PHASE BEHAVIOR OF CARBON DIOXIDE
AND OXYGEN IN THE IONIC LIQUID 1-HEXYL-3-METHYLIMIDAZOLIUM
BIS(TRIFLUOROMETHYLSULFONYL)IMIDE
A Thesis
Submitted to the Graduate School
of the University of Notre Dame
in Partial Fulfillment of the Requirements
for the Degree of
Master of Science
in Chemical Engineering
by
Katherine E. Wilbanks, B.Ch.E.
Joan Brennecke, Director
Graduate Program in Chemical and Biomolecular Engineering
Notre Dame, Indiana
April 2007
PHASE BEHAVIOR OF CARBON DIOXIDE
AND OXYGEN IN THE IONIC LIQUID 1-HEXYL-3-METHYLIMIDAZOLIUM
BIS(TRIFLUOROMETHYLSULFONYL)IMIDE
Abstract
By
Katherine E. Wilbanks
The phase behavior of carbon dioxide and oxygen mixtures in the ionic liquid 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide,
[hmim][Tf2N], at (40 ± 0.5)˚C was investigated. Experiments were carried out in a fixed volume high pressure view cell with the vapor phase being sampled and analyzed using a gas chromatograph. Solubility data was determined by difference. Measurements were carried out at equilibrium pressures from 18.7 to 72.5 bar and feed compositions of 25,
35, and 50 mole percent oxygen with the balance being carbon dioxide.
Noticeable enhancement was seen at pressures greater than 30 bar and was quantified by calculating enhancement factors and showing ternary diagrams. These results were compared with those of the same system at a lower temperature and carbon dioxide-expanded organic solvents.
For my mother, who gave me roots and wings.
ii
TABLE OF CONTENTS
List of Figures...... v
List of Tables ...... viiii
List of Symbols ...... ix
Acknowledgements ...... xii
Chapter 1: Introduction and Background...... 1 1.1 Introduction...... 1 1.2 Ionic Liquids as Green Engineering Liquids...... 2 1.3 Challenges facing Ionic Liquids as Green Engineering Fluids ...... 5 1.4 Objective of this work...... 7 1.5 Background ...... 7 1.5.1 Motivation ...... 7 1.5.2 Pure Gas Solubility Data ...... 9 1.5.3 Mixed Gas Solubility Background...... 14
Chapter 2: Theory ...... 21 2.1 Introduction...... 21 2.2 Phase Equilibrium in One-Component Systems ...... 21 2.3 Phase Equilibrium in Mixtures...... 24 2.4 Equation of State Selection...... 27 2.5 Calculations ...... 30
Chapter 3: Experimental Methods ...... 33 3.1 Apparatus...... 33 3.1.1 Experimental Configuration...... 33 3.1.2 Gas Chromatograph Configuration ...... 36 3.2 Materials...... 36 3.2.1 Gases...... 36
3.2.2 Ionic Liquid [hmim][Tf2N] ...... 37
iii 3.3 Procedure...... 39 3.3.1 GC Calibration ...... 39 3.3.2 Liquid Volume Calibration ...... 40 3.3.3 Combined Cell and Line Volume Calibration ...... 41
3.3.4 Oxygen-Carbon Dioxide-[hmim][Tf2N] Ternary System...41
Chapter 4: Results & Discussion...... 43 4.1 Introduction...... 43 4.2 Enhancement Factors ...... 43 4.3 Equal Fugacity Plots...... 51 4.4 Ternary Diagrams...... 54
Chapter 5: Conclustions...... 67 5.1 Conclusions ...... 67 5.2 Future Work ...... 68
Appendix A: Interpolating Oxygen Solubility Data ...... 70
Appendix B: Procedure for Mixed Gas Solubility Measurements...... 77
Appendix C: Gas Chromatograph Method Settings ...... 81
Appendix D: Error Analysis...... 84
Appendix E: Experimental Data shown in Fugacity Plots ...... 90
Bibliography ...... 94
iv
LIST OF FIGURES
Figure 1.1. Pure carbon dioxide solubility data collected in the static stoichiometric apparatus, IGA microbalance, and Rubotherm suspension microbalance at 10ºC, 25ºC, 40ºC, 50ºC and 60ºC. ....12
Figure 1.2. Pure oxygen solubility data in [hmim][Tf2N] at 10ºC, 25ºC, and 50ºC taken in the IGA microbalance...... 13
Figure 1.3. Solubility of carbon dixoide in [hmim][Tf2N] in the presence of oxygen with a feed gas composition of oxygen/carbon dioxide = 50/50 [52]...... 17
Figure 1.4. Solubility of oxygen in [hmim][Tf2N] in the presence of carbon dioxide with a feed gas composition of oxygen/carbon dioxide = 50/50 [52]...... 17
Figure 1.5. Phase diagram for nitrogen-carbon dioxide-ethane at 220 K and 4 MPa [53]. Data points represent experimental points from previous experiments [54]...... 19
Figure 1.6. Carbon dioxide-oxygen-methanol system at 40˚C and 30 bar shows little oxygen solubility enhancement [50]...... 20
Figure 3.1. Detailed view of the fixed volume view cell...... 35
Figure 3.2. The structure of the ionic liquid 1-hexyl-3-methylimidazolium
bis(trifluoromethylsulfonyl)imide, [hmim][Tf2N]...... 38
Figure 4.1. Solubility of carbon dioxide in [hmim][Tf2N] as a function of fugacity at (40 ± 0.5)oC...... 52
Figure 4.2. Solubility of oxygen in [hmim][Tf2N] as a function of fugacity at (40 ± 0.5)oC...... 53
v Figure 4.3. CO2 (1) – O2 (2) – [hmim][Tf2N] at (20 ± 2) bar and (40 ± 0.5)˚C...... 56
Figure 4.4. CO2 (1) – O2 (2) – [hmim][Tf2N] at (30 ± 4) bar and (40 ± 0.5)˚C...... 57
Figure 4.5. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 3) bar and (40 ± 0.5)˚C...... 58
Figure 4.6. CO2 (1) – O2 (2) – [hmim][Tf2N] at (50 ± 3) bar and (40 ± 0.5)˚C...... 59
Figure 4.7. CO2 (1) – O2 (2) – [hmim][Tf2N] at (60 ± 5) bar and (40 ± 0.5)˚C...... 60
Figure 4.8. CO2 (1) – O2 (2) – [hmim][Tf2N] at (70 ± 2) bar and (40 ± 0.5)˚C...... 61
Figure 4.9. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C and
pressures from 20 - 70 bar for a 25 mole % O2 feed with the
balance being CO2...... 62
Figure 4.10. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C and
pressures from 20 - 70 bar for a 35 mole % O2 feed with the
balance being CO2 ...... 63
Figure 4.11. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C and
pressures from 20 - 70 bar for a 50 mole % O2 feed with the
balance being CO2...... 64
Figure 4.12. Ternary phase behavior for CO2 (1)- O2 (2)- solvent (3) at (30 ± 4) bar and (40 ± 0.5)˚C where the organic solvents are
acetonitrile and methanol and the IL is [hmim][Tf2N]...... 65
Figure 4.13. Ternary phase behavior for CO2 (1)- O2 (2)- solvent (3) at (50 ± 3) bar and (40 ± 0.5)˚C where the organic solvents are
acetonitrile, acetone and methanol and the IL is [hmim][Tf2N]. ....66
Figure A.1. Solubility of pure oxygen in [hmim][Tf2N]. The experimental points at 10ºC, 25ºC, and 50ºC are shown...... 71
vi Figure A.2. The partial molar enthalpy of solution is the slope of ln H vs. 1 . The uncertainty in each data point is shown...... 75 T
! !
vii
LIST OF TABLES
Table 4.1. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ± 0.5)˚C with a feed composition of 25 mole percent O2 with the
balance CO2 ...... 46
Table 4.2. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ±
0.5)˚C with a feed composition of 35 mole percent O2 with the
balance CO2 ...... 47
Table 4.3. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ±
0.5)˚C with a feed composition of 50 mole percent O2 with the
balance CO2 ...... 48
Table 4.4. Vapor-liquid equilibrium for CO2-O2-solvent at (40 ± 0.5)˚C..49
Table 4.5. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ± 0.5)˚C with varying feed compositions...... 50
Table A.1. Henry's Law Constants as a function of temperature for
[hmim][Tf2N] ...... 73
Table E.1. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ±
0.5)˚C with a feed of 25 mole percent O2 with the balance CO2...91
Table E.2. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ±
0.5)˚C with a feed of 35 mole percent O2 with the balance CO2...92
Table E.3. Vapor-liquid equilibrium for CO2-O2-[hmim][Tf2N] at (40 ±
0.5)˚C with a feed of 50 mole percent O2 with the balance CO2...93
viii
LIST OF SYMBOLS
Special Notation
(underscore as in G) ...... property per mole
(overbar as in Gi ) ...... partial molar property ! !
! ! Symbols
d ...... total derivative symbol
f ...... pure component fugacity ! f i ...... fugacity of a species in a mixture ! G ...... Gibbs free energy
! N ...... number of moles
! P ...... pressure
! R ...... gas constant
! S ...... entropy
! T ...... temperature
! V ...... volume, mixture volume ! x i ...... mole fraction of species i in vapor phase
!
! !
ix Greek Symbols
"...... number of components
" ...... partial derivative symbol
!
! Subscripts
i ...... i th species, i = 1...."
! ! Superscripts !
I ,II ...... phase
IG ...... ideal gas property ! V ,L ...... vapor and liquid phase, respectively !
!
x ACKNOWLEDGEMENTS
It is my pleasure to thank the many people who have helped me in various ways throughout my education and this thesis work.
First, I’d like to thank my advisor, Dr. Joan Brennecke. Without her enthusiasm, knowledge, clarity and compassion, I would not have been able to get through the frustration of puzzling results. Throughout the past few months, her encouragement, sound advice and realistic opinions have been especially useful and appreciated.
I’d also like to thank Dr. Mark McCready and Dr. Ed Maginn for allowing me to be their Teaching Assistant. I thoroughly enjoyed running the recitation sessions and some of my fondest memories of the past two years come from interacting with the undergraduates.
I would be remiss without mentioning the professors from my undergraduate institution, Auburn University. Dr. Steve Duke taught my first undergraduate class and taught me what Chemical Engineering was and why it would be a great career. The love of the field he planted in me kept me going through some hard times and still remains today. Dr. Mario
Eden, Dr. Ram Gupta, Dr. Glennon Maples, and Dave Mills all taught me critical lessons. And lastly, Dr. Chris Roberts has been the most influential and admirable person in my career thus far. Since teaching me sophomore xi year and becoming department chair, I am constantly impressed by his energy, enthusiasm, passion and ability to make time for everyone.
This thesis work could not have been done were it not for the staff here at Notre Dame. I am grateful to the secretaries for deciphering my purchase orders, keeping me informed of pending deadlines and always being patient and understanding. Thank you Karen Jacobs, Marty Nemeth, and Jeanne Davids. I am also grateful to Jim Smith in the machine shop for fixing my apparatus and teaching me general machine knowledge, and
Jim Kerskey for his computer and electrical support.
I’d like to thank the post-docs in the Brennecke research group. Dr.
JaNeille Dixon has been a constant source of information relating to synthesis, purification and just overall knowledge. Dr. Zulema Lopez-
Castillo has helped me extensively with the gas chromatograph and general analysis of the data. I am grateful to both for their patience and help in the lab. I’d also like to thank the graduate students in the
Brennecke research group: Berlyn Mellein, Jes Anderson, Alex Chapeaux,
Luke Simoni, and Lindsay Ficke; and the undergraduate students: Kelsey
Montalto, Yamil Colon and Caitlin Lambert. Additionally, David Couling was a delight to work with in the lab Spring of 2006 and his knowledge of the system and work ethic were greatly appreciated.
I’d also like to thank the following people for their friendliness and fellowship over the past two years: Mike Dimino, Josh Enszer, Diana Hou,
Brian Novak, Prasad Sarangapani, Eric and Amy Smith, Jesse Sullivan,
xii Meredith Pierce, Paul Schramm, Lindsay Seders, Kate Waring, Andrea
Dunn, PJ LeBlanc, Brad Weldon, Shawn O’Brein and Jeremiah White. I’d especially like to thank Jason Gordon for being a constant support system and making ND a home away from home. I’d also like to wish the undergraduate classes of 2008 and 2009 the best of luck as it was a pleasure to get to know them.
Lastly, I’d like to thank my family for still loving me even though
I’ve missed many holidays and events being 700 miles from home. I’m looking forward to moving South.
xiii
CHAPTER 1:
INTRODUCTION AND BACKGROUND
1.1 Introduction
Ionic liquids (ILs) are a growing class of purely ionic organic salts that are liquids at low temperature (below 100˚C). Depending on the cation and anion choice, ILs are tailorable to exhibit properties for the desired
application [1]. [hmim][Tf2N], 1-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide, is the IUPAC standard IL and is therefore the focus of this research. Other ILs may prove more suitable for the application using a different cation such as a pyridinium base or a different anion such as tetrafluoroborate, triflate, or hexafluorophosphate.
Despite the numerous combinations of ILs possible, common properties are seen such as:
• High thermal stability • High ionic conductivity • Negligibly flammable • Negligible vapor pressure • Ability to dissolve a wide range of compounds • Low melting point (below 100˚C) • High heat capacity [2]
1 One key property that is exploited in industry is the negligible volatility. This allows for negligible volume loss due to vaporization which is beneficial for both worker exposure and replenishing the fluid. This property is also a key reason ILs are classified as “Green Engineering
Liquids.”
1.2 Ionic Liquids as Green Engineering Liquids
Ionic liquids have found use in the areas of heat transfer fluids, azeotrope-breaking liquids, the electrochemical industry, supported ionic liquid membranes, plasticizers, dispersants and surfactants, antimicrobial and embalming agents, anticorrosion coatings and electopolishing agents
[3], and especially new uses for reactions. Their potential industrial uses are growing as new ILs are synthesized with varying physical and chemical properties. Here, some of the more popular and interesting applications of
ILs as green engineering liquids will be discussed.
ILs have found particular use in reactions as solvents [4], co- solvents [5] and catalysts [6] bringing greater selectivity, specificity and yield. A key benefit of their use is their ability to be recycled [7]. As a
catalyst, 1-butyl-3-methylimidazolim tetrachloroferrate, [bmim][FeCl4], was found to be air stable for the Grignard reaction and after decanting the product, was recycled four times [8]. Bakos et al. used similar anion
ILs as both the catalyst and solvent and again, after recycling for four cycles saw neither the conversion nor enantioselectivity decreased [9].
2 Jana et al. observed similar results using 1-butyl-3-methylimidazolium hydroxide, [bmim][OH], as a catalyst and reaction media for a
Knoevenagel condensation of aliphatic and aromatic carbonyl compounds
[10]. Another interesting use is as a buffer for controlling the pH in non- aqueous media [11]. Additionally, organic products can be removed from
ILs through another green fluid, supercritical carbon dioxide (scCO2) [12].
Livingston et al. believe the major benefit of using ILs in reactions is the catalyst stability due to the higher viscosity the IL supplies. They used mixed solvents of an IL with a traditional organic compound to increase mass transfer for favorable results [13].
In addition to their use in reactions, many ionic liquids also exhibit unique gas solubility, transport and separation properties [14]. Supported
Ionic Liquid Membranes (SILMs) are growing in their industry use for both gas separations and reactions [15]. SILMs have distinct advantages over conventional Supported Liquid Membranes (SLMs) because the non- volatile viscous nature of ILs prevents evaporations and easy displacement from supporting media [16]. Additionally, enzymes can be suspended in the SILM providing greater stability than the analogous system [17]. SILMs have also been used to separate natural gas as they have exhibited selectivities towards hydrogen sulfide and carbon dioxde
[18].
Ionic liquids are capable of dissolving carbohydrates ranging from simple sugars to polysaccharides. Of particular interest is cellulose, the
3 most abundant renewable resource in the world. Once dissolved, the preparation of cellulose derivates and composites can be preformed.
There is great potential for this technology [19]. Chitin and chitosan, which usually only dissolve in strong acids, have been dissolved in 1-butly-
3-methylimidazolium chloride, [bmim][Cl]. This reversible dissolution allows for their use in cosmetics, drug delivery, heavy metal chelations, heterogeneous catalysis and more. Chitin and chitosan have even been
found to reversibly absorb carbon dioxide (CO2) which yields even more potential applications [20]. Additionally, banana pulp has even been dissolved in [bmim][Cl], to study the nutritional value over time with ripening using 13C NMR studies. The nutritional value is important as bananas are a staple in some tropical climates diet [21]. This carbohydrate dissolution is expected to occur due to the non-hydrated chloride ions present in the solution solvating the carbohydrates by forming hydrogen (H) bonds with their hydroxyl (OH) groups [22]. This interestingly indicates the solubility is influenced heavily by the anion
[23].
In addition to the uses discussed above, ILs have found use in other applications as high-performance liquid electrolytes [24], in solar cells
[25], lithium batteries [26], super capacitors [27], artificial muscles [28], and more [29]. Task-specific ionic liquids are used for both solubilizing and extracting metal ions from aqueous solutions [30]. New ILs with novel functionalities are being synthesized for specific industrial
4 applications and the known physical and chemical properties are being exploited for new applications.
1.3 Challenges facing Ionic Liquids as Green Engineering Fluids
The main challenges affecting ILs as green engineering fluids are availability, cost, purity, compatibility, environmental safety/health issues, and intellectual property [31]. These challenges are faced on a laboratory scale in this research, and must have solutions before expanding to an industrial scale. Under intense scrutiny in the literature this past year has been both the synthesis and purification methods and the environmental safety/health issues.
It is evident that the use of ionic liquids in some industrial applications is green, however, current methods of both synthesis and purification are unacceptable on an industrial scale. Ionic liquid preparation through microwave irradiation is growing, however, increased energy costs need to be considered compared to the conventional heating and longer times of traditional synthesis techniques. One key advantage of the microwave approach is the lack of solvent needed [32]. An additional
synthesis method employs scCO2 as the solvent for the reaction and purification giving a 100 percent yield [33], but again energy costs need to be considered. Aside from alternative synthesis methods, alternative purification methods are also being investigated. ILs physical properties are heavily dependent on their purity [34] and traditional purification
5 techniques involve extensive use of volatile organic compounds and columns. Electrolysis has been used successfully to remove bromide and chloride impurities [35]. These and other techniques will see increasing research as ILs near greater industrial use.
Again, before ILs can be used extensively in an industrial setting, health and safety issues need to be addressed and resolved. ILs are for the most part non-flammable, but they are not necessarily non- combustible. Energetic functionalities have been introduced such as nitric
oxide (NO2), cyanide ion (CN), and nitrogen ions (N3) which when heated with a small flame torch have been found to combust [36]. The IL we are
researching, [hmim][Tf2N], was not part of the study, but it is important to recognize the potential of ILs to combust in an industrial setting.
Additionally, the non-volatility of ILs only protects air quality, but IL disposal and the effect on organisms (anti-microbial properties) and water is still a concern due to some ILs miscibility with water [37]. Both modeling and experiments [38] have been done to analyze this effect.
Due to the increasing number of ILs synthesizable, it is very important to understand and predict the toxicity of these compounds through modeling. QSPR modeling results indicate that toxicity increases with the number of nitrogen atoms in an aromatic cation ring and that the anion plays a secondary role in toxicity [39].
It is apparent that before ILs are adopted by industry, more knowledge about both the synthesis and environmental health and safety
6 concerns must be determined. Additionally, the other challenges such as cost, purity and intellectual property must be addressed.
1.4 Objective of this work
The objective of this work was to investigate the ternary phase
behavior of the CO2–O2–[hmim][Tf2N] system. The IL chosen is the IUPAC standard. The system was studied from pressures of 18.7 to 72.5 bar at
(40 ± 0.5)ºC. O2 mole fractions were less than 0.50 as the focus was on
CO2 enhancement ability.
1.5 Background
1.5.1 Motivation
Understanding the solubility of CO2 and O2 in an IL is fundamental for two applications: IL as a solvent for a reaction and IL as a gas separation agent. These applications and background data on pure and mixed gas solubility in ILs will be discussed.
Using ILs as solvents for reactions is desired for two key reasons: the non-volatility of the IL prevents loss and the recycling and recovery of the IL allows the solvent to be used repeatedly without losing its desirable physical properties such as catalyst suspension. Common reaction gases
such as O2, hydrogen (H2) and carbon monoxide (CO) have low solubilities
in ionic liquids. It is hopeful that the presence of a high solubility gas, CO2, 7 will enhance the solubility of the reaction gas, in this work O2. This theory
has been tested in the literature for H2. Because of its low solubility and
molecular weight, reactions that occur in the presence of H2 were used as
probes. Leitner et al. found that their CO2-H2-IL system exhibited a
remarkable deviation from Henry’s law which predicts the H2 solubility in
the liquid solvent be only a function of the H2 partial pressure. In their experiments, the partial pressure was fixed, but the solubility increased as
it was diluted further with CO2 which lead to a higher reaction rate [40].
These results are expected to translate to other reaction gases. However,
CO2 is volatile, non-viscous, non-conduction, non-polar and unable to dissolve large and unsaturated compounds. Many ILs have exactly the opposite properties. Therefore, reaction performance is expected to differ between these two solvents and especially between different ILs [41].
As a separation agent, the varying solubilities (some orders of magnitude) of different gases can be exploited to selectively separate gases. This is especially popular in membrane systems where the IL is in a more stationary phase. Research is underway to determine how difficult it will be to remove the selectively solvated gas once the separation has occurred [42].
Essentially, if CO2 enhances O2 solubility well, the system would be better suited for reaction use. While, if there is little to no enhancement, a good gas separation could be made.
8 1.5.2 Pure Gas Solubility Data
Many trends on pure gas solubility data for CO2 can be seen in the literature. The two most basic are as pressure increases, solubility increases and as temperature increases, solubility decreases. Additionally, there is a greater anion affect than cation and increasing the alkyl chain length on the cation increases the solubility [43].
Pure solubility data of CO2 data in [hmim][Tf2N] at 25ºC, 40ºC and
60ºC has been collected in our laboratory in a static stoichiometric apparatus [44]. Additionally, data is available at 10ºC, 25ºC, and 50ºC from the IGA microbalance apparatus, and at 25ºC in the Rubotherm microbalance apparatus [45]. This data is shown in Figure 1.1. The static stoichiometric apparatus is a P-V-T apparatus consisting of glass equilibrium cells in which the pressure and temperature can be monitored.
It functions similarly to the experimental set-up used in this work and is discussed in detail elsewhere [46]. The IGA microbalance is a gravimetric microbalance (IGA 003, Hiden Analytical), details of which are described elsewhere [47]. The microbalance consists of a sample pan and a counterweight and needs only a small sample, about 75 mg. The balance has a 1 µg stable resolution and equilibrium was determined when the weight remained constant for at least 15 minutes. The measurements were! corrected for buoyancy using accurate equations of state. The
Rubotherm microbalance (gravimetric microbalance suspension balance,
Rubotherm), functions similarly to the IGA in that it is a gravimetric mass
9 balance, however, it has a larger sample size and associated equilibrium time. The key benefit of the Rubotherm microbalance is that the microbalance is sealed from the gas of interest allowing more noxious gases to be used. Details are found elsewhere [42].
In order to have a value at each pressure under study in this research, a polynomial is fit to the 40ºC data where y is the gas mole fraction and x is the pressure (bar).
"5 2 y = "9x10 x + (0.0163± 0.0004)x 1.1
This equation is found by applying a trendline in Microsoft Excel. The error in the second! term is derived from the error in the mole fraction. To find this uncertainty, three data sets were used based on the mole fraction data: a predicted value, higher case, and lower case. These data sets were then plotted against the pressure and trendlines compared to yield
Equation 1.1. The error is only in the second term due to the magnitude of the values.
O2 has not been investigated as extensively in the literature as CO2 leaving it with some uncertainty. Similar trends of increasing solubility
with decreasing temperature and increasing pressure are seen with O2 in
1-butyl-3-methylimidazolium hexafluorophoshate, [bmim][PF6] [48] and
1-butyl-3-methylimidazolium tetrafluoroborate, [bmim][BF4] [49].
However, the IL in this study, [hmim][Tf2N] behaves differently [50].
Data at 40˚C is not available and therefore must be interpolated.
There are two methods to interpolate the data that yield similar results,
10 discussed in detail in the appendix. The equation generated where y is the
O2 mole fraction and x is the pressure (bar) is
y = (0.0010 ± 0.0003)x 1.3
The error in the equation comes from fitting a linear trendline in excel.
The pressure is ! plotted against the mole fraction for three cases: the expected value, the higher uncertainty, and the lower uncertainty. The resulting fits are visually inspected to yield the uncertainty.
It is important to note that data was only available in the linear region, 0 - 15 bar, and therefore there is error associated with using this correlation beyond that region.
11
0.8
0.6
0.4
10 oC IGA Mole Fraction 2 25 oC IGA o CO 25 C Rubotherm 0.2 25 oC Kohn 40 oC Kohn 50 oC IGA 60 oC Kohn 0.0 0 20 40 60 80 100 120 140
Pressure (bar)
Figure 1.1. Pure carbon dioxide solubility data collected in the static stoichiometric apparatus, IGA microbalance, and Rubotherm suspension microbalance at 10ºC, 25ºC, 40ºC, 50ºC and 60ºC.
12
0.020
0.015
0.010 Mole Fraction 2 O 0.005 10 oC 25 oC 50 oC
0.000 0 2 4 6 8 10 12 14
Pressure (bar)
Figure 1.2. Pure oxygen solubility data in [hmim][Tf2N] at 10ºC, 25ºC, and 50ºC taken in the IGA microbalance.
13 1.5.3 Mixed Gas Solubility Background
CO2 enhancing the solubility of other gases in liquids is not a new area, however, experimental studies of this phenomena in ILs are. To show the desired effect and discuss previous results, mixed gas solubilities in conventional organic solvents will be the focus of this section with some discussion of how it compares to IL data. The three methods presented in the results discussion for quantifying enhancement will be used as examples.
The first method of quantifying enhancement is by calculating an enhancement factor. The enhancement factor is defined as the ratio of solubility in the mixture to the solubility of the pure gas at the same fugacity and temperature. Lopez-Castillo et al. found enhancement
factors for O2 in the presence of CO2 at 40˚C and pressures of 6 – 82 bar ranging from 0.77 - 1.31 in acetonitrile, 0.77 - 1.08 in acetone, and 0.98
- 1.27 in methanol. A value of 1 indicates no enhancement making it
apparent the CO2 provided only a modest increase in solubility [51].
Alternatively, Hert et al. found enhancement factors for O2 in the
presence of CO2 at (25 ± 0.5)˚C and pressures of 3.0 – 15.7 bar ranging
from 1.7 - 5.7 in [hmim][Tf2N] [52].
The second method of showing enhancement is essentially graphically showing enhancement factors. As evident in Figure 1.3, Hert
et al. found that rather than CO2 keeping its high solubility, it was
14 negatively effected by O2, lowering its solubility in [hmim][Tf2N] [52].
However, Figure 1.4 shows there was an increase in the O2 solubility in the IL. These results were surprising as they indicate that oxygen is competing with carbon dioxide for absorption sites or hindering the absorbance in some other way. It is believed that due to poor temperature control and systematic error more investigation into this system is needed at a broader pressure range. Considerable improvements of the experimental apparatus have also been made yielding better precision and accuracy of results.
15
0.35
0.30
0.25
0.20
0.15 Mole Fraction 2 0.10 CO
0.05 pure mixture
0.00 0 2 4 6 8 10 12 14 CO Fugacity (bar) 2
Figure 1.3. Solubility of carbon dioxide in [hmim][Tf2N] in the presence of oxygen with a feed gas composition of oxygen/carbon dioxide = 50/50 [52].
16
0.05
0.04
0.03
Mole Fraction 0.02 2 O
0.01 pure mixture 0.00 0 2 4 6 8 10 12 14 O Fugacity (bar) 2
Figure 1.4. Solubility of oxygen in [hmim][Tf2N] in the presence of carbon dioxide with a feed gas composition of oxygen/carbon dioxide = 50/50 [52].
17 The third and preferred method of showing the mutual effect of mixed gas solubilities is by graphically showing vapor-liquid ternary
behavior. Figure 1.5 shows an example of strong enhancement in a CO2-
nitrogen (N2)-ethane (C2H6) mixture. The dew line curving inward represents an enhancement of the nitrogen solubility in the ethane solvent. With no enhancement, the dew line would be linear. A more
representative system of the one under study is CO2- O2-methanol
(MeOH). Figure 1.6 is the same system that yielded enhancement factors of 0.98-1.27 [51]. It is evident from the ternary phase diagram there is little, if any, solubility enhancement of the oxygen. These two examples presented represent different temperatures, pressure, phases and solvents so should not be compared qualitatively.
It has been shown that CO2 can provide moderate enhancement of lower solubility gases in some conventional organic solvents and ILs. This has great potential for using ILs as the solvent for reactions such as oxidations. However, this increase in solubility generally comes at the cost of higher overall pressures, but potentially reduces flammability hazards and lowers viscosity for easier pumping. This enhancement can be shown by three ways: quantifying enhancement factors, equal fugacity plots, and ternary diagrams.
18
Figure 1.5. Phase diagram for nitrogen-carbon dioxide-ethane at 220 K and 4 MPa [53]. Data points represent experimental points from previous experiments [54].
19
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 MeOH (3) 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 CO (1) 2
Figure 1.6. Carbon dioxide-oxygen-methanol system at 40˚C and 30 bar shows little oxygen solubility enhancement. The enhancement factor in this system is 1.10 [50].
20
CHAPTER 2:
THEORY
2.1 Introduction
We will begin the theory section with a discussion of phase equilibrium in a one-component species. This will further be expanded to apply to mixtures. Equation of state selection and the calculation and data analysis will then be discussed.
2.2 Phase Equilibrium in One-Component Systems
. The second law of thermodynamics tells us that S gen " 0, with the equality being at equilibrium. Based on this premise, three conclusions can be reached: !
T I = T II 2.1
P I = P II 2.2 ! G I = G II 2.3 ! The following derivation will start from the equality of the Gibbs ! free energy in the liquid and vapor phases, 21 L V G (T,P) = G (T,P) 2.4
where by definition ! d G = "SdT +VdP 2.5
we can then take partial derivatives such that ! # "G & % ( = )S 2.6 "T $ 'P
and ! # "G & % ( = V 2.7 "P $ 'T
We will assume we have an equation of state from which we can compute
V as a function of T and! P .
Integrating Equation 2.7,
! ! ! P 2 2.8 G(T1,P2) "G(T1,P1) = # VdP P1
Using the ideal gas for illustrative purposes, we know ! IG RT V = 2.9 P
so that ! IG IG P2 RT G (T1,P2) "G (T1,P1) = # dP 2.10 P1 P
Subtracting Equation 2.10 from Equation 2.8 yields
! ) , # IG & # IG & P2 RT 2.11 %G (T1,P2) "G (T1,P2)( " %G (T1,P1) "G (T1,P1)( = / +V " .dP $ ' $ ' P1 * P -
To simplify, we set P 1 = 0 and recognize these are ideal gases so that ! 22 ! IG G (T1,P = 0) = G (T1,P = 0) 2.12
Removing the subscripts gives ! IG P # RT & G(T,P) "G (T,P) = %V " (dP 2.13 )O $ P '
We can now define the thermodynamic function, fugacity, f : ! # IG ' # ' % G(T,P) "G (T,P)% % 1 P * RT - % f = P exp$ ( = P exp$ ,V " /dP ( 2.14 0O % RT % &% RT ! + P . )% & )
and the fugacity coefficient, ": ! $ IG ( $ ( f & G(T,P) #G (T,P)& & 1 P + RT . & " = = exp% ) = exp% -V # 0dP ) 2.15 1O P !& RT & '& RT , P / *& ' *
Fugacity is important because of its relation to the Gibbs free
energy! in phase equilibrium calculations. We know from Equation 2.3 that
GI = GII with constant temperature and pressure in both phases. Using
this equality and the definition of fugacity given in Equation 2.14 gives
! I II IG f (T,P) IG f (T,P) G (T,P) + RT ln = G (T,P) + RT ln 2.16 P P
Canceling out like terms leaves
! I II f (T,P) = f (T,P) 2.17
and in terms of the fugacity coefficient
! I II " (T,P) = " (T,P) 2.18
Because these equations are derived directly from the equality of
the molar Gibbs energy! in each phase at phase equilibrium, Equations
2.17 and 2.18 can also be used as criteria for equilibrium [55]. 23 2.3 Phase Equilibrium in Mixtures
The analysis presented above for a pure component species will be further expanded to apply to a mixture. We will start with the definition of the Gibbs free energy in a mixture
# 2.19 dG = "SdT +VdP + $Gi dNi i =1 and use the communitive properties of the second derivatives of the thermodynamic functions!
" $ "G ' " $ "G ' & ) = & ) 2.20 "Ni % "T ( "T % "Ni ( T ,P ,N j #i P ,N j P ,N j T ,P ,N j #i and ! " $ "G ' " $ "G ' & ) = & ) 2.21 "Ni % "P ( "P % "Ni ( T ,P ,N j #i T ,N j T ,N j T ,P ,N j #i which yields two equations
! $ ' #G i S i & ) 2.22 = "& ) % #T ( P ,N j and
! # & "G i V i % ( 2.23 = % ( $ "P ' T ,N j
This gives us a form similar to the pure component
! P2 i i i 2.24 G (T1,P2, x) "G (T1,P1, x) = # V dP P1
The fugacity of species i in the mixture can then be defined by following the same progression! as Equations 2.9 - 2.14.
! 24 IGM # ' IGM % G i (T,P, x) "G i (T,P, x)% # 1 P* - ' f i (T,P, x) = xi P exp$ ( = xi P exp$ 0 ,V i "V i /dP ( % RT % & RT 0 + . ) & ) 2.25
The fugacity coefficient of species i in the mixture is then defined as !
IGM $ ( IGM f i &G i (T,P, x) #G i (T,P, x)& $ 1 P+ . ( " = = exp% ) = exp% V i #V i dP) i ! 10 - 0 xi P & RT & ' RT , / * ' * 2.26
We now want to develop a relationship between the fugacity of the ! pure component i to the fugacity of component i in a mixture. For a pure
component in a mixture, we know from Equations 2.7 and differentiating ! ! the natural log of Equation 2.14 with respect to pressure at a constant
temperature and composition that
# "G & # " lnf & % ( = V = RT% ( 2.27 "P dP $ 'T $ 'T
Similarly, for a species in a mixture
! # & # & "G i " lnf i % ( = V i = RT% ( 2.28 % "P ( % dP ( $ 'T ,x $ 'T ,x
To develop a relationship between the pure component i to the ! fugacity of component i in the mixture, we subtract Equation 2.27 from ! 2.28 and integrate between P = 0 and the pressure of interest, P. ! # ' # ' P % f i (T,P, x) % fi (T,P) RT ln$ ( * RT ln$ ( = + V i *V i dP 2.29 % % f (T,P " 0) P "0 ( ) &f i (T,P " 0, x)) & i )
!
25 Simplifying
" & $f i (T,P, x)$ P RT ln# ' = * V i )V i dP 2.30 $ x f (T,P) $ 0 ( ) % i i (
This tells us that for a mixture in which V i = V i , the fugacity of ! each species in the mixture is equal to the mole fraction times the pure ! component fugacity at the same temperature and pressure. If V i " V i ,
then f i and f i are related through the integral over all pressures of the ! difference between the species partial molar and pure component molar ! volumes.!
We would also like to develop an additional criterion for phase
equilibrium as well as an expression for the change in partial molar Gibbs
free energy at different compositions but the same temperature and
pressure. We will begin by stating
II I " G i = G i (T,P, x ) #G i (T,P, x ) 2.31
Substituting in the natural log of Equation 2.25 and recognizing that
! IGM IG
"G i T,P, x = G i (T,P) + RT ln x 2.32 ( ) i
through some algebraic manipulation, we arrive at
! # II ' II I % f i (T,P, x )% G i (T,P, x ) G i (T,P, x ) RT ln 2.33 " = $ I ( % % & f i (T,P, x ))
when state I is a pure component, this simplifies further to
! # ' % f i (T,P, x)% G i (T,P, x) "G i (T,P) = RT ln$ ( 2.34 ! % f (T,P) % & i ) 26
! This yields an additional criterion for phase equilibrium [55]
I " I % " II % II f i = f i $T ,P, x ' = f i $T ,P, x ' = f i 2.35 # & # &
The fugacity is needed for the gas mixtures to characterize the vapor phase ! above the IL. Due to the non-volatility and general lack of data related to the IL phase, data analysis is based on the changing number of moles in the gas phase even though we are characterizing the
IL rich liquid phase. The fugacity will be calculated for each gas and will
serve as a benchmark for enhancement. For example, pure O2 with a
fugacity of 5 bar will be compared to a mixture of 25 mole percent O2 with a fugacity of 5 bar but with a total mixture pressure of 30 bar.
2.4 Equation of State Selection
The central functions of the equation of state is to generate a calibration between the Gas Chromatograph (used to analyze the vapor phase) Area Counts and the number of moles of gas in the sample loop
(procedure discussed in Section 3.3.2) and to determine the fugacity of each gas species. Additionally, an equation of state is used as a check to insure the volumetric method of data analysis is correct (discussed in section 2.5).
The Peng-Robinson equation of state has been chosen due to its improved liquid densities as well as accurate vapor pressures and
27 equilibrium ratios over other equations that are also modifications of the
van der Waals equation [56].
The Peng-Robinson equation is a semi-empirical equation of the form
P = PR + PA 2.36
where P R is the repulsion pressure and P A is the attraction pressure. The ! repulsion pressure term is given by ! ! RT PR = 2.37 v " b
which is the same van der Waals hard sphere equation used in the Redlich-
Kwong Equation [57]. ! Where the Peng-Robinson equation differs is in
their attraction term, PA . They include a predicted critical compressibility
factor, the acentric factor and the reduced temperature. This gives an
equation of the! form
RT a(T) P = " 2.38 v " b v(v + b) + b(v " b)
Equation 2.38 can be rewritten as a cubic equation of state using
the compressibility! factor
3 2 2 2 3 Z " (1" B)Z + (A " 3B " 2B)Z " (AB " B " B ) = 0 2.39
where ! aP A = 2.40 R 2T 2
bP B = 2.41 ! RT
! 28 Pv Z = 2.42 RT
The Peng-Robinson equation can be further applied to mixtures by
using the following parameters!
a = "" xi x j aij 2.43 i j
where ! a = 1" # a1/2a1/2 2.44 ij ( ij ) i j
and is an empirically determined binary interaction coefficient found in " ij ! the literature and ! b = " xi bi 2.45 i
A spreadsheet developed by Lira [58] employing this equation of state
was used for the calculations.!
It is important to note that due to the non-volatility of ILs the
equation of state will be used only for the vapor phase which is a mixture
of only carbon dioxide and oxygen. These two gases are widely
characterized in the literature and parameters are available. The
empirically determined binary interaction coefficient, , is more widely " ij
depicted as k and is taken from the literature to be -0.04838 at 40˚C ij ! [53].
!
29 2.5 Calculations
At this point we have developed criteria for equilibrium and chosen the Peng-Robinson equation of state for characterizing the vapor phase.
For our criterion to be true, we must assume that at least one molecule of the ionic liquid is in the vapor phase. However, we know the vapor pressure of the ionic liquid is so close to zero, we consider this contribution negligible. This also prevents the ionic liquid rich phase from running through the gas chromatograph. For these reasons, our calculations are based on sampling the vapor phase above the ionic liquid for oxygen and carbon dioxide composition. This method is known more commonly as the stoichiometric method.
In order to determine the composition in the vapor phase, or headspace, we must first know the volume of the headspace. The vapor phase encompasses not only the space inside the viewing cell, but also the tubing around the cell. For this reason, we define the vapor space as the volume of the liquid subtracted from the total volume of the cell
V V V 2.46 headspace = total " liquid where the total volume, V total , of the cell is a constant (procedure ! discussed in section 3.3.3) and the liquid volume, V , in the cell liquid ! changes, expanding, as the experiment progresses. The liquid volume is found by measuring the height of the liquid ! in the view cell with the
30 cathetometer and correlating that value to the volume (correlation
discussed in section 3.3.2).
Once the volume of the headspace, V , is known, we need to headspace
find the number of moles of each component i in the headspace. The
vapor space can be sampled and analyzed! using the gas chromatograph ! which gives us the number of moles of each component in the sample
loop, N . The number of moles in the sample loop is then proportionally i ,loop
scaled to the number of moles in the headspace: ! V N = N " head 2.47 i ,headspace i ,loop V loop
The volume of the headspace, V , and the number of moles in headspace ! the headspace, N , need to be found at both the initial and i ,headspace ! equilibrium cell conditions. The number of moles absorbed into the ionic
liquid is simply! the difference of the two values.
N N initial Nequilibrium 2.48 i ,absorbed = i ,headspace " i ,headspace
Once we know the number of moles of each component absorbed,
we can calculate! the mole fractions in the liquid phase. The number of
moles of the carbon dioxide and oxygen absorbed will change each
experiment, however, the number of moles of ionic liquid remains
constant and is calculated from the mass added to the cell initially. No
ionic liquid is lost over time due to its negligible vapor pressure and
31 resistance to degradation. The mole fraction of component i in the mixture is calculated by: ! Ni ,absorbed xi = 2.49 "N j ,absorbed + Nionic liquid j =1
As we are using mixtures of gases, it is desirable to reference component fugacities! in addition to total system pressures. Equation
2.15 gives us the fugacity coefficient equation using the ideal gas equation, however we wish to use the Peng-Robinson equation of state.
Expressed in terms of the fugacity coefficient [57]
$ ' A Z + (1+ 2)B ln " = #ln(Z # B) # ln& ) + Z #1 2.50 & ) B 8 % Z + (1# 2)B (
At a given pressure and temperature the cubic equation 2.39 is ! solved for Z resulting in ", the fugacity coefficient. Again using the pressure, this yields the fugacity in the pure component case, ! ! f = "P 2.51 and for each component in the mixture, ! f i = yi "i P 2.52
We now have all the calculations outlined to properly analyze the data to construct ternary! diagrams visually showing the results and calculate enhancement factors to quantify the results.
Error propagation is discussed in the appendix.
32 CHAPTER 3:
EXPERIMENTAL METHODS
3.1 Apparatus
3.1.1 Experimental Configuration
The high pressure fixed volume view cell (FVVC) shown in Figure
3.1 was the sole experimental cell used in this work. The feed gas mixture entered the cell through valve 1. This feed gas came directly from the tank for lower pressures or from a syringe pump (ISCO 260D) equilibrated to 40˚C for the higher pressures. The cell was vented at the conclusion of an experiment through valve 2. Valve 3 allowed for nitrogen to bubble through the ionic liquid to remove trace gases between experiments (this valve could alternately be used for liquid phase sampling to the Gas
Chromatograph). The pressure of the cell was monitored by a gauge (Cole
Parmer 94785-00) with an accuracy of ± 2 psi connected to a pressure transmitter (Setra C206) and the temperature by a type K thermocouple
(Omega Engineering KMTSS-125-G-6) connected to a display meter
(Omega DP462). The vapor phase was sampled through valve 4 which led to multiport sampling valves and into the GC. A vacuum pump (Pfeiffer D-
33 35614 Asslar) was located after valve 5 that evacuated the vapor space and cell lines to below 4.0 x 10-3 mbar in between experiments.
The FVVC was made from SS316 and includes a borosilicate glass window to view the liquid level height. A specially designed Teflon stir bar was used to ensure mixing of both the vapor and liquid phases and was controlled by the magnetic stir plate.
The FVVC was surrounded by an air bath held at (40 ± 0.5)˚C. Air was circulated by 2 - 4 desk fans inside the air bath and the heating source was a 600 Watt light bulb controlled by a temperature controller
(Omega DP462). There were temperature gradients throughout the airbath, however, the FVVC occupied only about a quarter of the space, and with proper fan placement the temperature could be stabilized. The gas chromatograph (Varian 3800) used to analyze the gas was outside the air bath.
34
Figure 3.1. Detailed view of the fixed volume view cell.
35 3.1.2 Gas Chromatograph Configuration
The configuration of the GC parameters is crucial to ensuring the calibrations are performed at the same parameters of the experiment.
These parameters are changed in the method editor on the computer and are detailed in the appendix exactly as they appear on the screen. The column used was an AllTech HayeSep Q 80/100 stainless steel column measuring 6’ x 0.125” x 0.085”.
3.2 Materials
The materials used in this study are as follows.
3.2.1 Gases
The gases used in this study were all purchased from Mittler Supply
Inc. The list including grade and minimum purity is as follows:
• Carbon dioxide gas – Bone dry grade - 99.8 % minimum purity • Oxygen/carbon dioxide gas mixture – Custom grade (50 ± 2)% carbon dioxide • Oxygen/carbon dioxide gas mixture – Custom grade 25.11% oxygen • Nitrogen gas – Extra dry - 99.97 % minimum purity • Helium gas - High purity - 99.995 % minimum purity
36 3.2.2 Ionic Liquid [hmim][Tf2N]
The chemicals used in the synthesis of [hmim][Tf2N], including source, grade, and purifcation method (if any) are as follows:
• 1-Bromohexane – Aldrich – 98.0% minimum purity, redistilled • 1-methylimidazole – Aldrich – 99.0% minimum purity, redistilled over KOH • lithium bis(trifluoromethylsulfonyl)imide – 3 M – 97% minimum purity • Acetonitrile – Fisher – 99.9% minimum purity • Methylene chloride - Fisher – 99.9 % minimum purity
[hmim][Tf2N] was synthesized in our lab according to standard procedure detailed stepwise below [59].
1. Mix equal molar amounts of a nitrogen base (1-methylimidazole) and alkyl halide (1-Bromohexane) in a flask and allow to stir overnight with acetonitrile as a solvent. 2. Remove the solvent from the intermediate ionic liquid under vacuum. 3. Purify the IL, [hmim][Br], by washing with water. 4. Mix equal molar amounts of [hmim][Br] with lithium
bis(trifluoromethylsylfonyl)imide, [Li][Tf2N], over low heat allowing an anion exchange to occur producing the desired ionic
liquid, [hmim][Tf2N] 5. Dissolve the IL in methylene chloride and stir over activated charcoal to remove any coloured impurities. 6. Filter the solution removing the solvent in vacuo.
The identity of the IL was confirmed by 1H and 13C NMR. Impurity levels of the halide ions in the IL were measured using an Oakton Ion 510 meter with Cole-Parmer Ion Specific Probes yielding a bromide content less than 10 ppm. NMR results indicate that amine impurities were below
37 the detection limit with is roughly 5 wt%, however, the value is likely significantly less than this. The water content was measured with the
Brinkmann Metrohm Karl Fisher Coulometer and was 35 ppm at the start of the experiment. The IL was exposed to the atmosphere on several occasions, however, it was also exposed to the vacuum pump so this value is expected to have remained somewhat constant.
Figure 3.2. The structure of the ionic liquid 1-hexyl-3- methylimidazolium bis(trifluormethylsulfonyl)imide,
[hmim][Tf2N].
38 3.3 Procedure
3.3.1 GC Calibration
The GC was calibrated for CO2 and O2 using the pure gases. The volume of the sample loop, the temperature of the gas, and the pressure of the gas are known. These values can be inserted into the Peng-
Robinson equation of state [56] (equation of state selection discussed in
Section 2.4) in order to calculate the number of moles in the sample loop.
When the gas is sent from the sample loop through the GC, a peak is produced with a given peak area. This measurement is repeated at different pressures to generate a calibration curve of moles of the sample gas (determined from equation of state) and the peak area from the GC.
The calibrations generated are as follows where y is the moles of the gas and x is the GC area counts: ! For CO2: ! y = 3.62x10"10 ± 3x10"12 x 3.1 ( )
For O2:
! "10 y = 4.53x10 x 3.2
The least squares method of error analysis is discussed in the appendix
! -24 and the uncertainty in the O2 measurement is negligible (10 order of magnitude).
39
3.3.2 Liquid Volume Calibration
The volume of the liquid in the cell as a function of height is needed for determining both the liquid volume as well as the headspace volume. It is not possible to measure the volume in situ, so a cathetometer with
0.005 mm resolution (Gaertner Scientific Corporation 1417A) is used to measure the height.
To generate a calibration curve, the cell was disconnected from the
air bath and the N2 purge line was plugged. The thermocouple was left in place. A buret of tetradecane, the calibration fluid of choice due to it’s high surface wettability, was placed above the cell and known, variable amounts were added to the cell on the order of 0.5 mL. The liquid height change was recorded after each addition using a cathetometer thus generating a calibration curve. The cell was then disassembled, cleaned with a mixture of acetone and ethanol and allowed to air dry. The procedure was repeated to confirm results.
The calibration curve generated where y is the volume in mL and x is the liquid height in cm using the least squares method of error analysis is: ! ! y = (1.573± 0.001)x 3.3
40 !
3.3.3 Combined Cell and Line Volume Calibration
The combined cell and line volume is used to determine the volume of the headspace. The volume of the headspace is the combined cell and line volume minus the liquid volume calculated by the method discussed in the previous section.
To measure the combined cell and line volume, the cell was connected to the air bath and pressure tested. The cell was then vented and carbon dioxide from an Isco Syringe Pump was loaded into the cell and surrounding lines. The volume change in the pump is the total volume of the cell and lines. This procedure was repeated to confirm the results and yielded a volume of 19.8 mL.
3.3.4 Oxygen-Carbon Dioxide-[hmim][Tf2N] Ternary System
Using the apparatus and calibrations discussed above, measurements are taken using the stoichiometric technique. The advantage of this technique is that it is simple, however, it requires a larger sample size and great care has to be taken in removing all gases dissolved in the IL sample initially [43]. The key source of error with this technique lies in the gas adsorption on the metal surfaces of the FVVC and the small length of tubing not stirred which must be cleared before taking data.
41 For each measurement, the cell is held under vacuum (4.0 x 10-3 mbar) at 60˚C for at 24-36 hours to remove all traces of gas left over
from previous measurements. N2 is bubbled through the bottom to increase agitation while removing the residual gas. Once the cell is evacuated and the temperature in the cell and loop is constant at (40 ±
0.5)˚C, the ambient pressure and liquid level (volume) are measured and recorded. The cell is then charged with a known pressure of carbon dioxide-oxygen mixture and the gas mixture in the head space is sampled using the GC. This sampling must be done quickly to minimize the amount of gas dissolving into the ionic liquid. The GC gives peak areas which are then converted to moles of each gas using the calibrations discussed previously.
The stir bar is turned on and 12-72 hours are allowed for the cell to reach equilibrium which is determined by the pressure in the cell no longer dropping. Once equilibrium is reached the headspace is sampled again with the GC. The peak areas are converted to moles and the amount of each gas in the ionic liquid is calculated by difference.
Step-by-step detailed procedures are found in the appendix.
42 CHAPTER 4:
RESULTS & DISCUSSION
4.1 Introduction
Measurements of the solubility of CO2/O2 mixtures in [hmim][Tf2N] were taken at (40 ± 0.5)˚C and pressures from 18.7 to 72.5 bar. To analyze the data, three methods of characterization are used. The first is by quantifying an enhancement factor, the second by graphically comparing gas solubility at the same fugacity, and the third by showing tie-lines on ternary diagrams at similar total pressure will be shown and discussed here.
4.2 Enhancement Factors
One method of comparing the solubility of each gas in the mixture and its pure gas solubility under the same conditions is with the enhancement factor (EF). The enhancement factor is defined as the ratio of solubility in the mixture to the solubility of the pure gas at the same fugacity and temperature defined by Equation 4.1.
43 xmixture EF = gas 4.1 x pure gas
An enhancement factor equal (or very close) to 1 indicates the ! mixture of the gas does not affect the solubility. Due to pure CO2’s higher
solubility in [hmim][Tf2N] than O2, it is expected that oxygen will not
enhance the CO2 solubility. The converse is true for O2. It is expected that
the presence of CO2 will increase the solubility of oxygen giving enhancement factors greater than or equal to (≥) 1.
Tables 4.1, 4.2, and 4.3 show the enhancement factors for 25, 35,
and 50 mole percent O2 mixtures with the balance being CO2. It is evident
that enhancement of O2 is seen at all compositions and fugacities. This is
likely due to the dissolving of non-polar CO2 making the IL solution a more
accommodating environment for the non-polar O2.
Table 4.4 shows a comparison of organics (acetonitrile, methanol,
and acetone) at similar vapor phase composition and O2 fugacity [50]. It is apparent the IL experiences more significant enhancement than the
organics. For example, at a CO2 vapor phase composition of ~0.66 and O2
fugactiy of ~11.0 bar, the enhancement factor of O2 in the IL is 1.39 ±
0.10, while in the organic (acetonitrile) it is just 0.99. An even greater
difference is seen at higher pressures. At a CO2 vapor phase composition
of ~0.38 and O2 fugacity of ~40 bar, the enhancement factor of the O2 in the IL is 3.23 ± 0.23, compared to a value in an organic (acetonitrile) of 0f just 1.08. Similar results are also seen in the table for methanol.
44 Table 4.5 confirms the qualitative trend that at higher oxygen fugacities, there is a higher enhancement factor. The data presented in
Tables 4.1-3 are combined into one and sorted by O2 fugacity. The enhancement factors range from (1.28 ± 0.09) – (3.87 ± 0.27).
45
TABLE 4.1
VAPOR-LIQUID EQUILIBRIUM FOR CO2-O2-[hmim][Tf2N] AT (40 ± 0.5)˚C
WITH A FEED COMPOSITION OF 25 MOLE PERCENT O2 WITH THE BALANCE
CO2.
liquid phase vapor phase composition composition xO2 pure
Ptotal fO2 gas mole
(bar) xCO2 xO2 yCO2 yO2 (bar) fraction EF 20.0 0.275 0.009 0.654 0.346 6.8 0.007 1.28 ± 0.09 33.5 0.368 0.015 0.665 0.335 11.0 0.011 1.39 ± 0.10 41.7 0.429 0.021 0.661 0.339 13.7 0.014 1.56 ± 0.11 42.5 0.404 0.022 0.676 0.324 13.4 0.013 1.66 ± 0.12 48.2 0.462 0.031 0.669 0.331 15.4 0.015 2.02 ± 0.14 57.2 0.455 0.054 0.694 0.306 16.8 0.017 3.21 ± 0.23 58.7 0.462 0.049 0.691 0.309 17.4 0.017 2.79 ± 0.20 69.8 0.611 0.087 0.662 0.338 22.5 0.023 3.87 ± 0.27
46
TABLE 4.2
VAPOR-LIQUID EQUILIBRIUM FOR CO2-O2-[hmim][Tf2N] AT (40 ± 0.5)˚C
WITH A FEED COMPOSITION OF 35 MOLE PERCENT O2 WITH THE BALANCE
CO2.
liquid phase vapor phase composition composition xO2 pure
Ptotal fO2 gas mole
(bar) xCO2 xO2 yCO2 yO2 (bar) fraction EF 19.5 0.223 0.012 0.532 0.468 9.0 0.009 1.29 ± 0.09 31.2 0.328 0.022 0.535 0.465 14.2 0.014 1.57 ± 0.11 42.1 0.385 0.028 0.541 0.459 18.8 0.019 1.51 ± 0.10 52.4 0.462 0.048 0.541 0.459 23.2 0.023 2.09 ± 0.15 59.9 0.492 0.060 0.544 0.456 26.2 0.026 2.28 ± 0.16 65.4 0.498 0.082 0.543 0.457 28.6 0.029 2.86 ± 0.20 72.0 0.521 0.081 0.547 0.453 31.1 0.031 2.60 ± 0.18
47
TABLE 4.3
VAPOR-LIQUID EQUILIBRIUM FOR CO2-O2-[hmim][Tf2N] AT (40 ± 0.5)˚C
WITH A FEED COMPOSITION OF 50 MOLE PERCENT O2 WITH THE BALANCE
CO2.
liquid phase vapor phase composition composition
xO2 pure
Ptotal fO2 gas mole
(bar) xCO2 xO2 yCO2 yO2 (bar) fraction EF 20.4 0.068 0.017 0.477 0.523 10.5 0.011 1.66 ± 0.12 31.2 0.154 0.041 0.453 0.547 16.7 0.017 2.47 ± 0.18 33.8 0.180 0.054 0.444 0.556 18.3 0.018 2.97 ± 0.21 40.3 0.227 0.061 0.433 0.567 22.2 0.022 2.76 ± 0.20 45.0 0.235 0.064 0.441 0.559 24.4 0.024 2.61 ± 0.18 52.9 0.350 0.094 0.389 0.611 31.2 0.031 3.01 ± 0.21 58.4 0.399 0.097 0.378 0.622 34.9 0.035 2.77 ± 0.20 68.5 0.434 0.131 0.381 0.619 40.5 0.040 3.23 ± 0.23
48
TABLE 4.4
VAPOR-LIQUID EQUILIBRIUM FOR CO2-O2-SOLVENT AT (40 ± 0.5)˚C
liquid phase vapor phase composition composition
Ptotal fO2
(bar) xCO2 xO2 yCO2 yO2 (bar) EF solvent 33.5 0.368 0.015 0.665 0.335 11.0 1.39 ± 0.10 IL 29.6 0.16 0.005 0.6 0.38 11.1 0.99 Acetonitrile 30.1 0.08 0.005 0.55 0.43 13 1.1 Methanol 68.5 0.434 0.131 0.381 0.619 40.5 3.23 ± 0.23 IL 69.4 0.2 0.022 0.37 0.61 41.6 1.08 Acetonitrile
49
TABLE 4.5
VAPOR-LIQUID EQUILIBRIUM FOR CO2-O2-[hmim][Tf2N] AT (40 ±
0.5)˚C WITH VARYING FEED COMPOSITIONS
liquid phase vapor phase composition composition xO2 pure
Ptotal fO2 gas mole
(bar) xCO2 xO2 yCO2 yO2 (bar) fraction EF 20.0 0.275 0.009 0.654 0.346 6.8 0.007 1.28 ± 0.09 19.5 0.223 0.012 0.532 0.468 9.0 0.009 1.29 ± 0.09 20.4 0.068 0.017 0.477 0.523 10.5 0.011 1.66 ± 0.12 33.5 0.368 0.015 0.665 0.335 11.0 0.011 1.39 ± 0.10 42.5 0.404 0.022 0.676 0.324 13.4 0.013 1.66 ± 0.12 41.7 0.429 0.021 0.661 0.339 13.7 0.014 1.56 ± 0.11 31.2 0.328 0.022 0.535 0.465 14.2 0.014 1.57 ± 0.11 48.2 0.462 0.031 0.669 0.331 15.4 0.015 2.02 ± 0.14 31.2 0.154 0.041 0.453 0.547 16.7 0.017 2.47 ± 0.18 57.2 0.455 0.054 0.694 0.306 16.8 0.017 3.21 ± 0.23 58.7 0.462 0.049 0.691 0.309 17.4 0.017 2.79 ± 0.20 33.8 0.180 0.054 0.444 0.556 18.3 0.018 2.97 ± 0.21 42.1 0.385 0.028 0.541 0.459 18.8 0.019 1.51 ± 0.10 40.3 0.227 0.061 0.433 0.567 22.2 0.022 2.76 ± 0.20 69.8 0.611 0.087 0.662 0.338 22.5 0.023 3.87 ± 0.27 52.4 0.462 0.048 0.541 0.459 23.2 0.023 2.09 ± 0.15 45.0 0.235 0.064 0.441 0.559 24.4 0.024 2.61 ± 0.18 59.9 0.492 0.060 0.544 0.456 26.2 0.026 2.28 ± 0.16 65.4 0.498 0.082 0.543 0.457 28.6 0.029 2.86 ± 0.20 72.0 0.521 0.081 0.547 0.453 31.1 0.031 2.60 ± 0.18 52.9 0.350 0.094 0.389 0.611 31.2 0.031 3.01 ± 0.21 58.4 0.399 0.097 0.378 0.622 34.9 0.035 2.77 ± 0.20 68.5 0.434 0.131 0.381 0.619 40.5 0.040 3.23 ± 0.23
50
4.3 Equal Fugacity Plots
The following plots, Figures 4.1 and 4.2, show graphically the same
data as the enhancement factors. It is important to note that for CO2, at similar fugacities, the solubility does not change significantly, in contrast
to preliminary data presented by Hert et al. [52]. This means that CO2 is
not negatively affected by the presence of O2. The slight increase in the
solubility of CO2 at higher pressures can be attributed to the higher
overall pressure of the system. With O2, the solubility increases dramatically. Although this can partially be attributed to the higher total pressure, by comparing the different feed compositions, it is apparent the
presence of CO2 does enhance the O2 solubility, consistent with the results presented by Hert et al. [52]. Previous experimental data
presented in this way showed an obvious decrease in CO2 over a smaller and lower pressure range (3.0 – 15.7 bar) [52] but we now believe these results to be inaccurate. Considerable improvements of the experimental apparatus have been made to provide better precision and accuracy making the current results significantly more accurate and more reliable.
Tables of the data presented here are found in the appendix.
51
0.7
0.6
0.5
0.4 Mole Fraction 2 0.3 CO Pure 75% 0.2 65% 50%
0.1 0 10 20 30 40 CO Fugacity (bar) 2 Figure 4.1. Solubility of carbon dioxide in [hmim][Tf2N] as a function of fugacity at (40 ± 0.5)oC.
52
0.14
0.12
0.10
0.08
0.06 Mole Fraction 2 O 0.04 Pure 25% 0.02 35% 50%
0.00 0 10 20 30 40 50
O2 Fugacity (bar)
Figure 4.2. Solubility of oxygen in [hmim][Tf2N] as a function of fugacity at (40 ± 0.5)oC.
53 4.4 Ternary Diagrams
Ternary diagrams show a more accurate description of what is going on inside the system by showing both the liquid and vapor phase.
Figures 4.3-8 show ternary diagrams at pressures around 20, 30, 40, 50,
60, and 70 bar with pressure range detailed in each caption. Unlike enhancement factors which show a comparison at the individual component fugacity, ternary plots are at the total system pressure which gives a more accurate picture of the operating conditions.
On each diagram, the right side of the triangle is the vapor phase
and is composed of only the two gases, CO2 and O2. [hmim][Tf2N] has a negligible vapor pressure and is therefore not present in the vapor phase.
It is evident in Figure 4.9-11 by comparing the vapor phases over the pressure range that each feed composition falls to about the same
equilibrium gas phase composition. The 25 mole percent O2 feed falls to
~30-35 mole percent O2, the 35 mole percent O2 feed falls to ~45 mole
percent O2, and the 50 mole percent O2 feed falls to ~55-60 mole
percent O2 with the balance being CO2. This presents a potential way to control the equilibrium vapor phase by knowing how the vapor phase compositions change upon equilibrium.
The significant part of the ternary diagram is the lower left portion
where the O2 solubility in the IL can be seen. In a non-enhanced system, a straight line can be drawn between the two pure gas solubilities and each
54 data point should fall on this straight line. However, in an enhanced system, the tie lines would arch above the straight line between the two pure data points. It is apparent in Figure 4.4 that enhancement can be readily seen above pressure of 30 bar and even more noticeably at higher pressures.
Figures 4.12 and 4.13 show the organics (acetonitrile, methanol,
and acetone) as well as [hmim][Tf2N] at 30 and 50 bar, respectively, with pressure range detailed in caption. A few items are important to note in these figures. First, the organics have a vapor pressure so their vapor phase does not fall on the right ternary line as the IL vapor phase does.
However, at the pressures and temperatures study, the vapor phase is
mostly CO2 and O2 so it is comparable. Second, the tie-lines are fairly
similar sloped, however, there is a higher CO2 solubility in the IL than the
organics. Lastly, the O2 solubility is similar at 30 bar, however, at 50 bar
there is a noticeable increase in the O2 solubility in the IL.
The pure gas solubility data shown in Figures 4.3-4.8 is the same as that presented in section 1.5.2. In all graphs the circles represent pure
data, squares represent a feed gas of 25 mole percent O2, diamonds
represent a feed gas of 35 mole percent O2, and triangles represent a
feed gas of 50 mole percent O2 with the balance being CO2.
55
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.3. CO2 (1) – O2 (2) – [hmim][Tf2N] at (20 ± 2) bar and (40 ± 0.5)˚C.
56
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.4. CO2 (1) – O2 (2) – [hmim][Tf2N] at (30 ± 4) bar and (40 ± 0.5)˚C.
57
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.5. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 3) bar and (40 ± 0.5)˚C.
58
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.6. CO2 (1) – O2 (2) – [hmim][Tf2N] at (50 ± 3) bar and (40 ± 0.5)˚C.
59
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.7. CO2 (1) – O2 (2) – [hmim][Tf2N] at (60 ± 5) bar and (40 ± 0.5)˚C.
60
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
Pure 25% O2, 75% CO2 35% O2, 65% CO2 50% O2, 50% CO2
Figure 4.8. CO2 (1) – O2 (2) – [hmim][Tf2N] at (70 ± 2) bar and (40 ± 0.5)˚C.
61
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
20 bar 30 bar 40 bar 50 bar 60 bar 70 bar
Figure 4.9. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C and pressures from 20 - 70 bar for a 25 mole % O2 feed with the balance being CO2.
62 O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
20 bar 30 bar 40 bar 50 bar 60 bar 70 bar
Figure 4.10. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C and pressures from 20 - 70 bar for a 35 mole % O2 feed with the balance being CO2.
63 O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 IL (3) CO2 (1)
20 bar 30 bar 40 bar 50 bar 60 bar 70 bar
Figure 4.11. CO2 (1) – O2 (2) – [hmim][Tf2N] at (40 ± 0.5)˚C
and pressures from 20 - 70 bar for a 50 mole % O2 feed with
the balance being CO2.
64
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 solvent (3) CO2 (1)
IL acetonitrile methanol
Figure 4.12. Ternary phase behavior for CO2 (1)- O2 (2)- solvent (3) at (30 ± 4) bar and (40 ± 0.5)˚C where the organic solvents are acetonitrile and methanol and the IL is
[hmim][Tf2N].
65
O2 (2)
0.0 1.0
0.1 0.9
0.2 0.8
0.3 0.7
0.4 0.6
0.5 0.5
0.6 0.4
0.7 0.3
0.8 0.2
0.9 0.1
1.0 0.0 0.0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 solvent (3) CO2 (1)
IL acetone acetonitrile methanol
Figure 4.13. Ternary phase behavior for CO2 (1)- O2 (2)- solvent (3) at (50 ± 3) bar and (40 ± 0.5)˚C where the organic solvents are acetonitrile, acetone and methanol and
the IL is [hmim][Tf2N].
66 CHAPTER 5:
CONCLUSIONS
5.1 Conclusions
This work has shown that the presence of CO2 enhances the
solubility of O2. This enhancement is likely due to the presence of CO2 in the IL liquid solution increasing the free volume of the solution as well as
creating greater non-polar interactions with the O2. The CO2 may also be effective in shielding the charges on the cation and the anion. This enhancement was quantified three ways.
This work has presented enhancement factors ranging from (1.28 ±
0.09) – (3.87 ± 0.27) in pressures ranging from (18.7 – 72.5) bar at (40
± 0.5)˚C. Previous research on the same system at pressures ranging from (3.0 - 15.1) bar at (25 ± 0.5)˚C found enhancement factors ranging from 1.4 - 5.7 [52]; however, there was greater experimental uncertainty in these previous measurements that has since been fixed. In organic systems at pressures ranging from (13.7 - 52.4) bar at (40 ±
0.5)˚C, enhancement factors of 0.59 - 1.31 were found [51]. Thus, it is
apparent the presence of CO2 has a greater effect on the IL, consistent
with the preliminary results [52]. The presence of CO2 may also decrease
67 the solution viscosity and increase the rate at which the O2 dissolves in the solution, allowing equilibrium to be reached more quickly.
Plots of the component fugacity and mole fraction showed graphically the same data as the enhancement factor. Interestingly, there
is a slight increase in the CO2 mole fraction solubility. This is expected to be due to the increased overall pressure of the system and not the
presence of the lower solubility gas, O2.
Lastly, ternary diagrams showed the vapor-liquid equilibrium at the total system pressure. As expected, solubility increased with increasing system pressure and noticeable enhancements were seen above 30 bar.
Enhancement is seen in a ternary diagram by an arched curve between the two pure gas phases rather than a direct line.
These results are promising for IL use as reaction media. The
presence of CO2, a non-flammable gas, also makes the system safer.
However, more economic motivators are needed before this research sees widespread industrial application. Rate and selectivity in reactions as well as the potential of IL solvent recycling need to be investigated.
5.2 Future Work
There is great potential for future work in this project. Rarely is a pure gas found in industry, so understanding the phase behavior of mixed gas systems is crucial. For further research into IL use as a reaction
solvent, other reaction gases such as CO, H2, and CH4 can be studied at
68 similar pressure ranges of 20 – 70 bar. Additionally, data at more temperatures would allow for an enthalpy and entropy of absorption estimation.
There is also great potential to modify the system. The system volume is ideal with about 15 mL vapor space and 5 mL liquid, however, the geometry of the system causes limitations. Instead of a cylindrical cell, a rounder cell that allowed for a greater vapor-liquid interface and a different stir bar would decrease equilibrium time drastically. Additionally, all components of the system are piped in place. It would be beneficial to have the option to change the stir plate or allow for easier changing of tubing. The single most beneficial change to the system was the addition of the vacuum pump (pulls to 4.0x10-3 mbar) to remove residual gas in between measurements. However, bubbles can still be seen in the crevices of the cell so minimizing the places gas can absorb at the crevice of the cell seal and the stir bar would be ideal.
69
APPENDIX A:
INTERPOLATING OXYGEN SOLUBILITY DATA
There are two methods to interpolate the oxygen solubility data, the linear method and the Henry’s Law method. Each yields the same result and will be discussed.
A.1 Linear Method.
Pure solubility of oxygen data in [hmim][Tf2N] at 40ºC was not available. However, data at 10ºC, 25ºC and 50ºC has been taken in our laboratory using the IGA microbalance and is shown in Figure A.1. [50].
To interpolate comparable pure data at 40ºC, a linear trend line was fit to each of the three temperatures where y is the mole fraction and x is the pressure (bar).
For 10ºC
y = (0.0010 ± 0.0002)x A.1
!
70
0.020
0.015
0.010 Mole Fraction 2 O
0.005 10 oC 25 oC 50 oC
0.000 0 2 4 6 8 10 12 14
Pressure (bar)
Figure A.1. Solubility of pure oxygen in [hmim][Tf2N]. The experimental points at 10ºC, 25ºC, and 50ºC are shown.
71
For 25ºC
y = (0.0010 ± 0.0002)x A.2
For 50ºC ! y = (0.0014 ± 0.0004)x A.3
An equation was then linearly interpolated to predict pure oxygen solubility at 40ºC.! This equation was weighted 60 % on the 50ºC data and
40 % on the 25ºC data, the 10ºC data was not used for the interpolation.
This resulted in the following equation where y is the mole fraction and x is the Pressure (bar) at 40ºC.
y = (0.0012 ± 0.0003)x A.4
One fundamental uncertainty associated with this method is the assumption that !oxygen solubility is a linear function of temperature.
72 A. 2. Henry’s Law Method.
The second method of interpolating the data is more thermodynamically based and uses the same pure data shown in Figure
A.1. From this pure solubility data, a Henry’s Law Constant can be calculated from the graph of mole fraction vs. Pressure (bar) where the limiting linear slope is 1 and H is in bar. H
!
TABLE A.1
HENRY’S LAW CONSTANTS AS A FUNCTION OF TEMPERATURE FOR
[hmim][Tf2N]
Temperature (oC) Henry's Law (bar) 10 988.3 ± 101.2 25 874.5 ± 65.7 50 1107.5 ± 244.8
73
At this point it is important to note again that the Henry’s Law
Constants do not follow the expected trend of lower temperature causing
higher solubility. Once Henry’s Law Constants are known, they can be
graphed to find the enthalpy defined by the equation
# & % " lnH ( % ( = )h A.5 % " 1 ( $ T 'P
The enthalpy given in the equation of a line for where y is ln H and x is ! 1 is T ! y = ("308 ± 348)x + (7.9 ± 1.3) A.6 ! For 40˚C, 1 = 0.00319 giving ln H = 6.9 ± 0.3. The Henry’s Law T ! Constant is then 1028 ± 311 bar. From this point, we reverse the ! process.! We can generate an equation where y is the mole fraction, x is
the pressure (bar), the y-intercept is zero, and the slope is 1 . We can H
then insert a spread of pressure values and check the mole fraction
values generated. This method gives the equation !
y = (0.0010 ± 0.0003)x A.7
!
74
7.3
7.2
7.1
7.0 lnH
6.9
6.8
6.7
6.6 0.0030 0.0031 0.0032 0.0033 0.0034 0.0035 0.0036
1/T (1/K)
Figure A.2. The partial molar enthalpy of solution is the slope of ln H vs. 1 . The uncertainty in each data point is shown. T
! !
75 A.3. Comparison of Two Methods
Both methods generate a linear equation for solubility. Where y is the mole fraction, x is the pressure (bar), the y-intercept is zero, the equations generated are for the linear method:
y = (0.0010 ± 0.0002)x A.1 and for the Henry’s Law method: ! y = (0.0010 ± 0.0003)x A.7
The data generated from both methods is obviously similar, however, the Henry’s! Law fit is chosen due to its fundamental thermodynamic basis.
76
APPENDIX B:
PROCEDURE FOR MIXED GAS SOLUBILITY MEASUREMENTS
B.1 For Initial Point
B.1.1 GC Preparation
1. Turn Regulator on Helium Tank up to 60 psi. 2. Confirm the yellow valves are in the vapor phase sampling position and all valves connecting the 2 GCs are in the proper location. 3. Activate the “Sampling.mth” method. a. Click the Open icon next to the Method display and selecting the “Sampling.mth” method from the “Aki” folder inside the “Star” folder. 4. Turn the electronics on. a. Left-click on the Method display and select “View/Edit Method.” b. Click the “Electronics” tab and select “Turn Electronics On.” c. Save and Exit the Method editor. d. Left-click on the Method display and select “Re-activate Method.” 5. Create a sampling list. a. If no sample list is loaded, select “File” then “Open Sample List” and select aki.SMP from the Star Folder. b. Set up the list to have 2 blanks and your sample. Name the sample. c. Click “Data Files” on the bottom right corner of the sample list to indicate where to save the files (usually by day). d. Click “Begin” and hit “ok” however many times necessary to start the sample list e. If you cannot select “Begin,” you must first select “Suspend,” then go through steps c. and d.
77 B.1.2 Blank
1. Turn vacuum pump on. a. “Vent” and “Load” valves should both be open to the vacuum. 2. Hot button on GC a. Confirm the system is ready – the yellow “Waiting” text will flash on the control window. b. Hit the hot button on the GC c. Turn sampling valve to the “Inject” postion (right) then back to the “Load” position (left). d. Let the first blank run the full time until the temperature reaches 175˚C. e. Allow the GC to cool back down. 3. Repeat a. Confirm system is ready – the yellow “Waiting” text will flash on the control window b. Hit the hot button on the GC c. Turn sampling valve to the “Inject” position (right) then back to the “Load” position (left). d. Let the sample run for as long as it takes your samples to
come off. For CO2 and O2 this is only 5-6 minutes. e. Allow the GC to cool back down.
B.1.3 Initial Readings – while the blanks are running.
1. Measure & Record the ambient pressure in A53. 2. Measure the liquid level with the cathetometer a. Make sure the cathetometer is level. b. Measure & Record the bottom (reference) c. Turn off the stir bar. d. Measure & Record the liquid level e. Turn on the stir bar.
B.1.4 Charging the Cell & Initial Data Point
1. Confirm system is ready – the yellow “Waiting” text will flash on the control window 2. Open the valve on the top right of the air bath – valve 8 – to use gas straight from the tank, or the valve on the top left – unnumbered – to use gas from the ISCO pump. 3. Open the gas mixture tank or the ISCO pump.
78 4. Turn the regulator on the tank to 50-100 psi above your desired cell pressure. 5. Turn off the stir bar. 6. Measure & Record the Cell & Loop temperatures a. The cell is now 1. The loop is 4. b. Acceptable temperature range is 39.5-40.5˚C. 7. Confirm the Vent and Load valves are closed.****** 8. Open valve 12 (Kate Diagram Valve 1) slowly to let the mixture into the cell. Close the valve when you have reached your predetermined initial charge pressure. 9. Once the pressure stabilizes, record it. 10. Then open the Load valve (Valve 4). 11. Record the new pressure. 12. Hit the hot button on the GC. 13. Turn the sampling valve to the “Inject” position. 14. Close the Load valve (Valve 4). 15. Turn the sampling valve to the “Load” position. 16. Open the Vent valve (Valve 5).
B.2 For Equilibrium Point
B.2.1 GC Preparation
Repeat as for initial point.
B.2.2 Blank
Repeat as for initial point.
B.2.3 Initial Readings
Repeat as for initial point.
B.2.4 Equilibrium Data Point
1. Confirm system is ready – the yellow “Waiting” text will flash on the control window 2. Keep the stir bar on. 3. Measure & Record the Cell & Loop temperatures a. The cell is now 1. The loop is 4. b. Acceptable temperature range is 39.5-40.5degC. 4. Confirm the Vent and Load valves are closed (Valves 4 and 5). 5. Open the Load valve (Valve 4). 6. Record the new pressure.
79 7. Hit the hot button on the GC. 8. Turn the sampling valve to the “Inject” position. 9. Close the Load valve (Valve 4). 10. Turn the sampling valve to the “Load” position. 11. Open the Vent valve (Valve 5). 12. Repeat 2 more times. 13. Vent the cell and wait for gas to be removed.
B.3. Shutting the GC down.
After the GC has been used, confirm the electronics are turned off and the carrier gas has been turned down.
80
APPENDIX C:
GAS CHROMATOGRAPH METHOD SETTINGS
In addition to the following settings, the carrier gas, Helium, is set at 60 units.
Visible on the screen while the experiment is running:
Col Oven: 30 Front 1041: 200 Coolable 2: n/a Rear Valve: 50 Front TCD: 250 Heatable 2: n/a Heatable 3: n/a
Front Detector Status:
Ready: Yes Fault: No Front TCD Electronics: On Range: 0.5 Time Const: Fast Filament Temp: 225C Current: ~53 mA Balance Pct= -99% Polarity: Positive Detector Signal: ~0.034mV Bunch size: 4 Freq: 10.00 Hz
Inside the sampling.mth tab:
3800 GC Control
Sample Delivery: Front & Middle Valve Oven Installed •No Rear Valve Oven Installed •Yes Valve Oven •On Temp(C) 50 81 Valve 1: Gas Sampling
Injector: Middle & Rear Injector •None Front Injector Type 1041 Injector Oven •On Temperature (C) 200
Flow/Pressure: Middle & Rear EFC •None Front EFC Type 4 (for Valved Systems) Pressure (psi) 22.5 Hold (min), Time (min) 19.30 Time Initial Total Flow mL/min 30
Column Oven: Column Oven Coolant •Off Temp Rate(C/min) Hold(min) Time 30 6 6 175 15.0 4 19.67
Detector: Middle & Rear Detector •None Front Detector: Front Detector Type: TCD Detector Oven •On Temperature(C) 250 Electronics •On
Time: Initial Range: 0.5 Autozero: Yes Polarity: positive
Adjustments: Time Constant •Fast Carrier Gas: He Make-Up Flow: 0 Filament Temp Limit: 390 Ref/Make-Up Flow: 29
Output: Port A,B,C Installed •Yes Time: Initial Signal Source: Front Attenuation: 1
Data Acquisition: Detector Bunch Rate: 4
82 Noise Monitor Length: 64 FID/TSD: Front: IV; Middle: IV; Rear: IV
83
APPENDIX D:
ERROR ANALYSIS
D.1. Calibration Curve Error Analysis
The calibration curve error analysis is the same method for the CO2
GC calibration, the O2 GC calibration, fugacity fit and the liquid height to
volume calibration. The method of least squares is chosen because it is
objective, without systematic bias and the data is linear.
The calibration will be forced with have a y-intercept at zero so it
will fit the line:
y = mx D.1
The vertical deviation, d i , from this line is: ! d i = yi " y D.2 ! where y i is the measured value and y is the value from the line. ! The vertical deviation is squared to minimize the magnitude:
! ! 2 d 2 = y " y D.3 i ( i )
!
84 which gives the expression for the slope:
n x y " x y m # i i # i # i D.4 = 2 2 n x " x # ( i ) (# i )
where n is the number of points taken and x i , yi are experimental points. ! The standard deviation is then
! ! 1/2 % d 2 ( ' # i * " y = D.5 ' n $ 2 * & ( ))
and ! " 2 # n y D.6 "m = 2 2 n x $ x % ( i ) (% i )
This yields an equation for the line: ! y = m ± " x D.7 ( m )
D.2. Data Error Analysis!
Here, I will explain how error is propagated throughout analyzing an
experimental point. The raw experimental data includes:
• Reference liquid height (bottom of the cell) • Actual liquid height • Temperature of the cell (˚C) • Temperature of the sample loop (˚C) • Pre-load Pressure (psi) • Post-load Pressure (psi)
• Area counts O2
• Area counts CO2
85 The first thing to do is convert the area counts to moles of gas using the calibration curves developed. The calibrations generated are as follows where y is the moles of the gas and x is the GC area counts:
For CO2: ! ! y = 3.6171x10"10 ± 2.97x10"12 x D.8 ( )
For O2:
! "10 y = 4.5323x10 x D.9
-24 The O2 uncertainty is on the order of 10 and is therefore negligible.
Next, we !find the mole fraction of each gas in the sample loop. For
O2 mole fraction in sample loop (or vapor space):
N O2 xO = D.10 2 N + N O2 CO2
1/2 # 2& 2 # 2 2 & %# & "N + "N ( ! N "N % O2 CO2 ( O2 O2 ( ) ( ) "x = %% ( + % ( ( D. 11 O2 %% ( ( NO + NCO N N + N 2 2 $ O2 ' % O2 CO2 ( % $ ' ( $ ' and for CO2 mole fraction in the sample loop (or vapor space): ! N CO2 xCO = D.12 2 N + N O2 CO2
1/2 # 2& 2 # 2 2 & %# & "N + "N ( ! N "N % O2 CO2 ( CO2 CO2 ( ) ( ) "x = %% ( + % ( ( D.13 CO2 %% ( ( NO + NCO N N + N 2 2 $ CO2 ' % O2 CO2 ( % $ ' ( $ '
! 86 We find the height of the liquid by simply subtracting the Reference
Liquid Height from the Actual Liquid Height giving us a value in cm. This is converted to volume of the IL using a calibration curve where y is the volume (mL) and x is the height (cm).
y = (1.573± 0.001)x D.14
To find the volume of the vapor space, the IL volume is subtracted from the total volume ! of the cell and it’s lines, 19.8 mL. The error in the IL volume and vapor space volume is the same.
Next, we use a ratio to find the number of moles in the vapor space
(head space) for CO2 and O2. The volume of the loop is 0.00908 mL.
For O2:
" % Vhead NO ,head = $N O ,loop x ' D.15 2 $ 2 V ' # loop &
2 2 # & # & # & V "NO ,loop "V !"N = %N x head ( % 2 ( + % head ( D.16 O2 ,head % O2 ,loop ( % ( % ( $ Vloop ' NO ,loop Vhead $ 2 ' $ '
And for CO2: ! " V % N = $N x head ' D.17 CO2 ,head $ CO2 ,loop V ' # loop &
2 2 # & # & # & V "NCO ,loop "V !"N = %N x head ( % 2 ( + % head ( D.18 CO2 ,head % CO2 ,loop ( % ( % ( $ Vloop ' NCO ,loop Vhead $ 2 ' $ '
!
87 The calculations to this point are preformed for the data collected
at the initial and equilibrium stages. The following calculations are
performed only on the equilibrium data.
The number of moles absorbed in the IL is simply found by
difference.
For O2:
initial equilibrium NO ,absorbed = NO ,headspace " N D.19 2 2 O2 ,headspace
2 2 "N = N initial # Nequilibrium "N initial + "Nequilibrium O2 ,absorbed O2 ,headspace O ,headspace O2 ,headspace O ,headspace ! ( 2 ) ( ) ( 2 ) D.20 ! And for CO2:
initial equilibrium NCO ,absorbed = NCO ,headspace " N D.21 2 2 CO2 ,headspace
2 2 "N = N initial # Nequilibrium "N initial + "Nequilibrium CO2 ,absorbed CO2 ,headspace CO ,headspace CO2 ,headspace CO ,headspace ! ( 2 ) ( ) ( 2 ) D.22
! The mole fraction of each component in the liquid mixture can then be
found.
For O2:
N O2 ,absorbed xO = D.23 2 N + N + N CO2 ,absorbed O2 ,absorbed IL
1/2 # 2& 2 # 2 2 & # &% # & "N + "N ( ! N "N % O2 CO2 ( O2 % O2 ( ) ( ) ( "xO = % ( % ( + % ( D.24 2 % N + N + N (% % N ( N N ( $ CO2 O2 IL ' $ O2 ' % O2 + CO2 ( % $ ' ( $ '
88 ! For CO2:
N CO2 ,absorbed xCO = D.25 2 N + N + N CO2 ,absorbed O2 ,absorbed IL
1/2 # 2& 2 # 2 2 & # &% # & "N + "N ( ! N "N % O2 CO2 ( CO2 % CO2 ( ) ( ) ( "xCO = % ( % ( + % ( D.26 2 % N + N + N (% % N ( N N ( $ CO2 O2 IL ' $ CO2 ' % O2 + CO2 ( % $ ' ( $ '
And for [hmim][Tf2N]: ! NIL xIL = D.27 N + N + N CO2 ,absorbed O2 ,absorbed IL
1/2 # 2& # 2 2 & # &% "N + "N ( ! % O2 CO2 ( NIL % ( ) ( ) ( "xIL = % ( % ( D.28 % N + N + N (% N N ( $ CO2 O2 IL ' % O2 + CO2 ( %$ ' ( $ '
At this point, we have fully analyzed the data for both the liquid
and! vapor phases in order to construct ternary diagrams. The final error
propagation is in the calculation of the enhancement factor. The
enhancement factor is defined in chapter 4 as:
xmixture EF = gas D.29 x pure gas
2 2 xmixture # "xmixture & # "x pure & "EF = gas % gas ( + % gas ( D.30 ! pure mixture pure x % x ( % x ( gas $ gas ' $ gas '
!
89
APPENDIX E:
EXPERIMENTAL DATA SHOWN IN FUGACITY PLOTS
Tables E.1-3 show the experimental data graphed in the fugacity plots shown in Chapter 4 at feed compositions of 25, 35, and 50 mole percent
O2 with the balance being CO2. The pure gas mole fraction is at the same component fugacity.
90
TABLE E.1
VAPOR-LIQUID EQUILIBRIUM DATA FOR CO2-O2-[hmim][Tf2N] AT
(40 ± 0.5)˚C WITH A FEED OF 25 MOLE PERCENT O2 AND THE BALANCE
CO2.
liquid phase composition xCO2 pure xO2 pure
Ptotal gas mole gas mole
(bar) xCO2 xO2 fCO2 (bar) fO2 (bar) fraction fraction 18.7 0.282 0.011 10.9 6.6 0.158 0.007 20.0 0.275 0.009 11.9 6.8 0.172 0.007 33.5 0.368 0.015 18.9 11.0 0.274 0.011 41.7 0.429 0.021 22.5 13.7 0.326 0.014 42.5 0.404 0.022 23.4 13.4 0.339 0.013 48.2 0.462 0.031 25.4 15.4 0.368 0.015 57.2 0.455 0.054 29.7 16.8 0.430 0.017 58.7 0.462 0.049 30.2 17.4 0.437 0.017 69.8 0.611 0.087 32.1 22.5 0.466 0.023
------
91
TABLE E.2
VAPOR-LIQUID EQUILIBRIUM DATA FOR CO2-O2-[hmim][Tf2N] AT (40 ±
0.5)˚C WITH A FEED OF 35 MOLE PERCENT O2 AND THE BALANCE CO2.
liquid phase composition xCO2 pure xO2 pure
Ptotal gas mole gas mole
(bar) xCO2 xO2 fCO2 (bar) fO2 (bar) fraction fraction 19.5 0.223 0.012 9.4 9.0 0.137 0.009 31.2 0.328 0.022 14.3 14.2 0.208 0.014 36.6 0.378 0.043 16.3 16.7 0.236 0.017 42.1 0.385 0.028 18.5 18.8 0.269 0.019 52.4 0.462 0.048 21.8 23.2 0.316 0.023 59.9 0.492 0.060 24.0 26.2 0.348 0.026 65.4 0.498 0.082 25.5 28.6 0.369 0.029 72.0 0.521 0.081 27.1 31.1 0.393 0.031
------
92
TABLE E.3
VAPOR-LIQUID EQUILIBRIUM DATA FOR CO2-O2-[hmim][Tf2N] AT (40 ±
0.5)˚C WITH A FEED OF 50 MOLE PERCENT O2 AND THE BALANCE CO2.
liquid phase composition xCO2 pure xO2 pure Ptotal gas mole gas mole
(bar) xCO2 xO2 fCO2 (bar) fO2 (bar) fraction fraction 20.4 0.068 0.017 8.8 10.5 0.128 0.011 31.2 0.154 0.041 12.1 16.7 0.176 0.017 33.8 0.180 0.054 12.7 18.3 0.185 0.018 40.3 0.227 0.061 14.3 22.2 0.208 0.022 45.0 0.235 0.064 15.9 24.4 0.231 0.024 52.9 0.350 0.094 15.8 31.2 0.229 0.031 58.4 0.399 0.097 16.4 34.9 0.238 0.035 68.5 0.434 0.131 18.2 40.5 0.264 0.040
------
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