary and ternary systems. Finally Chapter 6 outlines the conclusions that were

drawn and presents future research directions.

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2. Literature study

2.1 Natural Gas Sweetening

As discussed previously, raw natural gas contains large amounts of acid gases such as carbon

dioxide (CO2) and hydrogen sulfide (H2S). Those gasses are referred to as acid gasses, because they

dissolve in water to form weak acids.2,6 Furthermore, natural gas can also contain other contaminants

such as carbonyl sulfide (COS), mercaptans (R-SH) and carbon disulfide (CS2).9 Many issues arise

from the presence of these contaminants. Acids formed by the dissolution of acid gasses to water can

corrode pipelines. Combustion of sulfur-based contaminants produces serious pollutants responsible

for acid rain. Furthermore, those sulfur compounds are inherently poisonous to humans and animals

and are corrosive to metals and other material used for the transportation and handling of natural

gas.2,9 Another important issue is that since CO2 is non-flammable, its presence reduces the heating

value and therefore the sale value of natural gas.2 It is therefore clear that the removal of acid gasses

and sulfur compounds is a major issue in natural gas processing.

The major challenge in natural gas sweetening processes is to remove all the contaminants to as

low a level as possible. Nowadays, there are several methods and processes deployed to sweeten

natural gas. However, since the concentration of CO2 and H2S in the raw natural gas to be processed,

as well as the permitted acid gas levels in the final product may vary significantly, no single process is

deemed markedly superior for all circumstances.2,9 Over the next pages, an overview of the most

commonly used processes for natural gas sweetening will be presented.

2.1.1 Liquid-phase processes

Liquid-phase processes can be classified into three categories based on the nature of the liquid

solvent that is being used. Namely there are chemical solvents, physical solvents and hybrid solvents

that are mixtures of amines together with physical solvents.

Chemical Solvent Processes

In sweetening processes that use chemical solvents, absorption of the acid gasses is achieved by

using alkanolamines or alkaline salts of various weak acids such as sodium and potassium salts of

carbonate. Regeneration of the solvent is achieved by reducing the pressure or by application of high

temperatures, which results in the desorption of the acid gases from the solvent. These processes are

suitable for removing H2S and CO2 from raw gas, but will still not remove organic sulfur components

such as carbon disulfide.9 Furthermore chemical solvents are specifically suitable when contaminants

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at low partial pressures need to be removed to very low concentrations. Their use is also advantageous

because of the low co-absorption of hydrocarbons.2

Amine Processes

Chemical absorption processes using aqueous amine solutions is a mature technology used in

the ammonia process, steam reforming and natural gas sweetening processes.7 Amines are compounds

derived from ammonia (NH3) by replacing one or more hydrogen atoms with another hydrocarbon

group, usually an alkyl or aromatic group.2 Depending on the degree of substitution of hydrogen

atoms by organic groups, amines are categorized as primary, secondary or tertiary. Replacement of a

single hydrogen atom results in a primary amine such as monoethanolamine (MEA) or diglycolamine

(DGA). Replacement of two hydrogen atoms results in a secondary amine such as diethanolamine

(DEA) or diisopropylamine (DIPA). Finally replacement of all three hydrogen atoms gives tertiary

amines such as methyl-diethanolamine (MDEA).9 The theoretical maximum CO2 loading of primary

and secondary amines is limited to approximately 0.5 mole of CO2 per mole of amine. In the case of

tertiary amines, the chemical mechanism is different allowing a theoretical maximum of 1 mole of

CO2 per mole of amine. Furthermore tertiary amines require a lower heat of regeneration.2,7,9

Regardless of the aqueous amine solution used as the sweetening agent, the general process flow

diagram for an amine sweetening plant does not change much. The sour gas will nearly always enter

the plant through a scrubber to remove any free liquids and entrained solids. It enters then the

absorber through its bottom and flows upward in intimate countercurrent contact with the amine

solution that absorbs acid gasses. Once at the top, the sweetened gas exits the absorber and passes

through an outlet separator in order for entrained amines to be recovered. Due to the presence of water

in the amine solution, the sweetened gas is now saturated with water so a dehydration step is

necessary before the gas is sold or fed to a cryogenic plant for LNG production. The rich amine

solution leaves the bottom of the absorber containing 0.20-0.81 mole of acid gas per mole of

amine.9,11 Before entering the top of the stripper, in order to be regenerated, the amine solution goes

through a flash tank to recover hydrocarbons that might have dissolved in it. In the stripper the amine

comes to contact with steam, which results in the amine-CO2 separation. A stream of lean amine is

removed from the bottom of the stripper and once cooled, returns to the absorber. Acid gasses leave

the stripper from the top and can then be vented, incinerated, or compressed for re-injection into a

suitable reservoir for enhanced oil recovery.12

Major disadvantages of the conventional amine gas treating processes are:

• Large amounts of energy required for the amine regeneration

• Relatively low CO2 loading capacity of amines which requires high solvent circulation rates

and large high-pressure absorber columns

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• Corrosive nature of amine solutions induce high wear of the processing equipment

• Degradation of the amines to organic acids

• Co-absorption and subsequent loss of hydrocarbons

• Volatility of the amine solvents causes environmental pollution

Potassium Carbonate Process

The potassium carbonate process is another chemical solvent process for treating gas streams. It

employs an aqueous solution of potassium carbonate (K2CO3) to remove CO2 and H2S. This process

requires high partial pressure of CO2 and usually a two-stage process is needed to sufficiently remove

acid gasses to acceptable low levels.9 Besides that though, the process flow steps are very similar to

those of an amine process. It is a particularly useful process for removing large quantities of CO2.

Advantages of using carbonate solutions are the high chemical solubility of CO2 in the

carbonate/bicarbonate system and the low solvent cost. Main disadvantage is the fact that potassium

carbonate is highly corrosive.

Physical Solvent Processes

In these processes, organic solvents are employed and acid gasses are removed from gas

streams by physical absorption instead of chemical reaction. The high solubility of acid gasses in

organic solvents (at high pressures) is the driving force for these processes. Solubility usually

increases as the temperature decreases and as pressure increases. Regeneration of the solvent is later

achieved either by heating or by pressure reduction. Compared to chemical solvents, the heat required

to regenerate physical solvents is much less, since the heat of absorption of acid gasses for physical

solvents is significantly lower. Furthermore, as most of the physical solvents absorb water, the

required capacity for dehydration of the processed gas is much less compared to aqueous amine

processes that saturate the sweet gas with water. Main weakness of physical solvents remain the issue

of relatively low acid gas absorption capacities (at low pressures) which makes physical solvent

processes competitive with amine processes only when the feed gas is available at high pressures.13

An advantage of the use of physical solvents however, is their less corrosive nature, compared to

aqueous amine solutions, whose handling requires more expensive and corrosion resistant materials.

An ideal physical solvent should have a high selectivity for acid gasses, very low vapor

pressure to minimize solvents losses, low viscosity, high thermal stability and should not be of

corrosive nature. Although many different compounds have been commercially used as physical

solvents for gas treating, none of them is markedly superior for all circumstances.13 Selection of the

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right solvent comes down to feed gas composition, process objectives and the special characteristics

of the solvent.

Physical solvents that have been commercially used for treating natural gas are: dimethyl-ether

of polyethylene glycol (DEPG), propylene carbonate (PC), N-methyl-2-pyrrolidone (NMP) and

methanol (MeOH).13 DEPG is used in the SELEXOL® process licensed by Dow. It is a process that

can remove CO2 simultaneously with H2S and water. SELEXOL® together with the RECTISOL®

process, form the leading physical absorption technologies for treating feed gas with very high CO2

concentrations.2 RECTISOL® is a process licensed by Lurgi GmbH that employs methanol as the

physical solvent and is used to remove H2S, COS and bulk CO2. NMP is employed in the PURISOL®

process (also licensed by Lurgi GmbH) to simultaneously remove H2S, CO2, RSH and H2O although

this particular solvent, is highly selective for H2S. Finally, propylene carbonate (PC) is used in the

Fluor process (licensed by Fluor Daniel Inc.) to remove CO2, H2S, COS, CS2 and H2O from natural

gas, achieving sweetening and dehydration of natural gas in one step.

Hybrid Solvents

Hybrid solvent processes employ mixtures of amines and physical solvents in order to take

advantage of the best characteristics of both. Depending on the combination of the physical solvent

and the amine that is being used, nearly complete removal of H2S, CO2, and organic sulfur

compounds is possible. Other advantages are higher acid gas loading, lower energy requirements for

regeneration, lower corrosion rates, and lower foaming tendency. The most widely known hybrid

solvent process is the Shell SULFINOL® process, which applies a mixture of sulfolane, water, and

diisopropanolamine (DIPA) or methyldiethanolamine (MDEA), Sulfinol-D, and Sulfinol-M,

respectively. This process is used to selectively remove H2S, COS, RSH, and other organic sulfur

compounds for pipeline specifications, while co-absorbing only part of the CO2.14–16

2.1.2 Dry-bed processes

Dry-bed processes use a fixed bed of solid material to remove acid gases either through ionic

bonding (physical adsorption) or through chemical reactions. When the bed is saturated with acid

gases, it needs to be regenerated or replaced. In the case of physical adsorption, regeneration of the

adsorbent is achieved by one or more simple temperature or pressure swing cycles. When the

adsorbed component reacts chemically with the bed material, the process is called chemisorption and

desorption is generally not possible. In this case the bed material needs to be replaced.

Using porous solid adsorbents for purification of gas mixtures is a process technology used in

the production of hydrogen, separation of oxygen and nitrogen from air streams and the capture of

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odorous pollutants from industrial processes. For natural gas treating, adsorption-based separation is

used to remove water, sulfur compounds, mercury and heavy hydrocarbons.17

2.1.3 Membranes

Although membrane separation technologies have been applied in the natural gas industry to

remove CO2, N2 and H2S since 1980, they still account for less than 5% of the market for new natural

gas processing equipment installed.18 Various membrane technologies have been deployed in the

natural gas industry. Of those technologies however, processes to capture CO2 have been the most

widely used and currently CO2 capture is the only natural gas separation process for which membrane

processes are competitive with the conventional amine technology.18 Membrane systems for natural gas processing employ polymeric membranes with the industry

standard currently being cellulose acetate.13,18 These membranes are of the solution-diffusion type, in

which a thin layer of cellulose acetate lays on top of a thicker layer of a porous support material that

provides the required mechanical strength.18 Gas separation is achieved by selective permeation of the

gas constituents in contact with the membrane. High-pressure gas is fed to one side of the membrane

while the other side is maintained at much lower pressure. After the gases have dissolved in the

membrane material, they move across the membrane barrier under the imposed partial pressure

gradient. The advantages of membrane technologies which makes them highly attractive process

separation units are the ability to separate chemical species without a phase change, low thermal

energy requirements, simple process flow schemes with few pieces of rotating equipment, compact

plant footprints and convenient start up and shutdown procedures. These features of membranes

systems are potentially attractive for remote, unmanned and footprint conscious sites. The main

limitation of their use is the significant loss of hydrocarbons in the effluent stream. This issue

however, can be handled by a hydrocarbon recovery stage downstream the membrane separation

stage. Furthermore, while membrane systems perform well at reduced feed flow rates, their

performance drops when design flow rates are exceeded. To tackle that, additional membrane

modules must be added in parallel which increases the overall cost.13

2.1.4 Cryogenic Fractionation

Cryogenic fractionation involves cooling the gases to a very low temperature so that the CO2

can be liquefied and separated. This technology requires substantial energy to achieve the low

temperatures needed. Furthermore, issues arise by the formation of CO2 solids during cryogenic

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distillation. To overcome those problems, two technological approaches have been pursued: (1)

extractive distillation by the addition of a heavier hydrocarbon to alter the solubility of components in

the column (Ryan/Holmes process) and (2) controlled freezing and re-melting of the solids

(Controlled Freeze ZoneTM® and CryoCells® processes).2 Main advantage of this method over

amine-based absorption processes, is the fact that the acid gas components are discharged as a high-

pressure liquid stream that can be easily pumped for geo-sequestration or for use in enhanced oil

recovery operations, while yielding a high-quality methane product.19 Furthermore, highly corrosive

aqueous amine solvents are avoided and process footprint is reduced, which may be important

consideration for offshore or floating production facilities.20

2.2 Ionic Liquids

Room-temperature ionic liquids (RTILs) are organic salts that melt below 100 °C and have

many interesting properties that make them excellent candidates for use in a variety of applications.

RTILs have negligible vapor pressures; are thermally and chemically stable and non-flammable.6,7

The negligible vapor pressures of RTILs along with their desirable gas solubility, enables them to be

used for various gas separation applications. CO2 capture from flue gas as well as for natural gas

sweetening purposes is a process field where the use of ILs as potential solvents has attracted the

interest of many researchers over the last decade. Primarily CO2 solubility in ILs has been the object

of most of the relevant research. The fact that the solubility and the selectivity of CO2 in RTILs can be

readily “tuned” by tailoring the structures of the cation and/or anion has given immense depth in the

research possibilities. Figure 1 shows the most commonly used anions and cations for IL synthesis.

Figure 1. Commonly used anions and cations for IL synthesis.7

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The solubility of CO2 in different ILs has been widely studied by several researchers.8,21–24

However, for gas separation processes, apart from solubility, gas selectivity is also a major factor that

dictates the choice of the most suitable solvent. Although dozens of CO2 solubility data are available

in the literature, selectivity data of CO2 in ILs are scarcely reported. In the case of natural gas

sweetening processes solubility of hydrocarbons like methane and ethane in ILs should also be

considered. Nevertheless research on methane solubilities in ILs is far less extensive than the

available research on CO2 solubilities.

2.3 CO2 /CH4 selectivity

In natural gas sweetening processes, the following selectivities are mainly relevant: CO2/N2,

CO2/H2, CO2/CH4. In this part however, only work relevant to CO2/CH4 selectivity is mentioned with

the majority of the research focusing mainly on ideal selectivities (i.e., the ratio of pure gas

solubilities). Determining real or actual selectivities in mixtures of gases cannot always be accurately

done from pure gas solubilities with the assumption of ideal mixing. Measuring mixed-gas solubilities

is significantly more difficult, therefore almost no mixture data can be found in the literature.

Although ideal selectivity of CO2/CH4 in ILs is comparable to that in conventional physical solvents,

real CO2/CH4 selectivity in ILs is expected to be lower than the ideal selectivity, since CH4 solubility

in many ILs increases for increasing temperature while this is the opposite for CO2.7

In their work, Anderson et al.25 measured the solubility of methane, among other gasses, in the

IL [hmpy][Tf2N]. The solubility of these gases (298 K) decreased by the following order: SO2 > CO2

> C2H4 > C2H6 > CH4 > O2 > N2. Similar gas solubility trends (i.e., CO2 > C2H4 > C2H6 > CH4 > O2)

were observed in the ILs [hmim][Tf2N], [bmim][PF6], and [bmim][Tf2N].25–27 In most of the ILs

studied, N2 and O2 solubilities are much lower compared to CO2 therefore the relevant selectivities are

high enough. Hydrocarbons however, such as methane, show moderate solubilities in ILs, thereby

reducing the CO2/hydrocarbon selectivity.

Applying regular solution theory, Camper et al. 28–33showed that the physical solubility of gases

in ILs was well correlated with the liquid molar volume of the IL and that ideal selectivities for

CO2/N2 and CO2/CH4 should increase as the molar volume of the IL decrease.

Finotello et al.33measured CO2, N2, and CH4 solubilities in [emim][Tf2N],[emim][BF4],

[hmim][Tf2N] and [mmim][MeSO4]. Their results show that, as temperature increases, the solubility

of CO2 decreases in all RTILs, the solubility of CH4 remains constant in [emim][Tf2N] and

[hmim][Tf2N] but increases in [mmim][MeSO4] and [emim][BF4]. Also, CO2/CH4 ideal solubility

selectivity increases as temperature decreases.

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Kumelan et al.34measured the solubility of methane and of xenon in the ionic liquid 1-n-butyl-

3-methylimidazolium methyl sulfate ([bmim][CH3SO4]) with a high-pressure view-cell technique

based on the synthetic method. Among his findings is that CH4 becomes less soluble in

[bmim][CH3SO4] with rising temperature.

Jacquemin et al. 35 reported experimental values for the solubility of carbon dioxide, ethane,

methane, oxygen, nitrogen, hydrogen, argon and carbon monoxide in 1-butyl-3-methylimidazolium

tetrafluoroborate, [bmim][BF4]. They conclude that carbon dioxide is the most soluble gas with mole

fraction solubilities of the order of 10-2. Ethane and methane are one order of magnitude more soluble

than the other five gases that have mole fraction solubilities of the order of 10-4.

In their work, Carvalho and Coutinho36 measured CH4 solubilities in imidazolium,

phosphonium, and ammonium ILs. Among the findings was the fact that an increase in temperature

had a small or even negligible impact on CH4 solubility. Furthermore, the solubility of CH4 was

shown to relate to the polarity of the IL.

Althulith et al.37present experimental measurements of the CH4 solubility in [emim][FAP].

Which is then compared to solubilities in other ILs. This comparison shows that the solubility of CH4

in the various ILs slightly decreases in the order: [hmim][Tf2N] > [emim][FAP] > [bmim][Tf2N].

Furthermore solubilities of CH4 are much lower compared with CO2 solubilities, thus making the

relevant ILs suitable for separating CO2 from natural gas.

Bara et al.38 measured the solubility and ideal selectivities of the gas pairs CO2/N2 and CO2/CH4

in imidazolium-based ILs functionalized with oligo(ethylene glycol). They showed that the CO2

solubility in these oligo(ethylene glycol) functionalized ILs were similar to their corresponding alkyl

analogues, but N2 and CH4 solubilities were lower corresponding to a higher ideal selectivity for the

two gas pairs. Similar results were reported by Carlisle et al.39 for nitrile-functionalized ILs. The

nitrile-functionalized ILs exhibited lower CO2, N2, and CH4 solubilities, but showed improved

CO2/N2 and CO2/CH4 selectivities compared to alkyl-substituted analogues.

12

3. Experimental Procedure

3.1 Phase Behavior Theory

According to Gibbs phase rule for non-reacting systems (eq. 3.1), the number of variables that

may be independently fixed in a system at equilibrium, is the difference between the total number of

variables that characterize the intensive state of the system and the number of independent equations

that can be written connecting those variables.40 This difference gives the degrees of freedom (F) of

the systems. For a system containing N chemical species and π phases, F is evaluated:

! = 2 − ! + ! (3.1)

A unary system (N = 1) must have at least one phase (π = 1) so the maximum degrees of freedom are

two. Therefore a single P-T diagram can describe phase behavior of pure substances. A typical P-T

diagram of a pure substance is shown in Figure 2. At the triple point, the unary system (N = 1), with

three phases (π = 3) reduces to a single point.40

Figure 2 P-T diagram of a pure substance

For a mixture of two components the degrees of freedom increase. A binary mixture (N = 2)

with at least one phase (π = 1) has a maximum of three degrees of freedom (F), namely: pressure,

temperature and mole fraction. These three dimensions can be projected in a three-dimensional space

P-T-x projection.

13

3.2 Experimental method

Experiments were carried out in a Cailletet apparatus, schematically shown in Figure 3. This

apparatus allows the measurement of phase equilibrium according to the synthetic method.41 In this

method, a mixture of fixed overall composition is injected into an equilibrium cell, which is then fully

sealed off throughout the whole experiment. That way, phase behavior is constrained to two degrees

of freedom, namely temperature and pressure. For each data point measured, the temperature of the

sample remains fixed at a set value and the pressure is slowly varied until phase change is observed

visually through the Cailletet transparent glass tube. Solubility experiments were executed on binary

systems of ionic liquids plus gas and pseudo-binary systems of ionic liquid plus a gas mixture of CO2

and CH4. In both cases bubble-point pressures were determined as the pressures at which the last

bubble of gas disappears in the liquid.

Figure 3. Schematic representation of the Cailletet apparatus: A, autoclave; B, magnets; C, capillary glass tube; D, drain; E, motor; F, metal stirrer; G, platinum resistance thermometer; H, rotating hand pump; Hg, mercury; I, thermostat liquid in; L, line to dead weight pressure gauge; M, mixture being investigated; Ma, manometers; O, thermostat liquid out; Or, hydraulic oil reservoir; P, closing plug; R, Viton-O-rings; S, silicone rubber stopper; T, mercury trap; Th, glass thermostat; V, valve.41

14

The ionic liquids used are highly hygroscopic. It is therefore essential that the amount of water

dissolved in the IL samples is as low as possible. In practice however, IL samples are considered

acceptable if water content is below the limit of 1200 ppm. To achieve that, the samples would have

to spend at least 24 hours at a temperature of 80°C in the vacuum oven. After that period, the exact

water content was measured using the Karl Fischer titration method. If the water content is within

acceptable limits, then the following steps are performed:

• An empty Cailletet tube (equilibrium cell) is weighed with an accuracy of 10-4 grams.

• Using a thoroughly washed and dried syringe, an arbitrary amount of ionic liquid is injected

in the tube, which is then weighed again to calculate the exact mass of the IL sample

(typically between 100 and 200 mg with an accuracy of 10-4 grams)

• Once the magnetic stirrer ball is inserted in the tube, the tube gets connected to a vessel of

calibrated volume at the gas rack. (Figure 4)

• The air and water in the sample are evacuated under high vacuum while the IL sample is kept

frozen by repeatedly being immersed in liquid nitrogen.

• The moles of gas in the calibrated vessel are calculated using the virial EoS. (Low pressure

ensures insignificant error by the use of the virial EoS40 )

• By adjusting the pressure of the gas or gas mixture in the calibrated volume vessel, the

desired mass percentage of gas in the final mixture is achieved. The gas rack used for filling

the Cailletet tube has an accuracy of ±0.1 mbar which leads to an uncertainty in the

composition of u(x)=±0.001.

Figure 4. Schematic representation of the gas rack.42

15

The gas partial pressure needed to achieve the desired gas mass percentage in the final sample

is calculated using the virial EoS (truncated to two terms for application at low pressures 40):

! = !!!!∗ − !(!)

(3.1)

Where: T : Temperature of the gas in the calibrated vessel R : Gas constant B(T) : Gas or gas-mixture second virial coefficient, which is only temperature dependent.40 (Detailed calculation of the second virial coefficients Appendix A) The molar volume of the gas or gas mixture is calculated:

!∗ = !!"##"$!

(3.2)

Where n are the moles of gas needed to achieve the desired mass fraction of gas in the sample.

Once the desired mass percentage of gas is achieved, the tube is sealed with mercury and then

transferred to the autoclave. The autoclave is connected to a hydraulic oil system in which pressure is

generated by means of a screw-type hand pump. Pressure measurements are done using a dead weight

pressure gauge with an accuracy of 0.05 bar.41 The temperature of the sample is kept constant at a set

value by means of a thermostat liquid, which is circulated through a glass jacket that surrounds the

Cailletet tube. A thermostat bath is used to maintain the temperature of the thermostat liquid to the

desired value with an accuracy of 0.01K 41. The temperature of the thermostat liquid is recorded with

an accuracy of 0.02 K using a platinum resistance thermometer. Finally proper agitation of the sample

is achieved using a stainless steel ball that is inserted in the tube along with the sample. The ball is

forced to move up and down by two reciprocating magnets.

16

3.3 Measurements and corrections

Once the Cailletet tube is connected to the autoclave, it gets covered with the glass jacket. The

jacket is then connected to the thermostat-bath and the flowing water within it, regulates the sample

temperature. After the sample reaches the desired temperature of the measurement, the pressure is

slowly varied using the dead-weight gauge. Bubble-point pressure for a specific temperature is then

determined as the pressure at which the last bubble of gas disappears.

For the measured bubble-point pressure to be correct, certain corrections need to be applied:

• Atmospheric pressure correction: The dead weight gauges always measure relative to the

atmospheric pressure. Therefore a barometer is employed with an accuracy of 0.1 mbar.

• Gravity correction for the laboratory latitude (52°):

!!,!"##$!%$& = !!"#$%&"' ∗9.812479.80665 = !!"#$%&"' ∗ 1.00059

(3.3)

• Temperature correction for the influence of the temperature !(°C)%on the piston area:

!!,!"##$!%$& = !!,!"##$!%$! ∗ [1 − ! − 20 ∗ 2.3 ∗ 10!!]

(3.4)

• Pressure transmitting fluid correction. In this case the pressure from the mercury column is

calculated using the temperature dependent density of mercury:

!!!" ! = −12.281 ∗ ! + 13595 (3.5)

!!" = !!" ! ∗ ! ∗ !ℎ

The actual experimental pressure is finally calculated: !!"#$% = !!,!"##$!%$& + !!"# − !!"

17

3.2 Ionic liquids used

Experiments were performed with the following ILs:

• 1-Ethyl-3-methylimidazolium diethyl phosphate ([emim]-[phosphate])

• Trihexyl-tetradecyl-phosphonium-dicyanamide([thtdp]-[dca])

• Trihexyl-tetradecyl-phosphonium-bis(2,4,4-trimethyl-pentyl)-phosphinate

([thtdp][phosphinate])

The ILs were purchased from SIGMA ALDRICH. In the table that follows are the information

provided by the company for these ILs. In Figures 5 and 6, the chemical structures of the ILs are

presented.

Thtdp-

dicyanamide

Thtdp-

phosphinate

Emim-

phosphate

Empirical Formula (Hill Notation) C34H68N3P C48H102O2P2 C10H21N2O4P

Molecular Weight (g/mole) 549.90 773.27 264.26

Density (g/mL at 20°C) 0.90 0.895 - Assay ≥ 95.0% ≥ 95.0% ≥ 98.0%

Table 2. Ionic liquid properties43

Figure 5. Chemical structures of [emim]-[phosphate] (left) and [thtdp]-[dicyanamide] (right).

Figure 6. Chemical Structure of [thtdp]-[phosphinate].

18

4. Modeling the phase behavior of binary systems

4.1 Peng-Robinson Equation of State

Using equations of state (EoS) is the most common method for phase equilibria correlation in

mixtures at high and low pressure. Cubic EoS, derived from the van der Waals EoS, are the most

common and industrially important EoS. Among them, the Peng–Robinson EoS has proven to

combine the simplicity and accuracy required for the prediction and correlation of volumetric and

thermodynamic properties of fluids. 44,45

The original Peng-Robinson EoS 46:

! = !!!! − ! −

! !! ! + ! + !(! − !)

(4.1)

Applying this equation at the critical point yields:

!(!!) = 0.45724!!!!!!!

!(!!) = 0.07780!!!!!

(4.2)

At temperatures other than the critical:

! ! = ! !! !!(!! ,!) ! ! = !(!!) (4.3)

Where !(!! ,!), is a dimensionless function of reduced temperature (!!) and acentric factor

(ω) that equals unity at the critical temperature.

! !!,! = 1 +! 1 − !!!.!!!

! = 0.37464 − 1.54226! − 0.26992!!

(

(4.4)

The extension of the PR EoS to mixtures, requires the use of mixing rules45:

!!"# = !!!!!!"!!

!!"# = !!!!!!"!!

(4.5)

Where

!!" = !!!! 1 − !!" = !!" !!" =12 (!! + !!)(1 − !!") !!! = !!! = 0 (4.6)

19

!! , !! !are the molar fractions of the components of the binary system and !!! , !!! are the binary

interaction parameters. To make the required calculations, numerical values of the critical temperature

and pressure (!! ,!!) as well as molecular weight and acentric factors (ω) are needed for CO2, CH4 and

the ILs

While these properties are well known for common substances such as carbon dioxide and

methane, they are not readily available for ILs. One reason for this is the fact that most of them start

to decompose at low temperatures and in some cases even at temperatures close to their boiling point.

As a result no experimental data exists and these properties must be estimated.

Group contribution methods correlate structural molecular properties with mathematical

functions representing a chemical property of a molecule. A group contribution method expresses the

thermodynamic property of a chemical compound as a function of the sum of contributions of smaller

groups of atoms constituting the molecule. A large variety of group contribution methods have been

proposed in the past years, differing in their field of applicability and in the set of experimental data

they are based on. Since the first group contribution methods were developed by Riedel in 1949 and

Lydersen in 1955, a large number of methods have been studied to achieve the most reliable results.

Valderrama et al.47 combined the best results of the Lydersen method with the best results of the

Joback-Reid method to propose a modified Lydersen-Joback-Reid method. This method gives good

results for molecules of high molecular weight such as ILs. In Table 3, critical properties and acentric

factors of CO2 andCH4 are reported. Table 4 reports critical properties and acentric factors for the ILs

Thtdp-dca and Thtdp-phosphinate. These values were calculated using group contribution methods in

the MSc thesis of T. Olasagasti42

Methane (CH4) Carbon Dioxide (CO2)

Molecular Weight (g/mole) 16.043 44.01 Critical Temperature (Tc) 190.6 K 304.1K

Critical Pressure (Pc) 4.599 MPa 7.376 MPa Acentric factor (ω) 0.012 0.239

Table 3. Critical properties and acentric factors of CO2 andCH4.

Thtdp-[dca] Thtdp-phosphinate Empirical Formula (Hill Notation)

C34H68N3P C48H102O2P2

Critical Temperature (Tc) 1525.5K 1878.8K Critical Pressure (Pc) 0.765 MPa 0.551 MPa Acentric factor (ω) 0.5822 -0.13188

Table 4. Critical properties and acentric factors of ILs [thtdp]-[dca] and [thtdp]-[phosphinate].

20

4.2 Vapor-liquid Equilibria (VLE)

When thermodynamics is applied to vapor/liquid equilibrium, the goal is to find by

calculations, the temperatures, pressures and compositions of phases in equilibrium. To achieve that,

appropriate models for the behavior of the system in vapor-liquid equilibrium are needed. The two

simplest models are Raoult’s and Henry’s law40.

The mathematical expression for Raoult’s law is:

!!! = !!!!!"#!!(! = 1,2,… ,!) (4.8)

In which%!! !and !! !are respectively the liquid and vapor phase mole fraction, !!!"#!the saturation

pressure of the pure component ! and ! the total pressure on the system. Two major assumptions are

required in order to reduce VLE calculations to Raoult’s law. Namely that the vapor phase is an ideal

gas and that the liquid phase is an ideal solution. The first assumption means that Raoult’s law only

applies for low to moderate pressures and the second assumption implies that it can have approximate

validity only when the species that comprise the system are chemically similar. It is however valid for

any species present at a mole fraction approaching unity, provided only that the vapor phase is an

ideal gas40.

Application of Raoult’s law to species !! %requires a value for !!!"#!at the temperature of

application, and thus is not appropriate for a species whose critical temperature is less than the

temperature of application. In this case, Henry’s law is the appropriate VLE model which in essence

defines the solubility of the supercritical gas in the solvent. For a species present as a very dilute

solute in the liquid phase, Henry’s law states that the partial pressure of the species in the vapor phase

is directly proportional to its liquid phase mole fraction. Thus,

!!! = !!!! (4.9)

Where !! is the Henry’s constant. Values of !! come from experiments.

Ideal solution behavior is described by the Lewis/Randal rule:

!!!" = !! !! (4.10)

The solid line in Figure 7 represents experimental values of !!. At !! = 1 the line becomes

tangent to the Lewis/Randal rule. In the other limit, !! → 0 !! becomes zero. The ratio !!/!! is

therefore indeterminate in this limit, and application of l’Hôpital’s rule yields:

21

lim!!→!

!!!!= !!!

!!! !!!!= !!

(4.11)

Equation 4.11 defines Henry’s constant !!, as the limiting slope of the !! − !! curve at !! = 0.

The equation of this tangent line expresses Henry’s Law:

!! = !!!! (4.12)

Which is applicable in the limit as !! → 0, but is also of approximate validity for small values

of !!.40

Figure 7. Composition dependence of liquid-phase fugacities for species i in a binary mixture.

22

4.3 Matlab

The implemented code, calculates the bubble point of a mixture using the fundamental relation

of any phase equilibrium calculations; the equality of fugacities of each species in each phase, that is:

!!! = !!!! → !!!!!! = !!!!! (4.7)

This condition is solved using the algorithm presented in the following figure:

Figure 8. Bubble point pressure algorithm for specified temperature and liquid composition.

Fugacity coefficients for both the liquid and gas phase are calculated using the Peng-Robinson

EoS.

23

5. Results

5.1 Binary systems

In this section the experimental results of the bubble-point pressure measurements will be

reported for the three binary systems that were studied. First P-T diagrams of the experimental data

are presented followed by P-x diagrams that include both experimental and modeling results. The

solubilities of CH4 and CO2 in each studied IL are then compared. The experimental data on the

systems [thtdp]-[dca]+CO2 and [thtdp]-[phosphinate]+ CO2 were available in the MSc Thesis work of

T.Z. Olasagasti.42

5.1.1 Thtdp-dicyanamide

The solubility of CH4 in the IL [thtdp]-[dca], was determined at temperatures ranging from

302.13 to 363.48 K and pressures up to 11.57 MPa by measuring the bubble-point pressures at various

compositions of CH4 in the IL. In Figure 9, the pressure in plotted versus the temperature for this

system. It can be concluded that the bubble-point pressures of CH4 in the IL increases almost linearly

with temperature, indicating that the solubility decreases with increasing temperature.

Figure 9. P-T diagram of the system thtdp-dicyanamide + CH4. The lines plotted are meant as a guide to the eye.

Among the three different ILs that were used for solubility measurements, [thtdp]-[dca] was the

less viscous. Due to the relatively low viscosity, the magnetic stirrer was able to operate at higher

speeds thus helping the system to reach an equilibrium state faster after small changes in pressure.

0!

5!

10!

15!

290! 305! 320! 335! 350! 365! 380! 395!

P"[M

Pa]"

T"[K]"

1.5!wt!%!CH4!

1.25!wt!CH4!

1!wt!%!CH4!

0.75!wt!%!CH4!

0.5!wt%!CH4!

0.35!wt!%!CH4!

0.25!wt!%!CH4!

24

Finally it is worth mentioning that due to gas leakage issues during sample preparation, most of the

measurements for this system had to be repeated up to three times. The issue was later resolved after

thorough rinsing of the gas rack was established before every filling.

As previously mentioned, the Peng-Robinson EoS is used to model the acquired experimental

data. Minimizing the deviation between the experimental and the calculated data gives the interaction

parameters of the model. These parameters along with the deviation between experimental and

calculated bubble-point pressures are listed in Table 3.

T(K)! k(1,2)! I(1,2)! Deviation!(%)!

303.15" 0.09! 0.25! 2.18!313.15" 0.08! 0.25! 1.88!

323.15" 0.08! 0.25! 1.71!

333.15" 0.07! 0.25! 1.64!343.15" 0.06! 0.25! 1.66!

353.15" 0.06! 0.25! 1.77!

363.15" 0.07! 0.22! 1.98! Table 5 Model interaction parameters and deviation between experimental and calculated pressures for the system [thtdp]-[dca]+CH4.

Figure 10. Model interaction parameters kij plotted against temperature for system [thtdp]-[dca]+CH4.

In Figure 11, the bubble-point pressures are plotted against the methane mole fraction at fixed

temperatures. In order to calculate these lines, measured data were extrapolated to exact temperatures.

It can be observed that higher pressures are needed for larger amounts of methane to be absorbed.

Furthermore, the bubble-point pressures increase slightly more than linearly with increasing CH4

0!

0.03!

0.06!

0.09!

0.12!

0.15!

300! 310! 320! 330! 340! 350! 360! 370!

kij"

T(K)"

25

concentration at given temperature. This is different from the typical behavior for IL + CO2 systems,

where the CO2 solubility curves exhibit a much stronger convex behavior at similar pressures.48 This

behavior of CO2 in ILs is more clearly illustrated in Figure 12.

Figure 11. P-x diagram of the system [thtdp]-[dca]+ CH4. Markers correspond to experimental bubble-point pressures and lines correspond to modeling results.

As it can be seen in Figure 11, calculated and experimental data are in good agreement, which

means that the interaction parameters are reliable.

Figure 12. P-x diagram of the systems [thtdp]-[dca]+ CH4 and [thtdp]-[dca]+CO2 42 at 303.15 K.

0!

2!

4!

6!

8!

10!

12!

0.06$ 0.11$ 0.16$ 0.21$ 0.26$ 0.31$ 0.36$

Pressure"(MPa)"

xCH4""

363K!exp.! 363K!model!353K!exp.! 353K!model!343K!exp.! 343K!model!333K!exp! 333K!model!323K!exp.! 323K!model!313K!exp.! 313K!model!303K!exp.! 303K!model!

0!

2!

4!

6!

8!

10!

12!

14!

0.0! 0.1! 0.2! 0.3! 0.4! 0.5! 0.6! 0.7! 0.8! 0.9!

P(MPa)"

x"

CH4!

CO2!!

26

In Figure 12, the solubility of CH4 and CO2 in the IL [thtdp]-[dca] are plotted in the same

graph. Experimental data of CO2 solubility in the IL [thtdp]-[dca] were available in the MSc Thesis

work of T.Z. Olasagasti.42 The clearly shown difference in solubility of CO2 and CH4 in the IL

illustrates the potential of this IL to be used in relevant gas separation processes.

T(K) HCO2 (MPa) HCH4 SCO2/CH4 (HCH4/HCO2)

298.15" 2.70! 16.71! 6.18!

303.15" 3.00! 17.24! 5.75!

313.15" 3.58! 18.24! 5.10!

323.15" 4.14! 19.15! 4.63!

333.15" 4.67! 19.96! 4.27!

343.15" 5.18! 20.68! 3.99!

353.15" 5.68! 21.30! 3.75!

363.15" 6.15! 21.83! 3.55!

Table 6. Henry’s constants of CH4 and CO2 and ideal CO2/CH4 selectivity in the IL thtdp-[dca].

Table 6 compares the Henry’s constants for the solubility of CH4 and CO2 in the IL thtdp-[dca]. The

ideal selectivity for the CO2/CH4 separation is calculated by dividing the Henry’s constant of CH4 by

the Henry’s constant of CO2. The separation ratios of CO2/CH4 range from 3.75 to 6.18. Maximum

separation ratio is achieved at the lowest temperature. Henry constants are plotted against temperature

in Figure 13.

Figure 13. Henry constants of CH4 and CO2 in the IL [thtdp]-[dca] as a function of temperature.

0!

5!

10!

15!

20!

25!

300! 310! 320! 330! 340! 350! 360! 370!

H"(MPa)"

T(K)"

CH4! CO2!

27

5.1.2 Thtdp-phosphinate

The second binary system that was studied is that of IL [thtdp]-[phosphinate] + CH4. The

solubility of CH4 in this IL was determined at temperatures ranging from 302.00 to 363.27 K and

pressures up to 12.05 MPa. For this system, the pressure is plotted versus temperature in Figure 14.

While the behavior of this system is very similar to the system [thtdp]-[dca]+CH4, the bubble-point

pressures for similar gas mole fractions are much lower, which means an increased solubility. This

particular IL is very viscous, making the preparation of the samples and bubble-point pressure

measurements harder. Establishing an equilibrium state of the system, after small changes in pressure,

takes longer compared to other systems studied.

Figure 14. P-T diagram of the system [thtdp]-[phosphinate] + CH4. The lines plotted are meant as a guide to the eye.

0!

5!

10!

15!

290! 305! 320! 335! 350! 365! 380! 395!

P(MPa)"

T(K)"

2.0!wt%!CH4!

1.5!wt%!CH4!

1!wt!%!CH4!

0.8!wt%!CH4!

0.5!wt%!CH4!

0.25!wt%!CH4!

28

The Peng-Robinson EoS was also employed for modeling this system. In the following table,

model interaction parameters and deviation between experimental and model-calculated bubble-point

pressures are reported.

T(K)! k(1,2)! I(1,2)! Deviation!(%)!

303.15" 0.13! 0.47! 5.31!313.15" 0.14! 0.44! 5.11!323.15" 0.15! 0.42! 4.98!333.15" 0.16! 0.4! 4.85!343.15" 0.17! 0.38! 4.70!353.15" 0.17! 0.37! 4.55!363.15" 0.17! 0.36! 4.46!

Table 7. Model interaction parameters and deviation between experimental and calculated pressures for system [thtdp]-[phosphinate]+CH4.

Figure 15. Model interaction parameters kij plotted against temperature for system [thtdp]-[phosphinate]+CH4.

0!

0.04!

0.08!

0.12!

0.16!

0.2!

0.24!

300! 310! 320! 330! 340! 350! 360! 370!

kij"

T"(K)"

29

Figure 16. P-x diagram of the system [thtdp]-[phosphinate]+CH4. Markers correspond to experimental bubble-point pressures and lines correspond to modeling results.

As can be seen in Figure 16, the pressure increases sharply with increased CH4 mole fraction.

Deviation between experimental and calculated data is larger for this system than the [thtdp]-

[dca]+CH4 system.

Figure 17. P-x diagram of the systems [thtdp]-[phosphinate]+ CH4 and [thtdp]-[phosphinate]+CO2 42 at 303.15 K.

In Figure 17, the solubility of CH4 and CO2 in the IL [thtdp]-[phosphinate] are plotted in the same

graph. Experimental data of CO2 solubility in this IL were available in the MSc Thesis work of T.Z.

Olasagasti.42. Table 8 compares the Henry’s constants of CH4 and CO2 in this IL and reports the ideal

0!

2!

4!

6!

8!

10!

12!

14!

0.06$ 0.11$ 0.16$ 0.21$ 0.26$ 0.31$ 0.36$ 0.41$ 0.46$ 0.51$ 0.56$

P(MPa)"

xCH4"

363K!exp.!363K!model!353K!exp.!353K!model!343K!exp.!343K!model!333K!exp.!

0!

2!

4!

6!

8!

10!

12!

0.0! 0.1! 0.2! 0.3! 0.4! 0.5! 0.6! 0.7! 0.8! 0.9!

P(MPa)"

x"

CH4! CO2!

30

selectivities. Separation ratios of CO2/CH4 range from 5.58 to 14.19. Maximum separation ratio is

achieved at the lowest temperatures. Henry constants are plotted against temperature in Figure 18.

"

"

"

"

"

"

"

Table 8. Henry’s constants of CH4 and CO2 and ideal CO2/CH4 selectivity in the IL [thtdp]-[phosphinate].

Figure18. Henry coefficients of CH4 and CO2 in the IL [thtdp]-[phosphinate] as a function of temperature.

T(K) HCO2 (MPa) HCH4 SCO2/CH4 (HCH4/HCO2)

303.15" 0.66! 9.38! 14.19!313.15" 0.71! 9.92! 13.99!323.15" 0.86! 10.42! 12.15!333.15" 1.08! 10.89! 10.10!343.15" 1.37! 11.32! 8.27!353.15" 1.73! 11.72! 6.77!363.15" 2.16! 12.09! 5.58!

0!

2!

4!

6!

8!

10!

12!

14!

300! 310! 320! 330! 340! 350! 360! 370!

H(MPa)"

T(K)"

CH4! CO2!

31

5.1.3 Emim-phosphate

The final binary system that was studied is the system [emim]-[phosphate] and CO2. The

solubility of CO2 in this IL was determined at temperatures ranging from 302.1 to 362.19 K and

pressures up to 11.24 MPa. For this system, experimental bubble-point pressures for various

temperatures and concentrations of CO2 are shown in Figure 19. Solubility of CO2 in this IL sharply

decreases for increasing concentration of CO2 and increasing temperature.

Figure19. P-T diagram of the system [emim]-[phosphate]+CO2 for various concentrations of CO2.The lines plotted are meant as a guide to the eye.

Figure20. P-x diagram of the system [emim]-[phosphate]+CO2, for various temperatures. The lines plotted are meant as a guide to the eye.

0!

2!

4!

6!

8!

10!

12!

290! 305! 320! 335! 350! 365! 380! 395!

P"[M

Pa]"

T"[K]"

10!wt!%!CO2!

7.5!wt!%!CO2!

5.0!wt%!CO2!

2.5wt!%!CO2!

1.5!wt!%!CO2!

0!

3!

6!

9!

12!

0.05$ 0.1$ 0.15$ 0.2$ 0.25$ 0.3$ 0.35$ 0.4$ 0.45$ 0.5$

P"(MPa)"

XCO2"

363K!

353K!

343K!

333K!

323K!

313K!

303K!

32

Figure 21. Henry constants of CO2 in the ILs [emim]-[phosphate], [thtdp]-[dca] and [thtdp]-[phosphinate] as a function of temperature.

In Figure 20, the solubility of CO2 in the IL [emim]-[phosphate] is plotted as a function of CO2

mole fraction for various temperatures. Bubble-point pressures increase sharply as the CO2 mole

fraction in the binary increases. The low solubility of CO2 in this IL is also reflected in the Henry

constants of this system shown in Figure 21. Comparing the Henry constants of CO2 in the three ILs

studied leads to the conclusion that the performance of [emim]-[phosphate] as a CO2 capture solvent

is relatively poor. This fact was a determining factor to not execute CH4 solubility measurements in

this IL, since investigating CO2 selectivity in a solvent with unsatisfactory CO2 solubility would be

rather pointless.

0!

4!

8!

12!

16!

300! 310! 320! 330! 340! 350! 360! 370!

H(MPa)"

T(K)"

emim[phosphate]!

thtdp[dca]!

thtdp[phosphinate]!

33

5.2 Ternary Systems

The experiments executed so far provide useful data on single gas (CO2 and CH4) solubilities in

various ILs. Use of this data made also possible to determine the ideal selectivities of the examined

solvents. Nevertheless, as natural gas processes operate at high pressures, ideal selectivity should not

be expected for the CO2/CH4 separation. Mixed gas solubility data are very scarce in the literature and

in fact there is no single straightforward way to calculate real selectivities. In the present work, mixed

gas solubilities are measured in ternary systems of IL+CO2+CH4. The mixed gas bubble-point

pressures measured are later compared to the sum of the single gas bubble-point pressures of the

binary systems. This way it is determined whether or not there is significant interaction between the

two gases that eventually results in real selectivity deviating from ideal selectivity.

Figure 22. Qualitative ternary diagram of all the binary and ternary systems that contained the IL thtdp-[dca].

The experimental procedure was exactly the same as in the case of binary mixtures, only this

time gas mixtures of known CO2 and CH4 concentrations were introduced in the Cailletet tubes,

together with a single IL each time, to form ternary systems. The three gas mixtures were ordered and

delivered in separate gas cylinders each containing respectively: 25 mole% CO2 - 75 mole% CH4, 50

mole% CO2 - 50 mole% CH4, 75 mole% CO2 - 25 mole% CH4. For each IL, four ternary systems were

34

prepared with each of the three gas mixtures. In total 24 ternary systems were studied in a temperature

range roughly from 303 to 363K. In figures 22 and 23 a more detailed representation of the executed

experiments is given.

Figures 22 and 23 are triangular ternary diagrams. In each of the three axes of these diagrams,

the mole fraction of each component is shown ranging from 0 to 1. Each red point located in the

triangular plane, represents a fixed ternary composition for which bubble-point pressure

measurements were performed for temperatures ranging from 303 to 363 K. The red points located on

the axes of the diagram represent fixed binary compositions for which again bubble-point pressures

were measured. Quantitative 2-D ternary diagrams are drawn for a single temperature and pressure.

While the temperature range of each studied system is the same (303-363K) the pressure range is

different. A quantitative representation of all the available data would require a 3-D ternary diagram.

Figures 22 and 23 constitute qualitative representations of the executed experiments in the sense that

they illustrate, in a simple 2-D plane, all the systems whose different pressure would position them in

different planes along a third pressure axis (perpendicular to the triangular plane) of a 3-D ternary

diagram.

Figure 23. Qualitative ternary diagram of all the binary and ternary systems that contained the IL thtdp-[phosphinate].

35

Figure 22 includes all the systems that contain the IL thtdp-[dca] while Figure 23 the systems with

thtdp-[phosphinate]. The red points located on the CO2 mole fraction axes, represent experimental

CO2 solubility data available in the Thesis work of T. Olasagasti.42

Each of the blue lines represents an axis where the mole/mole ratio of CO2/CH4 remains

constant. There is such an axis for each of the three gas mixtures used. The green lines on the other

hand represent axes of constant mole/mole CH4/IL ratios. This illustrates the set-up of the

experimental procedure. Every binary system of CH4+IL that was measured had a certain CH4/IL

ratio. For each of these CH4/IL ratios, three ternary systems with the same CH4/IL ratio were prepared

and measured, using the three different gas mixtures. That way, on each green line, there are 3 ternary

systems and one binary system that maintain a constant CH4/IL ratio. The CO2/IL ratio for every

ternary system however, changes as a different gas mixture is used for each one of them. As

mentioned earlier, the goal of this study is to investigate and compare mixed gas to single gas

solubilities in the same ILs. In order for such a comparison to be possible, it was necessary to

calculate theoretical mixed gas pressures based on the available single gas pressure data and with the

assumption that there is no interaction between the two gasses. From binary mixture data of CH4+IL,

single gas pressures were extrapolated to match the exact CH4/IL ratios found in ternary systems. In

the same manner, using binary mixture data of CO2+IL42, single gas pressures were also calculated for

the same CO2/IL ratios as in the ternary systems. The calculated (theoretical) mixed gas bubble-point

pressure was achieved by adding the pressures (for CO2 and CH4) obtained from the binary systems.

These theoretical mixed gas pressures are compared with the actual (real) ternary bubble-point

pressures (for various temperatures) in the following figures.

Figure 24. Calculated (theoretical) and measured (actual) mixed gas solubility for various temperatures and ratios of CH4/IL in the ternary system thtdp-[phosphinate]+ (50 mole% CO2 – 50 mole% CH4).

0$

1$

2$

3$

4$

5$

6$

7$

8$

9$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&

Pcalc.$for$CH4/IL=0.487$ Pternary$for$CH4/IL=0.487$Pcalc.$for$CH4/IL=0.387$ Pternary$for$CH4/IL=0.387$Pcalc.$for$CH4/IL=0.242$ Pternary$for$CH4/IL=0.242$Pcalc.$for$CH4/IL$=$0.119$ Pternary$CH4/IL=0.119$

36

Figure 24 compares theoretical to actual mixed-gas solubility in the ternary system thtdp-

[phosphinate] + (50mole%CO2 –50mole%CH4). The lines represent the calculated pressures while the

markers refer to actual (measured) ternary bubble-point pressures. Each set of lines and markers refers

to one of the four different CH4/IL ratios studied. Figure 25 illustrates theoretical and actual mixed-

gas solubility for the ternary systems of thtdp-[phosphinate] + (25mole%CO2-75mole%CH4) while

Figure 26 refers to the ternary system thtdp-[phosphinate] + (75 mole%CO2-25mole%CH4.)

Figure 25. Calculated (theoretical) and measured (actual) mixed gas solubility for various temperatures and ratios of CH4/IL in the ternary system thtdp-[phosphinate]+ (25mole%CO2-75mole%CH4).

Figure 26. Calculated (theoretical) and measured (actual) mixed gas solubility for various temperatures and ratios of CH4/IL in the ternary system thtdp-[phosphinate]+ (75mole%CO2-25mole%CH4).

0$

1$

2$

3$

4$

5$

6$

7$

8$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&


0$

1$

2$

3$

4$

5$

6$

7$

8$

9$

10$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&

Pcalc.$for$CH4/IL=0.387$ Pternary$for$CH4/IL=0.387$

Pcalc.$for$CH4/IL=0.242$ Pternary$for$CH4/IL=0.242$

Pcalc.$for$CH4/IL$=$0.119$ Pternary$for$CH4/IL=0.119$

37

In Figures 24 to 26, it is illustrated that theoretical and actual mixed gas solubilities have very

small deviation. The highest deviation observed was of 0.2 MPa. As the ratio of CH4/IL increases the

deviation between theoretical and actual solubility appears to increase. This however, has to do with

the extrapolation error in calculated pressures. Theoretical pressures are extrapolated by binary

systems to match the exact CH4/IL and CO2/IL ratios found in the ternary systems. To achieve the

needed extrapolation, data of CH4/IL ratios versus binary bubble-point pressures need to be fitted. The

same is done for data of CO2/IL ratios versus binary bubble-point pressures. While solubility data for

CH4 are satisfactorily fitted with a second order polynomial, CO2 solubility data become highly non-

linear for increasing CO2/IL ratios and therefore fitting this data induces larger error. Fitting binary

CO2 solubility data was tried using alternatives from linear to fifth order polynomial. The smallest

error was observed when using a third-order polynomial and therefore that was the method adopted to

calculate theoretical mixed gas solubilities. In Figure 27, theoretical and actual mixed-gas solubilities

are reported for the ternary system thtdp-[dca] + (50mole%CO2-50mole%CH4). Figures 28 and 29

illustrate the same information for the ternary systems thtdp-[dca] + (25mole%CO2-75mole%CH4)

and thtdp-[dca] + (75mole%CO2-25mole%CH4) respectively. Small deviation between calculated and

experimental mixed-gas solubilities is also observed for these systems. Deviation again increases with

increasing CH4/IL ratios, which again has to do with extrapolation error.

Figure 27. Calculated (theoretical) and measured (actual) mixed gas solubility for various temperatures and ratios of CH4/IL in the ternary system thtdp-[dca]+ (50mole%CO2-50mole%CH4).

0$

1$

2$

3$

4$

5$

6$

7$

8$

9$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&


38



In Figure 28 there are no measured pressure points for the temperatures of 353K and 363K.

This is because the Cailletet tube broke before these points could be measured. Making an entire new

filling just to obtain these two points was deemed unnecessary.

0$

2$

4$

6$

8$

10$

12$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&


0$

1$

2$

3$

4$

5$

6$

7$

8$

9$

300$ 310$ 320$ 330$ 340$ 350$ 360$ 370$

P[MPa

]&

T[K]&


39

Ternary Systems Summary

For these experiments, three different gas mixtures were used, namely 25mole%CO2-

75mole%CH4, 50mole%CO2-50mole%CH4, 75 mole%CO2-25mole%CH4. For each of the two

phosphonium based ILs used, bubble-point pressures were measured for 12 systems (4 systems per

gas mixture) and for temperatures between 303K and 363K. Regardless of which IL formed the

ternary system measured, common observation was the fact that the deviation between single gas and

mixed gas solubility appeared to increase as the CO2 mole fraction increased in the system. In other

words, deviation between calculated and actual mixed-gas solubility appeared higher in the ternary

systems formed using the 75 mole%CO2-25mole%CH4 gas mixture. Deviation in the system of IL+

25mole%CO2-75mole%CH4 was between 0.003 MPa and 0.18MPa (0.33-4.1%). Deviation in the

systems of IL+ 50mole%CO2-50mole%CH4 was between 0.002 MPa and 0.22MPa (0.03-4.8%).

Finally deviation in the systems of IL+75mole%CO2-25mole%CH4 was between 0.002 MPa and

0.34MPa (0.1-5.3%). As mentioned before, these deviations are induced by extrapolation errors due to

the highly non-linear behavior of the system IL+CO2, as pressure increases. Temperature increase did

not affect the calculated deviation. Besides the fact that the found deviations between calculated and actual (measured) mixed-gas

solubility are low (maximum 5.3%), the certainty that the aforementioned deviations are due to

extrapolation error comes from the fact that for a given CO2/IL (extrapolated) ratio, the error depends

strongly on the order of the polynomial that was chosen for data fitting. For low CO2/IL ratios a

second order polynomial fitting gives good results. As we move to higher CO2/IL ratios, the highly

non-linear behavior of CO2 is best fitted by a much higher order polynomial. Fitting the entire CO2/IL

ratios versus bubble-point pressure curves with the lowest error, was achieved by using a third order

polynomial. Overall it is safe to conclude that there is indeed no strong interaction between the gasses,

such that would result in real selectivity to deviate significantly from ideal selectivity. For the studied

separation with phosphonium based ILs, ideal selectivity can be assumed.

40

6. Conclusion and Future direction

6.1 Conclusion

The experiments executed during the course of the present work provide extensive and reliable

CO2 and CH4 solubility data in certain ILs. Furthermore mixed gas solubilities in ILs are also

measured and reported. In detail the following concluding remarks can be drawn:

• The first measurements to be executed were those of the binary systems [thtdp]-

[dca]+CH4 and [thtdp]-[phosphinate]+CH4. These two systems have similar behavior in

the sense that CH4 solubility decreases with increasing CH4 mole fraction with the

same trend. The actual pressure however, needed to dissolve the same amount of gas is

significantly higher for the [thtdp]-[dca] which means that the solubility of CH4 is

higher in the [thtdp]-[phosphinate]. This is also illustrated in the Henry constants of the

two systems.

• For all the systems studied (binary as well as ternary) common observation was the fact

that solubility of CH4 and/or CO2 decreased for increasing temperature.

• Modeling of the above binary systems using the Peng-Robinson EoS showed that

calculated values and experimental data are in good agreement, which means that the

interaction parameters are reliable.

• Literature data on CO2 solubility42 in [thtdp]-[phosphinate] and [thtdp]-[dca] show a

better performance of the IL [thtdp]-[phosphinate]. This together with the findings of

the present work on CH4 solubilities in the same ILs leads to higher ideal CO2/CH4

selectivity in the [thtdp]-[phosphinate]. This IL is therefore more suitable to be

considered as a gas separation solvent.

• Although gas separation performance for the IL [thtdp]-[phosphinate] seems superior

to [thtdp]-[dca], [thtdp]-[phosphinate] is much more viscous. This made measurements

more difficult and time consuming since more time was needed to accurately achieve

equilibrium. While viscosity decreases with increased temperature, so does gas

solubility in the IL requiring higher pressures.

• Apart from the two phosphonium based ILs mentioned above, gas solubility

measurements were also executed with an imidazolium based IL ([emim]-[phosphate]).

The solubility of CO2 in this IL was measured and found to be very low compared to

the phosphonium based ILs.

Comparing the Henry constants of CO2 in the three ILs studied leads to the conclusion

that the performance of [emim]-[phosphate] as a CO2 capture solvent is relatively poor.

This fact was a determining factor to not execute CH4 solubility measurements in the

41

imidazolium based IL, since investigating CO2 selectivity in a solvent with

unsatisfactory CO2 solubility would be rather pointless.

• Measuring mixed gas solubilities was the second part of the present work. As natural

gas processes operate at high pressures, ideal selectivity should not be expected for

CO2/CH4 separation. According to the findings of the present work however, single gas

and mixed gas solubilities have very slight deviations.

• Three different gas mixtures were used, namely 25mole%CO2-75mole%CH4,

50mole%CO2-50mole%CH4, 75 mole%CO2-25mole%CH4. With each of the two

phosphonium based ILs, 12 ternary systems were measured. The deviation between

single gas and mixed gas solubilities ranges between 0.01 to 0.34 MPa, which is

equivalent to a maximum deviation of 10% between calculated and actual mixed gas

solubilities. These deviations however, are induced by extrapolation errors and due to

the highly non-linear CO2 phase behavior as pressure increases.

• The above experimental finding leads to the conclusion that ideal selectivity can in fact

be expected for the CO2/CH4 separation with phosphonium based ILs.

42

6.2 Future Directions

ILs are a new category of solvents with potentially broad industrial applicability. Further

research is however needed in the vast synthetic landscape provided by ILs, especially for carbon

capture or gas processing purposes.

Specifically for natural gas sweetening processes, while CO2 solubility data are increasingly

available, CH4 solubility data are still rather scarce. Further research on CH4 solubility in ILs with

known high CO2 solubility may lead to IL solvents with increasingly favorable properties for natural

gas sweetening.

In natural gas sweetening processes, apart from the main separation (CO2/CH4), CO2/H2S and

CO2/N2 selectivities are also very relevant. Further research is needed in order to experimentally

acquire relevant selectivity data.

The present work focuses on experimental determination of solvent properties in a lab-scale. In

order however to conclude about the efficiency and cost-effectiveness of these solvents in real

processes, scale-up studies are needed. Real pilot processes using ILs as solvents need to be

researched and designed before we can assess the feasibility of ILs on industrial scale. Industry almost

always demands demonstration on pilot scale before commercialization.

Despite the many advantages of ILs as industrial solvents, there is a profound lack of relevant

safety, health, and environmental studies. To prevent environmental pollution, biodegradability and

toxicity of candidate ILs should be thoroughly assessed.

The current lab-scale price of ILs is extremely high. The price will drop for a large scale

production of ILs, but a price level typical for conventional solvents should not be expected since ILs

are complex molecules, requiring advanced synthesis and purification steps. This is an important

limiting factor for the large-scale implementation of ILs as industrial solvents.

43

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39.! Carlisle,!T.!K.,!Bara,!J.!E.,!Gabriel,!C.!J.,!Noble,!R.!D.!&!Gin,!D.!L.!Interpretation!of!CO2!Solubility!and!

Selectivity!in!NitrileeFunctionalized!RoomeTemperature!Ionic!Liquids!Using!a!Group!Contribution!

Approach.!7005–7012!(2008).!

40.! Smith,! J.M.,! Van!Ness,!H.C.,! Abbott,!M.!M.! Introduction* to*Chemical*Engineering*Thermodynamics.!

(McGraweHill,!2005).!

41.! Raeissi,! S.! &! Peters,! C.! J.! Experimental! determination! of! highepressure! phase! equilibria! of! the!

ternary!system!carbon!dioxide+limonene+linalool.!J.*Supercrit.*Fluids!35,!10–17!(2005).!

42.! Zuzuarregui!Olasagasti,!T.!CO2!Capture!with!Ionic!Liquids.!99!(2010).!

43.! SIGMAeALDRICH,!http://www.sigmaaldrich.com/nederland.html!

44.! Álvarez,! V.! H.! &! Aznar,! M.! Thermodynamic! modeling! of! vapor–liquid! equilibrium! of! binary!

systems! ionic! liquid+supercritical! {CO2! or! CHF3}! and! ionic! liquid+hydrocarbons! using! Peng–

Robinson!equation!of!state.!J.*Chinese*Inst.*Chem.*Eng.!39,!353–360!(2008).!

45.! Sandler,!S.!Chemical,*Biochemical*and*Engineeting*Thermodynamics.!(John!Wiley!&!Sons,!2006).!

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Acentric!Factor!of!Another!200!Ionic!Liquids.!1318–1330!(2008).!

48.! Raeissi,! S.! &! Peters,! C.! J.! High! pressure! phase! behaviour! of! methane! in! 1ebutyle3e

methylimidazolium!bis(trifluoromethylsulfonyl)imide.!Fluid*Phase*Equilib.!294,!67–71!(2010).!

49.! Dymond,! J.!H.!&!Smith,!E.!B.!The*Virial*Coefficients*of*Pure*Gases*and*Mixtures.! (Clarendon!Press,!

1980).!

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51.! Mallu,!B.!V.!Compression!factors!and!second!coefficients!of!H2!and!CH4.!(1990).!!

46

Appendix A

Calculation of the second virial coefficients of pure gasses and gas mixtures.

• Second Virial coefficients for single gases:

!(!) = !!!!"#!!"#

− 1!!!!

!!!

Where Ai are the virial coefficient of every gas, Tref=298.15 and Tatm is the temperature of the gas

when the filling is carried out. In the table that follows the virial coefficients of CO2 and CH4 are

reported 49:

A1 A2 A3 A4

CH4 -43,46 -113.5 -19.33 -6.66

CO2 -127.3 -287.6 -118.3 0

• Second Virial Coefficients for gas mixtures:

Mixture second virial coefficient (Bm) is a know function of temperature but also of the gases’

mole fractions comprising the mixture (!!!!):

!! = !!!!!!"!

!!!

!

!!!

The !!" !is a function of temperature only, and is calculated according to the following

correlations (Tsonopolos 1974)50 for a binary non-polar gas mixture:

!!"!!!"!!!!"

= !(!) !!!" + !!(!) !!!"

Where: ! ! !!!" = 0.1445 − !.!!"!!!"

− !.!"#$!!!"

! − !.!"#"!!!"

! − !.!!!"!#!!!"

! = !!

! ! !!!" = 0.0637 + 0.331!!!"!− 0.423!!!"!

− 0.008!!!"!= !!

47

So !!" =!!!!"(!!!!!!)

!!!"

Where:

!!!" = !!!!!!"!! (1 − !!")

!!!" =!!!!"

!!!" =!!!

!! + !!!

!!

!

8

!!!" = 0.5(!!! + !!!)

!!!" =4!!!"(!!!

!!!!!!

+ !!!!!!!!!)

!!!!! + !!!

!!

!

Where: !!! ,!!! ,!!! ,!!! ,!!! ,!!! !are pure component critical temperatures, pressures and

volumes respectively. !!" is the binary interaction parameters.

So at this point !!" = !!(!!" , !"#$!!"#$"%&%'!!! ,!! ,!!) The procedure now goes iteratively. !!" is randomly chosen and then !!"!!"#$. is calculated for

various temperatures.

Experimental !!"! for binary CO2-CH4 mixtures was available in literature (Mallu 1990)51.

By iteratively minimizing the objective function ! = !!"#$%.!!!"!!"#$.!!"#$%.

, the optimal binary

interaction parameter !!" is calculated. Finally the optimal value of !!" is calculated for each

temperature.

48

Appendix B

Experimental Measurements

BubbleVpoint*pressures*for*the*system:"Thtdp(dicyanamide+CH4"!

!0.25%wt%%CH4%(0.0794%mole%fraction)% % 0.35%wt%%CH4%(0.1076%mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.21% 20.27% 1.427% % 302.25% 19.51% 2.027%313.21% 20.34% 1.502% % 322.28% 19.89% 2.213%323.20% 20.00% 1.576% % 332.32% 19.98% 2.298%333.26% 20.00% 1.651% % 342.39% 20.39% 2.374%343.21% 19.63% 1.696% % 352.38% 20.49% 2.454%353.25% 19.73% 1.752% % 362.39% 20.46% 2.524%363.28% 19.70% 1.792% % 312.28% 19.61% 2.117%!0.5%wt%%CH4%(0.1481%mole%fraction)% % 0.75%wt%%CH4%(0.2052 mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.34% 20.34% 2.963% % 302.46% 20.06% 4.457%312.17% 20.39% 3.118% % 312.58% 20.10% 4.733%322.29% 20.19% 3.243% % 322.77% 19.91% 4.887%332.32% 20.58% 3.363% % 332.67% 20.01% 5.067%342.33% 20.57% 3.478% % 342.66% 20.00% 5.228%352.41% 20.52% 3.593% % 352.69% 20.00% 5.378%362.35% 20.52% 3.683% % 362.83% 20.05% 5.523%!1%wt%%CH4%(0.2573 mole%fraction)% % 1.25%wt%%CH4%(0.3030 mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.38% 19.73% 6.294% % 302.13% 18.76% 7.855%313.39% 19.56% 6.535% % 312.14% 18.86% 8.195%323.47% 20.00% 6.799% % 322.25% 19.86% 8.536%333.52% 19.75% 7.056% % 332.29% 19.86% 8.821%343.53% 20.09% 7.266% % 342.31% 19.86% 9.086%353.46% 19.77% 7.451% % 352.36% 19.82% 9.327%363.48% 20.09% 7.621% % 362.34% 19.88% 9.537%!1.5%wt%%CH4%(0.3432 mole%fraction)%T/K% Troom/%°C% Pactual/MPa%302.18% 19.66% 9.638%312.25% 19.68% 10.048%322.28% 19.49% 10.433%332.32% 19.77% 10.758%342.36% 20.05% 11.059%352.44% 20.00% 11.329%362.43% 20.02% 11.569%*

49

BubbleVpoint*pressures*for*the*system:"Thtdp(phosphinate+CH4"!

!

!

!0.25%wt%%CH4%(0.1067%mole%fraction)% % 0.5%wt%%CH4%(0.1949%mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.65% 19.68% 1.0154% % 303.25% 20.90% 2.1434%312.05% 18.83% 1.0655% % 313.25% 21.80% 2.2329%322.07% 20.50% 1.1307% % 323.19% 23.00% 2.3545%332.09% 20.20% 1.1709% % 333.22% 23.95% 2.4495%342.12% 20.08% 1.2159% % 343.25% 24.90% 2.5446%352.11% 19.85% 1.2610% % 353.26% 25.90% 2.6446%362.15% 20.22% 1.3011% % 363.27% 26.80% 2.7346%!!0.8%wt%%CH4%(0.2792 mole%fraction)% % 1%wt%%CH4%(0.3276 mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.42% 20.16% 3.3487% % 302.70% 20.00% 4.3846%312.04% 20.29% 3.5288% % 312.56% 19.92% 4.6098%322.05% 20.47% 3.7075% % 322.66% 20.00% 4.8748%332.04% 20.15% 3.8782% % 332.62% 19.95% 5.0492%342.05% 20.45% 4.0385% % 342.63% 20.18% 5.2394%352.15% 20.30% 4.1887% % 352.70% 20.36% 5.4246%362.20% 20.11% 4.3189% % 362.68% 20.22% 5.5949%%!1.5%wt%%CH4%(0.4238 mole%fraction)% % 2%wt%%CH4%(0.4963 mole%fraction)%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.97% 19.85% 6.5697% % 302.00% 19.92% 9.7081%312.30% 19.85% 6.9050% % 311.98% 20.26% 10.1734%322.07% 19.96% 7.2144% % 322.07% 20.70% 10.6281%332.02% 19.26% 7.5148% % 332.09% 20.55% 11.0234%342.12% 19.58% 7.8050% % 342.11% 20.16% 11.4038%352.17% 19.76% 8.0752% % 352.13% 19.12% 11.7344%362.16% 19.12% 8.3206% % 362.16% 20.60% 12.0492%!!

50

BubbleVpoint*pressures*for*the*system:""Emim(phosphate+CO2""!!

!1.5%%wt%%CO2% % 2.5%%wt%%CO2%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.69% 20.40% 0.5487% % 302.21% 19.78% 0.9096%311.97% 21.83% 0.6782% % 312.05% 20.32% 1.1205%322.02% 21.81% 0.7983% % 322.02% 20.12% 1.3506%332.01% 21.48% 0.9236% % 332.08% 19.76% 1.6057%342.08% 20.41% 1.0744% % 342.11% 21.55% 1.8758%352.14% 20.60% 1.2295% % 352.10% 19.40% 2.1512%362.13% 20.23% 1.3897% % 362.18% 20.21% 2.4360%!!5%%wt%%CO2% % 7.5%%wt%%CO2%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.01% 19.77% 1.8963% % 302.13% 20.56% 2.9485%312.00% 20.27% 2.3066% % 312.05% 20.25% 3.6141%322.04% 20.06% 2.7619% % 322.05% 20.42% 4.3345%332.03% 19.90% 3.2823% % 332.07% 20.38% 5.1451%342.11% 20.00% 3.8576% % 342.09% 20.40% 6.0156%352.11% 19.89% 4.4531% % 352.12% 20.60% 6.9808%362.16% 20.14% 5.0885% % 362.19% 20.22% 8.0065%%!10%%wt%%CO2%T/K% Troom/%°C% Pactual/MPa%302.07% 20.02% 3.8174%312.00% 19.96% 4.7981%322.02% 20.13% 5.8237%332.02% 20.42% 6.9394%342.11% 20.44% 8.2702%352.16% 20.43% 9.7111%362.18% 20.32% 11.2421%!!! !

51

BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[phosphinate]"+"(25mole%CO2G"75mole%CH4)"""

!XCH4%=%0.1031%%%XCO2%=%0.03454% % XCH4%=%0.1828$$$$XCO2%=%0.0613&T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.44% 21.25% 1.0509% % 304.01% 22.28% 2.1758%312.63% 20.52% 1.1130% % 312.58% 21.33% 2.2664%322.63% 22.60% 1.1765% % 322.64% 21.62% 2.4174%332.66% 22.58% 1.2366% % 332.61% 21.20% 2.5268%342.64% 22.61% 1.2967% % 342.62% 21.54% 2.6469%352.75% 22.46% 1.3618% % 352.70% 21.55% 2.7771%362.81% 20.77% 1.4135% % 362.72% 21.66% 2.8972%!XCH4%=%0.2553% % %XCO2%=%0.08556& % XCH4%=%0.2953&&&XCO2%=%0.09896&T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.29% 21.51% 3.4826% % 303.16% 22.54% 4.4273%313.18% 21.22% 3.6980% % 313.17% 21.31% 4.6878%323.21% 21.97% 3.9230% % 323.23% 21.69% 4.9988%333.22% 22.85% 4.1031% % 333.28% 21.66% 5.2590%343.20% 22.54% 4.2982% % 343.30% 21.09% 5.4633%353.19% 22.61% 4.4833% % 353.16% 21.47% 5.7034%363.29% 22.53% 4.6784% % 363.26% 21.57% 5.9435%%BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[dicyanamide]"+"(25mole%CO2G"75mole%CH4)"!XCH4%=%0.0772% % %XCO2%=%0.0258& % XCH4%=%0.104&&&XCO2%=%0.0348&T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.16% 20.94% 1.517% % 303.23% 20.48% 2.1978%312.26% 20.98% 1.6021% % 313.23% 20.32% 2.3128%322.49% 20.28% 1.6822% % 323.27% 21.31% 2.4385%332.48% 20.41% 1.7673% % 333.24% 21.09% 2.5330%342.26% 21.16% 1.8425% % 343.27% 20.77% 2.6281%352.36% 20.98% 1.9026% % 353.58% 21.31% 2.7277%362.52% 20.74% 1.9721% % 363.40% 21.79% 2.8177%%XCH4%=%0.1413&&&XCO2%=%0.0473& % XCH4=%0.1921$$$XCO2%=%0.0644%T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.39% 21.48% 3.1386% % 303.60% 20.60% 4.7314%312.63% 20.76% 3.3025% % 313.95% 20.30% 4.9962%322.73% 20.94% 3.4878% % 323.19% 21.77% 5.210%332.58% 20.46% 3.6429% % 333.18% 21.39% 5.4211%342.67% 21.73% 3.8027% % 343.24% 20.89% 5.6360%352.74% 21.09% 3.9180% %362.69% 20.57% 4.0337% %*

52

BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[phosphinate]"+"(50"mole%CO2"–50"mole%CH4)"""

!XCH4%=%0.0961%%%XCO2%=%0.09687% % XCH4%=%0.1626$$$XCO2%=%0.1639$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.25% 21.88% 1.0826% % 303.27% 20.17% 2.3570%313.25% 21.65% 1.1576% % 313.37% 20.88% 2.5028%323.23% 19.91% 1.2422% % 322.23% 19.83% 2.6777%333.20% 19.67% 1.3225% % 332.24% 19.68% 2.8429%343.15% 20.17% 1.4125% % 342.60% 19.58% 3.0481%353.23% 20.42% 1.5075% % 352.70% 24.3% 3.2240%363.12% 19.56% 1.6080% % 362.76% 20.21% 3.4095%!XCH4%=%0.2181$$$XCO2%=%0.2199$ % XCH4%=%0.2461$$$$XCO2%=%0.2480$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.26% 20.06% 3.9138% % 303.25% 19.36% 4.9207%313.25% 20.13% 4.1585% % 313.24% 19.75% 5.2609%323.24% 21.42% 4.4244% % 323.23% 19.45% 5.6061%333.22% 20.96% 4.6694% % 333.26% 20.02% 5.9117%343.28% 20.47% 4.9146% % 343.34% 20.65% 6.2279%353.32% 20.54% 5.1841% % 353.31% 20.19% 6.5828%363.29% 20.57% 5.4742% % 363.26% 19.98% 6.9431%%BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[dicyanamide]"+"(50"mole%CO2"–50"mole%CH4)"!XCH4%=%0.07338$$$XCO2%=%0.07397$ % XCH4%=%0.0971$$XCO2%=%0.0979$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.52% 19.37% 1.6701% % 302.86% 20.21% 2.3968%312.86% 19.6% 1.7853% % 312.37% 20.24% 2.5669%322.45% 19.75% 1.8856% % 322.40% 20.27% 2.7279%332.45% 19.75% 2.0008% % 332.47% 20.24% 2.8780%342.54% 19.76% 2.1109% % 342.44% 20.27% 3.0232%352.49% 19.74% 2.2111% % 352.55% 19.91% 3.1627%362.55% 19.7% 2.3112% % 362.45% 20.17% 3.2930%!XCH4%=%0.129$$$XCO2%=%0.1301$ % XCH4%=%0.1701$$$XCO2%=%0.1715$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.22% 20.37% 3.4497% % 303.05% 20.64% 5.1344%313.27% 20.52% 3.6798% % 312.41% 20.97% 5.4596%323.29% 20.05% 3.9257% % 322.39% 19.57% 5.8238%333.29% 19.66% 4.1408% % 332.45% 19.49% 6.1541%343.31% 19.81% 4.3509% % 342.46% 19.96% 6.4643%353.28% 19.94% 4.5510% % 352.51% 19.81% 6.7696%363.34% 19.72% 4.7461% % 362.50% 19.94% 7.0498%!

53

BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[phosphinate]"+"(75"mole%CO2"–25"mole%CH4)""

!XCH4%=%0.0812$$$XCO2%=%0.2397$ % XCH4%=%0.1237$$$XCO2%=%0.3653$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.19% 21.67% 1.4918% % 303.32% 21.44% 3.2780%313.08% 21.39% 1.6323% % 313.19% 21.42% 3.5831%323.31% 21.66% 1.7772% % 323.26% 21.44% 3.9482%332.52% 21.67% 1.9173% % 333.19% 21.44% 4.2033%342.08% 22.17% 2.0736% % 343.17% 21.63% 4.5634%351.97% 22.35% 2.2388% % 353.15% 21.95% 4.8535%362.74% 22.02% 2.4428% % 363.16% 22.11% 5.1937%!XCH4%=%0.1530$$$XCO2%=%0.4519$ %T/K% Troom/%°C% Pactual/MPa% %303.31% 20.90% 5.3047% %312.54% 21.03% 5.7801% %322.81% 20.46% 6.3703% %332.57% 20.92% 6.8309% %342.71% 21.63% 7.3958% %352.86% 21.05% 7.8863% %362.97% 21.13% 8.4417% %%BubbleVpoint*pressures*for*the*ternary*system:"""ThtdpG[dicyanamide]"+"(75"mole%CO2"–25"mole%CH4)"!XCH4%=%0.06435$$$XCO2%=%0.190$ % XCH4%=$0.08167$$$$XCO2%=%0.24113$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%303.63% 20.95% 2.1740% % 303.21% 22.17% 3.0291%312.58% 20.92% 2.3542% % 313.16% 22.25% 3.2993%322.58% 20.95% 2.5734% % 323.13% 20.95% 3.5850%332.59% 20.40% 2.7687% % 333.23% 21.02% 3.8601%342.58% 21.63% 2.9595% % 343.20% 20.83% 4.1453%352.61% 22.39% 3.1448% % 353.40% 21.77% 4.4198%362.74% 22.28% 3.3250% % 363.53% 21.68% 4.6850%!XCH4%=%0.10308$$$$XCO2%=%0.30435$ % XCH4%=%0.1277$$$$$$$XCO2%=0.3771$T/K% Troom/%°C% Pactual/MPa% % T/K% Troom/%°C% Pactual/MPa%302.57% 21.97% 4.3086% % 303.19% 21.32% 6.4017%312.44% 22.11% 4.7089% % 313.11% 21.29% 7.0071%322.56% 21.80% 5.1394% % 323.20% 21.91% 7.6623%332.55% 21.12% 5.5297% % 333.23% 21.97% 8.2826%342.56% 22.01% 5.9249% % 343.23% 22.01% 8.8929%352.60% 22.11% 6.3302% % 353.23% 21.97% 9.4833%362.67% 22.00% 6.7155% % 363.29% 22.10% 10.0635%

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