Supercritical Fluids and Gas-Expanded Liquids

Chapter 7

Supercritical Fluids and Gas-Expanded Liquids Larissa P. Cunico and Charlotta Turner Lund University, Lund, Sweden

7.1 INTRODUCTION A supercritical fluid (SCF) is any substance for which both temperature and pressure are above the critical point (CP, Fig. 7.1). In SCF state, liquid, and gas phases do not exist. If a liquid substance is in equilibrium with its vapor in a closed vessel and heated to a temperature above its critical temperature (Tc), leading to a pressure over its critical pressure (Pc), the interface between the two phases (liquid, L, and vapor, V) will diminish, and only one phase remains filling the entire space of the vessel, i.e., an SCF phase. At equilibrium conditions this means moving along the vaporization line toward the CP and beyond (Fig. 7.1). A visualization of such phase transfer is shown in Fig. 7.2 [1]. What characterizes an SCF is “gas-like” viscosity and “liquid-like” density (Table 7.1). The low viscosity inherently leads to fast diffusivity, which is important to achieve fast mass transfer in separation processes. SCFs are sometimes called “compressed gases,” since they are commonly obtained from substances that are gases at ambient conditions, but also reflecting the physical properties of an SCF (gas-like viscosity and compressed, liquid-like density). In addition, the compressibility of an SCF is much larger than that of a liquid, and therefore, a relatively small change in pressure leads to a quite large change in volume and density. This makes SCFs interesting in separation processes since solubility can easily be tuned by changing temperature and pressure. There are several options of substances to be used as SCFs in separation processes [2], such as propane, ethylene, ammonia, nitrous oxide and carbon dioxide, of which the latter is by far the most commonly used (Table 7.2). Supercritical carbon dioxide (scCO2) has several benefits as compared to the other SCFs, in terms of inertness, high purity at low cost, nonflammability, relatively low toxicity, and easily attainability CP (Tc 5 31
C, Pc 5 7.4 MPa) The Application of Green Solvents in Separation Processes. DOI: http://dx.doi.org/10.1016/B978-0-12-805297-6.00007-3 © 2017 Elsevier Inc. All rights reserved.

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SECTION | II Green Solvents

Pc

Supercritical fluid (SCF)

Liquid

Solid Melting Freezing

Pressure

156

ion zat n ori satio p Va den n Vapor Co

n io at on im ti bl osi u S ep D

Triple point

Tc

Critical point (CP)

Gas

Temperature

FIGURE 7.1 Schematic phase diagram (pressure vs temperature) showing the three classical states of matter: solid, liquid, and gas, as well as the supercritical fluid (SCF) state. Sublimation, melting, and vaporization lines are indicated, as well as the triple point and the critical point (CP). Pc is the critical pressure and Tc is the critical temperature.

(Table 7.2). In this chapter, mainly scCO2 will be considered as a solvent since this is the most commonly used SCF in separation processes. In addition, CO2 is produced as a byproduct from many processes such as biogas production, fermentation of sugars into ethanol, hydrogen production from methane, and in sodium phosphate manufacturing, making scCO2 a potential green solvent. In terms of solubility, scCO2 is a solvent of extremely low relative static permittivity (dielectric constant), similar to hexane or even smaller (depending on pressure and temperature [3,4]. This means that scCO2 can be used to dissolve mainly nonpolar low-molecular-weight compounds. For instance, compressed liquid CO2 or scCO2 is used in cleaning processes to remove fats and oils, in decaffeination of coffee beans, extraction of aromatic compounds to produce perfumes, and in dry cleaning of textiles. In analytical chemistry, scCO2 is used as a solvent in extraction (supercritical fluid extraction, SFE) [510] and in chromatography (supercritical fluid chromatography, SFC) [1114]. See also Chapter 11, Environmentally Benign Supercritical Fluid Extraction, (SFE) and Chapter 16, Supercritical Fluid Chromatography (SFC) for further information about these techniques. Since scCO2 is a solvent of low polarity, with a dipole moment of zero Debye, it has its limitations in separation processes. The main solution to enhance polarity is to add a cosolvent, also called entrainer or modifier, to the scCO2. There are a large number of publications describing SFE and

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157

FIGURE 7.2 Pictures of a variable volume view-cell with a transparent sapphire window showing the phase transfer of CO2 from (A) vapor-liquid phase equilibrium where a clear meniscus is shown; (B) with an increasing temperature the meniscus begins to disappear; (C) the densities of the liquid and vapor phases are more and more similar; and finally (D) the CP is reached and the two phases are no longer distinguishable. Here, the CO2 is in the SCF state. Taken from University of Leeds, School of Chemistry, http://www1.chem.leeds.ac.uk/ (accessed February, 2016)

SFC using scCO2 with mainly either ethanol or methanol added in small volume percentage [1519]. Recently, it has been shown that instead of adding a cosolvent to scCO2, compressed liquid CO2 can be added to an organic solvent, thereby creating a so-called gas-expanded liquid (GXL). In terms of physical properties, a GXL has similar density compared to the organic solvent without CO2 added, and viscosity is somewhere between a conventional liquid and an SCF. GXLs are interesting options to SCFs, since more diverse physicochemical properties can be obtained from a variety of different green organic solvents, than for scCO2 with or without a cosolvent added. This has been demonstrated by Jessop et al., for methanol, ethanol, and acetone among other solvents, giving a range of obtainable dielectric properties [20]. Applications of GXLs as solvents have mainly been demonstrated in particle

158

SECTION | II Green Solvents

TABLE 7.1 Typical Ranges of Density, Viscosity, and Diffusivity for Gases, SCFs, and Liquids Substance (State of Matter)

Density (g/mL)

Viscosity (cP)

Diffusivity (mm2/s)

Gas

1023

0.01

200

SCF

0.21.0

0.020.1

0.010.1

Liquid

1.0

0.32.0

0.001

TABLE 7.2 A Selection of SCFs [2] SCF

Tc (
C)

Pc (MPa)

Other Aspects

Carbon dioxide

31

7.4

Low polarity

Water

374

22.1

Corrosive with dissolved oxygen

Ethane

32

4.9

Flammable

Propane

97

4.3

Extremely flammable

Ethylene

9

5.0

Flammable

Methanol

239

8.1

Flammable

Ethanol

241

6.1

Highly flammable

Toluene

319

4.1

Highly flammable

Sulfur hexafluoride

46

3.8

Potent greenhouse gas

Dinitrogen monoxide (nitrous oxide)

33

7.4

Enhances combustion of other substances

Ammonia

132

11.3

Flammable, toxic

formation, polymer processing and in homogeneous and heterogeneous catalysis (see an excellent review article by Jessop and Subramaniam [21]), but only in a few cases as extraction solvent [2226]. Using strict definition, a GXL is produced by adding compressed liquid CO2 to an organic solvent and letting the mixture reach equilibrium conditions. At vapor-liquid equilibrium (VLE), there should be a dense liquid phase containing dissolved CO2, and a vapor phase with mainly CO2 and some organic solvent vapor. It is the liquid phase of this system that could be used as a GXL solvent in separation processes. To emphasize that there are two phases in equilibrium (VLE), this type of GXL is called “two-phase GXL” in this chapter (Fig. 7.3).

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Supercritical Fluids and Gas-Expanded Liquids Chapter | 7

P1

P2 P3

Vapor Increase P

Add CO2 Two-phase GXL

Liquid

(A)

(B)

One-phase GXL (C)

FIGURE 7.3 Schematic drawing showing (A) organic liquid solvent at ambient pressure; (B) compressed liquid CO2 is added, giving a liquid phase (two-phase GXL) in equilibrium with a vapor phase; and (C) pressure is increased (in this case by decreasing the volume) to a magnitude that phase transfer occurs giving only one liquid phase (one-phase GXL).

The same or similar physicochemical properties can however be obtained at higher pressure in a one phase liquid phase in which the organic liquid contains significant amount of dissolved compressed liquid CO2. In this chapter, we will still consider such compressed fluid as a GXL, although there is no vapor phase in equilibrium with the liquid phase due to the higher pressure. The current situation is that there is no proper terminology to describe such liquid; in the literature it is called “compressed fluid,” “binary fluid,” “enhanced fluidity liquid,” or “subcritical GXL.” In this chapter, we will call this fluid a “one-phase GXL” (Fig. 7.3). It would be good if the International Union of Pure and Applied Chemistry (IUPAC) could come with a recommendation of a proper terminology, for both types of GXLs. The aim of this chapter is to describe SCFs and GXLs in terms of their physical properties (density, compressibility), mass transfer properties, dielectric properties, phase equilibria, and solubility. Experimental methods to produce SCFs and GXLs will be described, as well as methods to characterize the compressed fluids. Finally, aspects on greenness of SCFs and GXLs will be discussed.

7.2 PHYSICOCHEMICAL PROPERTIES OF SCFS In this section, properties of SCFs will be described in more detail, with focus on neat (pure) scCO2. However, similar physicochemical properties are found for SCFs containing small volume percentage of an organic cosolvent, as long as the mixture is in SCF state.

160

SECTION | II Green Solvents

7.2.1 Density As already mentioned, a relatively small change in pressure of an SCF leads to a relatively large change in density, as compared to in liquids. Fig. 7.4 shows that more the density of an SFC is influenced by pressure the closer it is to the CP [27]. This means that an SCF closer to the CP is more compressible than an SCF further away from this point. As a consequence, high- and low-density regions (i.e., density fluctuations) have been observed near the CP. In Table 7.3, density values are shown for scCO2 as a function of temperature and pressure [28]. Clearly, the higher the temperature is, the higher the pressure is needed to achieve the same high density. A density of 0.80 g/mL and higher is marked with gray shading in Table 7.3, which is valuable in applications where relatively high density is needed. One such example is in the extraction of lipids [5]. In general, density is important because the solubility of many compounds in scCO2 strongly depends on the density of the SCF, which will be further discussed in Section 7.5. Density of scCO2 can be calculated by an equation of state (EOS), for instance, the PengRobinson EOS [29]. EOS is described in more detail below in Section 7.4. Wang et al. [30] have compared different EOSs for the correlation of scCO2 density. The authors have showed that when density increases linearly with pressure (reduced parameter), most of the predictions using EOS were accurate, while in the region where density increases TRef = TT = 0.8 C 0.9 2.0

1.0

PR = PP C

1.1 1.2 1.0

CP 1.55

0

0.1

1.0 PR = PP C

10.0

FIGURE 7.4 Variation in the reduced density of a pure component (e.g., CO2) in the vicinity of its critical point, as a function of reduced pressure for different isotherms. Reduced parameters are used to enable a comparison of different SCFs. Reprinted from McHugh, M.A., Krukonis, V.J., 1994. Supercritical fluid extraction. Principles and practice, In: Brenner, H. (Ed.). Butterworth-Heinemann with permission from Butterworth-Heinemann (Elsevier).

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161

TABLE 7.3 Density in Milligram per Milliliter for scCO2 as a Function of Pressure up to 60 MPa and Temperature up to 100
C T (
C)

P (MPa) 7.4

10

20

30

40

50

60

31

0.42

0.76

0.89

0.94

0.99

1.02

1.04

40

0.23

0.62

0.83

0.91

0.96

0.99

1.02

50

0.19

0.38

0.78

0.88

0.92

0.96

0.99

60

0.17

0.29

0.72

0.83

0.89

0.93

0.97

70

0.16

0.25

0.66

0.78

0.86

0.90

0.94

80

0.15

0.22

0.59

0.74

0.82

0.88

0.92

90

0.14

0.21

0.54

0.70

0.79

0.85

0.89

100

0.13

0.19

0.48

0.66

0.76

0.82

0.87

Gray shaded area shows density values of 0.80 g/mL or higher. Source: Data from Peace Software. http://www.peacesoftware.de/einigewerte/co2_e.html (accessed November, 2015).

drastically with a change in pressure, the accuracy of the density predictions is low. EOS developed exclusively for scCO2 can also be found in the literature, such as the one developed by Span and Wagner [31]. Similar evaluation of different EOSs for the prediction of the density of scCO2 was done by B¨ottcher et al. [32]. Significant difference between the considered EOS was also observed by the authors [32]. Correlations using empirical equations (in some cases polynomials) can be found in the literature for scCO2 density such as in Bahadori et al. [33] and Wang et al. [30]. A large database containing experimental data for the density of scCO2 is found online in the NIST Chemistry Webbook [34]. Since a large number of experimental data were considered in the correlations used in the website, the values of density for certain conditions of temperature and pressure present good accuracy. Furthermore, there is software available online to carry out such calculations, for instance, the Peace Software [28].

7.2.2 Viscosity and Mass Transfer In extraction, mass transfer is a key factor controlling the kinetics of the process, while in chromatography, mass transfer affects the efficiency of the separation. The slowest and rate-limiting type of mass transfer is diffusion. SCFs enable high mass transfer rates due to the low viscosity and fast diffusivity. It is important to distinguish between self-diffusion (also known as Brownian motion or tracer diffusion), which according to IUPAC’s definition

162

SECTION | II Green Solvents

is the diffusion coefficient (D i) of species (molecules) i in the absence of a chemical potential gradient [35]; and binary diffusion (also called chemical diffusion) (D12), which describes the molar flux of a compound in a media (solvent) occurring due to a concentration (or chemical potential) gradient of the compound. Self-diffusion coefficients can be described by the StokesEinstein equation and its variations (infinite dilution condition, low Reynold number): Di 5

κB T 6πηr

ð7:1Þ

where κB is Boltzmann’s constant, T is the system temperature, η is the dynamic viscosity, and r is the radius of the spherical particle. A common equation used to correlate the self-diffusion coefficients based on the StokesEinstein equation is the WilkeChang [36] and its variations. The equation of WilkeChang is as follows: ð7:4:10215 ÞTMS 0:6 ηVeb

1=2

Di 5

ð7:2Þ

where T is the temperature, MS is the molecular weight, η is the solvent viscosity, and Veb is the molar volume at the vaporization temperature and atmospheric pressure. Other equations for the calculation of self-diffusion coefficients are available in the literature. For example, Olesik [37] has used the method of Giddings and Seager [38] for analytes near to the infinite dilution. Coelho et al. [39] have used perturbation theory and molecular dynamics. Liu and Macedo [40] have considered empirical equations based on the density expansion for self-diffusion coefficients calculation. Sua´rez-Iglesias et al. [41] have used the relationship between viscosity and self-diffusion for the calculation of the self-diffusion coefficients. Moreover, the rough hard sphere model and its modifications have also been used for the calculation of diffusion in neat SCFs by different authors [42]. Binary diffusion of a compound moving a certain distance in a media is described by Fick’s first law of diffusion: J 5 2 D12

@C @x

ð7:3Þ

where J is the diffusion flux per area, D12 is the binary diffusion coefficient, C is the concentration of the compound that is diffusing, and x is the perpendicular position (length). Because the concentration is generally not constant with time, mathematical modifications were proposed for this equation, creating Fick’s second law of diffusion: @C @2 C 5 D12 2 @t @x

ð7:4Þ

Supercritical Fluids and Gas-Expanded Liquids Chapter | 7

163

When the binary diffusion coefficient varies with time, which is often the case in, e.g., extraction, the following equation can be used:
 @C @ @ 5 D12 ð7:5Þ @t @x @x Experimental diffusion coefficients are scarce in the literature due to technical difficulties of the measurements [41,43,44]. One of the most used methods for measurement of the binary diffusion coefficients in SCFs is the Taylor dispersion method, in which dispersion of a solute in a laminar flow through a capillary tube is measured. Using this method, the binary diffusion coefficient of acetone, benzene, naphthalene, 1,3,5-trimethylbenzene, phenanthrene, pyrene, and chrysene have been measured in scCO2 by Sassiat et al. [36]. Clearly, the diffusion coefficients decrease with increasing molecular weight and size (i.e., molar volume), just as expected. One exception is heavy isotopes, which have higher diffusion coefficients due to their lower molar volume (bonds are shorter), as compared to their normal isotopes’ counterparts [45]. Further, the diffusion coefficients decrease with increasing density and viscosity of the SCF. Effects of temperature is marginal—the diffusion coefficients of scCO2 only increase with 10% when temperature increases from 30
C to 60
C [36]. This is in contrast to liquids, for which diffusion coefficient almost double in the same temperature range [46]. Fig. 7.5 shows the self-diffusivity of scCO2 as a function of temperature for different isobars. Viscosity is an important mass transfer property, and as shown in Eqs. (7.3) and (7.4), it is inversely proportional to the diffusion coefficient. Viscosity is also an important property for the prediction of pressure drops and heat transfer rates in processes. Experimental data for viscosity of scCO2 can be found in the literature such as in Pensado et al. [47], Vesovic et al. [48], and in the NIST Webbook [34]. Viscosity of scCO2 increases with an increase in pressure and a decrease in temperature [47] (Fig. 7.6). With the addition of an organic solvent, such as methanol or ethanol, the viscosity of the mixture increases in comparison with pure scCO2 at the same conditions of temperature and pressure [4952]. For the calculation of viscosity of scCO2, EOS and correlated models can be used [53]. The viscosity of a fluid can be expressed in function of three different terms [54,55]: ηðT; ρÞ 5 η0 ðTÞ 1 ΔηðT; ρÞ 1 ΔηC ðT; ρÞ

ð7:6Þ

where η0 ðTÞ is the dilute gas viscosity at determined temperature (zero-density limit), ΔηðT; ρÞ is the excess viscosity that determines the effect of elevated pressures, and ΔηC ðT; ρÞ is the critical enhancement that considers the fluctuations near the CP. The last term, ΔηC ðT; ρÞ, can be neglected for applications away from the CP [55]. The same principle of the three terms (low-density limit and high-density limit) was also applied to self-diffusion coefficients by Boned et al. [56].

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SECTION | II Green Solvents

10–2

Diffusivity (cm2/s)

10–3

10–4

Pressure (bar) 70 80 100 Saturated vapor Critical point

150 200

Saturated liquid

Typical diffusivities of solutes in normal liquids

10–5

0

20

40

60

80

100

Temperature (°C) FIGURE 7.5 Self-diffusion coefficients (D, cm2/s) of scCO2 as a function of temperature for different isobars. Reprinted from McHugh, M.A., Krukonis, V.J., 1994. Supercritical fluid extraction. Principles and practice, In: Brenner, H. (Ed.). Butterworth-Heinemann with permission from Butterworth-Heinemann (Elsevier).

7.2.3 Dielectric Properties Polarity is a microscopic property quantifying the dipole moment of a molecule, and has the unit Debye (1 debye 5 3.3 3 10230 C m). Polarity is unaffected by changes in temperature. The relative static permittivity (dielectric constant) on the other hand, is a macroscopic property describing how easily a solvent can become polarized, i.e., how easily an electric field can be established across the bulk of the solvent when exposed to an electric field. The relative static permittivity (εr) can be measured by a capacitor, relating the capacitance in vacuum (C0) to the capacitance in the liquid (Cx): εr 5

C0 Cx

ð7:7Þ

The relative static permittivity depends on the temperature, as well as on the pressure. One of the most used equations for the relative static permittivity

Supercritical Fluids and Gas-Expanded Liquids Chapter | 7

165

0.12 0.11 0.10

Viscosity (CPS)

0.09

37°C

0.08 77°C

0.07 0.06

47°C 0.05 0.04

PC

0.03 0.02 0.01 40

100 Pressure (bar)

1000

FIGURE 7.6 Dynamic viscosity of scCO2 as a function of pressure for different isotherms. Reprinted from McHugh, M.A., Krukonis, V.J., 1994. Supercritical fluid extraction. Principles and practice. In: H. Brenner (Ed.). Butterworth-Heinemann with permission from ButterworthHeinemann (Elsevier).

for polar compounds considering the temperature and pressure dependence and based on molecular properties is the Kirkwood-Frohlich [57]: ðεr 2 εN Þð2εr 1 εN Þ M 4πNA 2 5 gμ 2 ρ 9kB T ðεr 12Þ εr

ð7:8Þ

where εN is the infinite relative permittivity, M is the molecular weight, ρ is the density, NA is Avogadro’s number, kB is Boltzmann’s constant, T is the temperature of the system, μ is the dipole moment of the fluid molecule in the vacuum, and g is the Kirkwood parameter. The parameter g measures the local order among molecules (nearest neighbors) and can be calculated as:   g 5 1 2 z cosψ ð7:9Þ where z is the coordinator number and ψ is the angle between the test dipole and the neighbor. The infinite relative permittivity (εN ) can be calculated using the ClausiusMossotti equation. As shown in Fig. 7.7, CO2 has a dipole moment of zero. When a cosolvent is used, the intermolecular interactions, e.g., hydrogen bonding formed in the mixture leads to an increase in density of the mixture and a decrease in the total molar volume of the mixture. When the pressure increases, less space between the molecules is available, which creates more interaction

166

SECTION | II Green Solvents

FIGURE 7.7 Carbon dioxide has zero dipole moment, but it is a quadrupole that interacts with dipoles.

between the molecules and a larger number of hydrogen bonds [58]. However, when temperature increases, the intermolecular interaction decreases significantly, which can be explained by the breaking of the hydrogen bonds [58,59]. Liquid and scCO2 have lower relative static permittivity than hydrocarbons [60]. For instance, scCO2 at 10 MPa and 80
C has a relative static permittivity of 1.11, while hexane at ambient conditions has a relative static permittivity of 1.88 [61]. Even when taking density into account, CO2 has a significantly lower relative static permittivity than hydrocarbons such as hexane or heptane (Table 7.5). Further, CO2 has a quadrupole moment that enables dipolequadrupole interactions with other molecules, and it acts both as a weak Lewis acid and a Lewis base. An example of the dielectric properties of scCO2 is its higher solubility of water in comparison to carbon monoxide. The explanation for this effect is the hydrogen bonding between the oxygen in CO2 and the hydrogen in water [60].

7.3 PHYSICOCHEMICAL PROPERTIES OF GXLS Information about physicochemical properties of GXLs in the literature is rather scattered. It is fairly easy to find information about GXLs using strict definition, i.e., two-phase systems within the binary fluid phase envelope. It is, however, more difficult to find information about one-phase GXLs, since there is no clear definition of such binary fluid, and the search terms are not that obvious. In this chapter, we are making an attempt to gather the most relevant information about physicochemical properties of both one-phase and two-phase GXLs. In general, GXLs are nonideal mixtures, which bring challenges in the predictability of their physicochemical characteristics [62]. Jessop and Subramaniam [21] classified GXLs into different classes according to their properties. Class 1 liquids are defined as solvents that do not dissolve compressed liquid CO2, like water, and consequently do not change their properties significantly with exception of acidity. Class 1 GXLs will not be considered here. Class 2 liquids are solvents that dissolve compressed liquid

167

Supercritical Fluids and Gas-Expanded Liquids Chapter | 7

CO2 such as methanol and hexane, e.g., and as a consequence, their physicochemical properties change significantly. Class 3 liquids dissolve only moderate amounts of CO2, e.g., ionic liquids, polymers, and crude oil. For this class, some properties change moderately such as volume expansion and polarity, but viscosity can change significantly [21].

7.3.1 Density In separation processes, using a GXL as the solvent, it is important to know which change in density that is obtained when varying pressure, temperature, and composition. If the liquid in a two-phase GXL is used as a solvent, it is the density of the liquid phase that is relevant. A two-phase GXL has the advantage of being a rather robust VLE system as long as T and P are fixed, since changes in the overall composition does not affect the composition in each phase [6370] (Fig. 7.8). Pressure in such system is often increased simply by adding more compressed liquid CO2. In this case, what does change is the volume of the liquid phase, which naturally affects the density. An example of a two-phase GXL is given for CO2/ethanol at 7 MPa and 60
C, in which the composition of each phase is marked with an arrow (Fig. 7.8). On the other hand, for a one-phase GXL the situation is different since such liquid is not in equilibrium with a vapor (“liquid” regime in Fig. 7.8). In this case, density is a result of pressure, temperature, and composition, and all these variables need to be perfectly controlled. 18.0 SCF

16.0

Pressure (MPa)

14.0

CP Liquid

12.0

80°C e

lin int

o

10.0

ep bbl

60°C

Bu

8.0

40°C 6.0

Dew point line

4.0 VLE

2.0 0.0 0.0

Vapor 0.1

0.2

0.3

∗

0.4 0.5 0.7 0.6 Mole fraction CO2 (x1)

0.8

0.9

∗∗ 1.0

FIGURE 7.8 Phase equilibrium for a binary mixture of CO2 and ethanol at 40, 60, and 80
C.  Mole fraction of CO2 in the liquid phase and  mole fraction of CO2 in the vapor phase for a certain pressure and temperature (in this case 7 MPa and 60
C). Data from Refs. [6370].

168

SECTION | II Green Solvents

The volume expansion of a liquid in equilibrium with a vapor phase has been studied before by different authors [7175]. The volume expansion of mixtures containing CO2 and alcohols was obtained considering the experimental values of density by Aida et al. [73]. For mixtures containing alcohol and CO2, the increasing volume expansion is less for lower concentrations of CO2 (x , 0.6) and larger for higher concentrations (x . 0.6) [73]. The authors [73] showed that the partial molar volume of alcohols is negative at high concentration of CO2 (x . 0.8), and the same was observed in mixtures with other organic compounds. (The partial volume of a compound is the change in total volume due to the variation in composition/moles of CO2 added.) This can be explained by the fact that by adding more CO2, the space between the alcohol molecules is filled with the CO2 molecules, which results in the breaking of the hydrogen bonds present in the mixture. Similar conclusion was done by Souvignet and Olesik [76] for mixtures of methanol and CO2. This effect can also be explained by the approach of the CP or phase transition. For higher pressures, where a single-phase liquid is present in the system, P¨ohler and Kiran [71] have determined the volume expansion of mixtures containing CO2 and ethanol. This was done by using a variable-volume vessel connected with a sensor to determine the volume variation. A ferromagnetic slug was connected to the piston and its position was detected by a linear variable differential transformer [77]. The study showed that the volume increases with pressure and decreases with temperature [71]. Moreover, as Kordikowski et al. [72] and P¨ohler and Kiran [71] observed, the expansion of the liquid phase is independent of the type of organic solvent (acetonitrile, 1,4-dioxane, ethyl acetate, dimethyl sulfoxide (DMSO), N,N-dimethylformamide (DMF), and ethanol) when increasing the molar fraction of CO2. However, Abbot et al. [78] observed that not all organic solvents expand in the same magnitude with the addition of CO2 due to differences in the solubility of the CO2 in the organic solvent. The same authors showed that the CO2 solubility is similar at 5 MPa for a wide range of organic solvents including acetone, ethanol, and propan-1-ol, e.g. The volume expansion is shown in Fig. 7.9 for a range of CO2 compositions (molar fraction) in acetone [79], propan-2-ol [74], ethyl acetate [72], and methanol [80]. The volumetric properties of other solvents were studied in binary mixtures with CO2, such as ethyl acetate [81], pentane [82], sulfur hexafluoride [77], toluene [83], and acetone [84]. Other types of SCFs were also considered in the study of volume expansion, such as ethane [72]. The excess molar volume, i.e., the difference between the ideal (sum of volume of two compounds) and the value of volume obtained for a mixture, can be obtained by: X VmE 5 V mixture 2 x i Vi ð7:10Þ

Supercritical Fluids and Gas-Expanded Liquids Chapter | 7

169

Volume expansion (%)

1200 1000 800 600 400 200 0 0.0

0.2

0.4 0.6 Mole fraction CO2 (x1)

0.8

1.0

FIGURE 7.9 Volume expansion (%) of binary mixtures containing: KCO2 (1) 1 acetone (2) at 40
C, V CO2 (1) 1 2-propanol (2) at 40
C, ’ CO2 (1) 1 ethyl acetate (2) at 40
C and ▲ CO2 (1) 1 methanol (2) at 35
C. Data from Refs. [72,74,79,80].

where V mixture is the value of volume obtained for a mixture, xi and Vi are the composition (molar fraction) and the molar volume of each solvent compound i in the mixture, respectively. The excess molar volume can also be expressed in terms of density: X VmE 5 xi Mi ðρ21 2 ρ21 ð7:11Þ i Þ i51

where ρ is the density of the mixture, Mi is the molar mass of the each solvent compound i, and ρi is the density of each solvent compound i. In the cited examples of experimental data available in the literature for density of mixtures containing CO2, inside of the phase diagram (two phases), it is possible to notice a slightly increasing density with increasing pressure and molar fraction of CO2 [73,85] (Fig. 7.10). Furthermore, above a certain pressure the density drops suddenly. This is when the bubble point line is reached (Fig. 7.8) and there is a phase transition from a two-phase system (VLE) to a one-phase liquid. During the transition, the densities of the two phases rapidly approach one another to become one homogeneous liquid solution. This density change can also be observed in Fig. 7.11, in which density is plotted versus pressure for CO2/ethanol mixtures of different molar fractions [71]. With increasing pressure, density is rapidly increased when being just above the bubble point line, however, moving further away from the phase change regime, the change in density with pressure is less dominant. It can also be seen that the higher the molar fraction of CO2 to ethanol, the larger the change in density with pressure (i.e., higher compressibility).

170

SECTION | II Green Solvents

FIGURE 7.10 Density scheme for (VLE) mixtures containing CO2 and alcohol at 40
C: K CO2 1 2-propanol, V CO2 1 methanol and ▲ CO2 1 ethanol. Mole fractions of CO2 range between 0.1 (pressure 1.84 MPa) and 0.9 (pressure 7.75 MPa) for CO2 1 2-propanol, between 0 (0 MPa) and 0.62 (7.55 MPa) for CO2 1 methanol, and between 0 (0 MPa) and 0.62 (7.72 MPa) for CO2 1 ethanol. Plotted with data from Yaginuma, R., Nakajima, T., Tanaka, H., Kato, M., 1997. Densities of carbon dioxide plus 2-propanol at 313.15 K and pressures to 9.8 MPa, J. Chem. Eng. Data, 42 814816 and Aida, T., Aizawa, T., Kanakubo, M., Nanjo, H., 2010 Relation between volume expansion and hydrogen bond networks for CO2-alcohol mixtures at 40
C, J. Phys. Chem. B, 114 (2010) 13628-13636.

FIGURE 7.11 Density scheme for one-phase GXL mixtures containing CO2 1 ethanol at 50
C, pressure ranges between 8.02 and 61.26 MPa, and different compositions of CO2: V wCO2 5 0.5, ¨ ’ wCO2 5 0.7, ▲ wCO2 5 0.8, s wCO2 5 0.9, and K wCO2 5 1.0. Plotted with data from Pohler and Kiran [71].

7.3.2 Compressibility Another interesting property is the isothermal compressibility, κ, which is the fractional change in volume of a system as the pressure changes at constant temperature (i.e., the increase in pressure when the volume is decreased). An SCF is in general far more compressible than a liquid, as

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171

TABLE 7.4 Compressibility Factors (κT ) for SCFs, GXLs at Different Mole Fractions of Solvent/CO2 and a Conventional Liquid Solvent, All Data at a Temperature of 35
C Solvent

P (MPa)

κT (MPa21)

Liquid ethyl acetate

10

 0.0013

GXL (ethyl acetate/CO2, 72/28)

10

 0.0017

GXL (ethyl acetate/CO2, 57/43)

10

 0.0026

GXL (ethyl acetate/CO2, 33/67)

10

 0.0058

scCO2 (near CP, 0.7 g/mL)

10

 0.0527

scCO2 (1.0 g/mL)

40

 0.0038

Source: Data from Falco, N., Kiran, E, 2012. Volumetric properties of ethyl acetate plus carbon dioxide binary fluid mixtures at high pressures. J. Supercrit. Fluids. 61, 924; Velasco, I., Rivas, C., Martinez-Lopez, J.F., Blanco, S.T., Otin, S., Artal, M., 2011. Accurate values of some thermodynamic properties for carbon dioxide, ethane, propane, and some binary mixtures, J. Phys. Chem. B 115, 82168230.

0.40

K(atm–1) (E-2)

0.30

0.20

0.10 0.00 70 80 90 100 110 120 130 140 150 160 170 180 190 200 Pressure (atm) FIGURE 7.12 Isothermal compressibility factor, κ, of methanol/CO2 mixtures as a function of pressure: (’) methanol/CO2 (10/90, mole fraction), (&) methanol/CO2 (20/80, mole fraction), (K) methanol/CO2 (30/70, mole fraction), (1) methanol/CO2 (40/60, mole fraction). Reprinted from Souvignet, I., Olesik, S.V., 1995. Solvent-solvent and solute-solvent interactions in liquid methanol/carbon dioxide mixtures, J. Phys. Chem., 99, 1680016803 with the permission from ACS Publications.

evident in Fig. 7.4. In Table 7.4, κ values for a selection of SCFs and GXLs are listed and compared with a typical liquid [81,86]. For SCF and GXL binary mixtures, the isothermal compressibility increases with increasing CO2 content [76,81], as shown in Table 7.4 for ethyl acetate [81] and in Fig. 7.12 for methanol [76]. The same was observed for mixtures of scCO2 and propane and octane [87], and acetone [88].

172

SECTION | II Green Solvents

The isothermal compressibility decreases with increasing pressure. However, close to the critical region, the isothermal compressibility of pure components presents a different behavior, having a maximum value. For instance, scCO2 has its maximum isothermal compressibility at 8 MPa for 35
C (close to the CP), for a range of pressures between 5 and 13 MPa [89]. This is quite obvious due to the rapid change in density near the CP as explained above. The same is observed for mixtures of fluids such as methane 1 CO2 [90] or water 1 CO2 [90] or for very small concentrations of cosolvent, e.g., 0.01 molar fraction of ethanol in scCO2 [91]. For GXL mixtures containing scCO2 and solvent (e.g., ethyl acetate or acetone), the isothermal compressibility decreases with an increase in pressure as well as with a decrease in temperature [81,87,88]. However, the change in compressibility of a mixture levels off at a certain pressure, i.e., the isothermal compressibility values become almost constant (Fig. 7.12). The isothermal compressibility of SCFs and GXLs can be calculated as a function of the molar volume or density, as follows [81]:

 1 @Vm 1 @ρ 5 ð7:12Þ κT 5 2 Vm @P T ρ @P T When the content of CO2 is high, i.e., the binary mixture is an SCF, large variations in isothermal compressibility are observed with relatively small changes in pressure [81]. This has a large impact on separation processes, especially in chromatography. A large compressibility leads to large density fluctuations and thereby large local variations in solubility of the target compounds. In SFC separation methods, where gradient elution is commonly used in which the mobile phase initially consists of neat scCO2 or scCO2 containing only small volume percentage of a cosolvent, changing the flow rate will affect the inlet pressure and thereby have a large impact on the density of the mobile phase [92]. At higher volume percentage of cosolvent, the mobile phase is a one-phase GXL, and compressibility is lower, resulting in less effects in terms of density when changing the flow rate—just like in conventional high-performance liquid chromatography (HPLC). However, a gradual increase in cosolvent amount also increases the viscosity and the pressure drop over the column, naturally bringing other disadvantages that are well known for HPLC. A tutorial on ultrahigh-performance SFC by Novakova et al. is found in Ref. [93] and a review by Guiochon and Tarafder on preparative SFC is found in Ref. [11].

7.3.3 Viscosity and Mass Transfer Viscosity and mass transfer properties of GXLs are somewhere between those for conventional liquids and SCFs. In general, the same relations can be used to describe viscosity and diffusion coefficients for GXLs as for conventional liquids and SCFs.

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173

The viscosity of CO2-expanded methanol was studied by Sih et al. [51] for temperatures between 25
C and 40
C and pressure between 0.1 and 7.0 MPa. Not surprisingly, the viscosity decreased with increasing temperature and with the addition of CO2. The viscosity reduction seemed to be linear for composition of CO2 in the mixture higher than 0.5 mole fraction of CO2 [51]. The same research group [52] has also measured the viscosity of CO2-expanded acetone. Kariznovi et al. [94] has measured the viscosity and density of binary mixtures containing CO2 and methanol, ethanol, or propan-1-ol. The self-diffusion coefficients (referring to no concentration/chemical potential gradient) for alcohols and CO2 was calculated by Aida et al. [73] using the StokesEinstein equation variation (Eq. 7.1). The self-diffusion coefficients increased linearly for low concentrations of CO2, but a fast increment was observed at higher concentrations (x . 0.6) [73]. The selfdiffusion coefficients decreased with increasing alkyl chain number [73], and for the same molecular weight compound, branched alkyl alcohols presented higher self-diffusion coefficients than linear ones. There are few studies in the literature that present direct measurements of mass transfer coefficients in GXLs. For instance, the binary diffusion of benzene was studied by Sassiat et al. [36] for the entire range of CO2 concentrations in methanol. Similarly, the binary diffusion of benzonitrile in CO2-expanded ethanol was studied by Li and Tan [95]. It was observed that in CO2-expanded ethanol, higher diffusion coefficients were obtained in comparison with pure ethanol or methanol, and the diffusion coefficients increased with increasing CO2 concentration. This behavior is shown in Fig. 7.13 for the binary diffusion coefficients of benzonitrile in CO2expanded ethanol [95]. Further, the binary diffusion coefficients decrease

FIGURE 7.13 Binary diffusion coefficients of benzonitrile in CO2-expanded ethanol at 40
C and different compositions (mole fractions) of CO2: & x1 5 0.0, 3 x1 5 0.1, V x1 5 0.2, ▲ x1 5 0.3, s x1 5 0.4 and K x1 5 0.5. Plotted with data from Lin, I.-H., Tan, C.-S., 2008. Diffusion of benzonitrile in CO2-expanded ethanol, J. Chem. Eng. Data, 53, 18861891.

174

SECTION | II Green Solvents

with increasing pressure, which can be explained by the increase in the collisions between the solute and solvent [95], see also Fig. 7.10 showing the density versus pressure. Regarding the diffusion coefficient calculations for GXL, Shukla et al. [96] have used molecular dynamics simulation and TaylorAris dispersion technique to obtain the diffusion coefficient for both CO2-expanded acetone and CO2-expanded methanol. The same was done by Eckert et al. [97], which have measured transport properties of solutes in GXL using TaylorAris diffusion techniques and compared the results with molecular dynamics simulation.

7.3.4 Dielectric Properties Dielectric properties of a solvent determine the type of intermolecular interactions that can be established between the solvent and the solute molecules to be dissolved. Examples of intermolecular interactions or bonding between solvent molecules and solutes are hydrogen bonding, dipoledipole interactions, dipoleinduced dipole interactions, induced dipoleinduced dipole interactions, dipoleion interactions and induced dipoleion interactions. For ionic liquids, which are not described in this chapter, ionion interactions are also possible. Dielectric properties of a solvent can be quantified by the KamletTaft solvatochromic parameters scale [98100]. With this experimental spectrophotometric method, a chromophore (probe molecule) is added to the solvent, and the size of a shift in UV/Vis maximum absorbance is observed. The stronger the intermolecular interactions with the solvent, the larger the wavelength shift for the observed peak will be. Obtained shifts are normalized into three different scales—polarizability (π ), the α-parameter, and the β-parameter, and these numbers give a clear idea about the overall dielectric properties of the solvent, and can be used to create solvent property maps [101]. Polarizability is the capacity of the solvent to induce any kind of electrostatic interaction with a dissolved solute, which is a measure of the “polarity” of the solvent. Hence, polarizability is a macroscopic property just like the dielectric constant. Using the KamletTaft solvatochromic method, polarizability (π ) was determined using 1-methoxy-4-nitrobenzol, N, N-diethyl-3-nitroaniline, 1-methoxy-4-[(E)-2-nitroethenyl]benzene, 1-ethyl-4nitrobenzene, N-methyl-2-nitroaniline and N,N-diethyl-4-nitroaniline as suggested chromophores [100]. Cyclohexane is given the value zero and DMSO is given the value one. Most other solvents will have values in between 0 and 1. The KamletTaft α-parameter is used to scale the solvent hydrogenbond donating ability (i.e., the acidity). In this case, it has been proposed to assign the value zero for hydrocarbons, ethers, esters, tertiary amines, and N, N-substituted amides [98]. Methanol was assigned as 1 by the authors. The

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175

chromophores tested for the hydrogen-bond acceptor were 4-nitroaniline, N, N-diethyl-4-nitroaniline, 4-nitrophenol, and 1-methoxy-4-nitrobenzol [98]. The KamletTaft β-parameter is used to scale the solvent hydrogen-bond accepting ability (i.e., the basicity). In this case, it has been proposed to assign the value zero for hexane, heptane, and cyclohexane and the value 1 for hexamethylphosphoramide (HMPA) [99]. The chromophores tested for the hydrogen-bond acceptor were 1-methoxy-4-nitrobenzol, Reichardt’s dye, Brooker’s merocyanine, 4-carbomethoxy-1-ethylpyridinium iodide, bis[α-(2pyridyl benzylidene)-3,4-dimethyl aniline]bis(cyano)iron(II) (Fe(LL)2(CN)2), and N,N-diethyl-4-nitroaniline [99]. The KamletTaft parameters (π , α, and β) can be used in a linear dependence called general linear solvation energy relationship (LSER). As described by Taft et al. [102], the LSER can be used in the correlation and prediction of the solvent effects and bring more knowledge about the effects in the molecular level. The LSER is written as [103]: 

XYZ 5 XYZ0 1 sπ 1 aα 1 bβ

ð7:13Þ

where XYZ0 is the value of the considered property for the inert solvent without solvation abilities, and s, a, and b are constants that represent the sensibility (dependence) of the solvent parameters [104]. It has been used to describe reactions rates and equilibrium constants, e.g. Bulgarevich et al. [105] has reported that if the constants s, a, and b are known for the liquid state of a compound and using π , α and β for the SCF state of the same compound, the behavior of the studied property can be predicted in the SCF state. In Jessop et al. [20], solvatochromic parameters including the KamletTaft  π , α, and β for green solvents have been determined experimentally and compared with literature values. The authors observe that there is a lack of green solvents to replace more toxic/environmentally burdensome organic solvents that are aprotic (low acidity), highly polar and demonstrating either low basicity (like acetonitrile and dichloromethane) or high basicity (like DMSO and DMF). The chromophores cited by the authors [20] are 4-nitroaniline and 4nitrophenol for the β-parameter, N,N-dimethylbenzamine and 4-carbomethoxy1-ethylpyridinium iodide for the α-parameter and 1-methoxy-4-nitrobenzol and  N,N-diethyl-4-nitroanisole for polarizability, π . For protic solvents, there are several options of green solvents, for instance, water, methanol, ethanol, butan-1-ol, and glycerol. Using the solvent dielectric property map obtained by the KamletTaft solvatochromic method, Jessop et al. [101] also showed how the properties change upon the addition of compressed liquid CO2 to any of the green solvents. Figs. 7.14A and B show a selection of solvents’ KamletTaft solvatochromic parameters [20,106109]. As shown for methanol and acetone, polarizability as well as basicity decrease when adding compressed CO2 to the solvent, enabling the coverage of a larger map of dielectric properties as compared to using only neat green solvents. Similar effects of decreasing polarizability can be

176

SECTION | II Green Solvents

(A)

(B) FIGURE 7.14 KamletTaft solvatochromic parameters of the most common organic solvents (blue labels) and a selection of green solvents including GXLs (green labels). (A) Protic solvents (α-parameter . 0.5) and (B) aprotic solvents (α-parameter , 0.5). THF, tetrahydrofuran; ACN, acetonitrile; DMSO, dimethyl sulfoxide; DCM, dichloromethane; and CHCl3, chloroform. Data for these solvents and several more are found in Table 7.5. Plotted with data from Refs. [20,106109].

obtained by heating the green solvent at saturation pressure; however, an elevated temperature may bring disadvantages such as thermal degradation of the analytes/solutes in the separation process. Fig. 7.14B further shows the effect of pressure (i.e., density) on the polarizability of scCO2. LennardJones mixtures (i.e., mathematical model to describe intermolecular interactions in mixtures containing neutral atoms or molecules—such as CO2CO2 interaction) were studied using molecular dynamics by Chialvo and Debenedetti [110]. The authors [110] observed that for both attractive (e.g., CO2CO2 interaction) and repulsive mixtures (e.g., xenonxenon interaction), the number of the solute molecules is higher in the vicinity of other solute molecules in comparison with this number in the bulk. Near to the CP of mixtures containing attractive forces, there is a solvent enrichment

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177

around the solute. The density increase in this region (near the CP) leads to an enhancement of collisions between the solute molecules. In repulsive mixtures, the opposite happens—there is a solvent depletion around the solute near the CP of the solvent [110]. KamletTaft solvatochromic parameters have been measured by Bulgarevich et al. [105] for mixtures involving different alcohols (methanol, ethanol, propan-1-ol, pentan-1-ol and hexan-1-ol) and CO2. The authors [105] showed that even a small amount of cosolvent (between 0.01 and 0.05 molar fraction) changes the KamletTaft parameters. Shukla et al. [96] have used molecular dynamic simulation and found the local solvation and transport effects of different GXL and SCF solvent mixtures. The local density enhancement was less intense for GXLs in comparison with SCFs, although the trend was similar [96]. This variation in density is due to the changes in solutesolvent interactions. The KamletTaft solvatochromic parameters for GXLs (CO2 1 methanol and CO2 1 acetone) were determined by Wyatt et al. [106]. The results showed a linear relationship between the solvatochromic properties and the bulk molar fraction. The KamletTaft polarizability decreased with increasing CO2 concentration as expected. However, nearly constant values of α- and β-parameters were obtained for different concentrations of CO2. Eckert et al. [111] has showed that the polar intermolecular interactions with the solute decreases while increasing CO2 concentration in mixtures containing acetone, acetonitrile and methanol GXL, with the decline being substantial at molar fraction of CO2 of over 0.6. The authors [111] show that close to the solute (Coumarin 153 in this case), CO2-expanded methanol shows a high concentration of methanol around the acceptor side of the molecule (the side containing oxygen and fluorine atoms). In Table 7.5, physicochemical properties of a range of common solvents along with scCO2 and GXLs are listed, based on data that could be found in the literature. The table includes the properties density, viscosity, selfdiffusivity, dipole moment, dielectric constant, and KamletTaft solvatochromic parameters, and ought to be valuable in the selection of a proper solvent in separation processes.

7.4 PHASE EQUILIBRIA AND PHASE DIAGRAMS 7.4.1 Introduction to Phase Equilibria—SCFs and GXLs In separation processes, it is of primary value to have a one-phase solvent throughout the separation step. Since one of the components is a compressed gas (CO2), care has to be taken that the CO2 that is part of the solvent remains in liquid or supercritical state, since the formation of a gas phase will lead to a drop in solubility of the analytes (target compounds, also called solutes) as well as nonrepeatable data. Hence, the following section

TABLE 7.5 Physicochemical Properties of a Selection of SCFs and GXLs Solvent

Liq. CO2

scCO2

Density (g/cm3)

Viscosity Dynamic (cP)

Self-Diffusivity Di (cm2/s) 3 104

Dipole Moment (Debye)

Relative Static Permittivity

KamletTaft Parameters α

β

π

0.975 [34] (0
C, 10 MPa)

0.114 [34] (0
C, 10 MPa)

1.24 [112] (0
C, 10 MPa)

0.00

1.60 [113] (0
C, 5.06 MPa)







0.834 [34] (25
C, 10 MPa)

0.076 [51] (25
C, 10.98 MPa)

1.88 [112] (25
C, 10 MPa)

0.632 [34] (40
C, 10 MPa)

0.048 [34] (40
C, 10 MPa)

2.69 [112] (60
C, 30 MPa)

0.00

1.23 [61] (80
C, 15 MPa)

0.000

20.070 [115] (23
C, 27.4 MPa)

0.910 [34] (40
C, 30 MPa)

0.094 [34] (40
C, 30 MPa)

1.25 [114] (35
C, 8.05 MPa)

2 0.100 [106] (35
C, 20.68 MPa)

0.222 [34] (80
C, 10 MPa)

0.138 [34] (40
C, 70 MPa)

1.50 [114] (35
C, 15.03 MPa)

0.746 [34] (80
C, 30 MPa)

0.022 [34] (80
C, 10 MPa)

1.14 [114] (50
C, 8.50 MPa)

0.948 [34] (80
C, 70 MPa)

0.064 [34] (80
C, 30 MPa)

1.43 [114] (50
C, 14.47 MPa) 1.49 [114] (50
C, 22.28 MPa)

20.250 [115] (50
C, 8.82 MPa) 2 0.080 [106] (35
C, 20.68 MPa)

Acetone/CO2

0.750 [71] (80/ 20) (50
C, 11.29 MPa)

0. 257[52] (80/20) (25
C, 1.03 MPa)

1.40 [36] (30
C, 11.5 MPa)





0.210 [116] (78.4/21.6) (35
C, 1.38 MPa)

0.500 [116] (78.4/21.6) (35
C, 1.38 MPa)

0.650 [116] (78.4/21.6) (35
C, 1.38 MPa)





8.20 [78] (36.1/ 63.9) (25
C, 0.5 MPa)

0.280 [116] (21.9/78.1) (25
C, 0.5 MPa)

0.460 [116] (21.9/78.1) (25
C, 0.5 MPa)

0.432 [78] (36.1/63.9) (25
C, 0.5 MPa)

0.202 [52] (80/20) (40
C, 4.88 MPa) Acetone/CO2

Ethanol/CO2

0.855 [78] (36.1/ 63.9) (25
C, 0.5 MPa)

0.150 [52] (35/ 65) (25
C, 3.71 MPa)

0.808 [71] (50/ 50) (50
C, 11.54 MPa)

0.140 [52] (35/ 65) (40
C, 4.88 MPa)

0.677 [71] (20/ 80) (50
C, 10.16 MPa)

0.580 [117] (80/ 20) (25
C, 2.38 MPa)

0.745 [71] (50/ 50) (50
C, 9.85 MPa)

0.240 [117] (41.4/58.9) (30
C, 5.94 MPa)

0.400 [116] (21.9/78.1) (25
C, 0.5 MPa) 



3.35 [118] (21.2/ 78.8) (30.5
C, 9.8 MPa)

0.535 [119] (36.4/64.6) (25
C, 5 MPa)



8.86 [78] (31.5/ 68.5) (25
C, 0.5 MPa)

0.291 [78] (31.5/68.5) (25
C, 0.5 MPa) 0.537 [119] (36.4/ 64.6) (25
C, 5 MPa)

0.822 [78] (31.5/ 68.5) (25
C, 0.5 MPa) Ethyl acetate/CO2

0.873 [120] (20/ 80) (40
C, 10.25 MPa)



1.59 [121] (35
C, 10.28 MPa)











0.896 [120] (80/ 20) (40
C, 10.28 MPa)

(Continued )

TABLE 7.5 (Continued) Solvent

Methanol/CO2

Density (g/cm3)

0.811 [73] (77.3/ 22.7) (40
C, 4.01 MPa)

Viscosity Dynamic (cP)

Self-Diffusivity Di (cm2/s) 3 104

Dipole Moment (Debye)

Relative Static Permittivity

KamletTaft Parameters α

β

π

0.426 [51] (80/ 20) (25
C, 5.62 MPa)

2.91 [122] (77.1/ 22.91) (50
C, 11 MPa)



23.30 [123] (78.6/ 21.4) (35
C, 9.51 MPa)

1.000 [106] (77.4/22.6) (35
C, 3.45 MPa)

0.660 [106] (77.4/22.6) (35
C, 3.45 MPa)

0.510 [106] (77.4/ 22.6) (35
C, 3.45 MPa)

7.74 [122] (18.9/ 81.1) (50
C, 11 MPa)



5.10 [123] (26.7/ 73.3) (35
C, 10.34 MPa)

0.980 [106] (28.8/71.2) (35
C, 6.72 MPa)

0.490 [106] (28.8/71.2) (35
C, 6.72 MPa)

0.598 [78] (32.5/ 67.5) (25
C, 0.5 MPa)

1.170 [107]

0.470 [107]

1.090 [107]

0.352 [51] (80/ 20) (40
C, 7.59 MPa) Methanol/CO2

0.799 [78] (32.5/ 67.5) (25
C, 0.5 MPa)

0.151 [51] (25/ 75) (25
C, 5.56 MPa) 0.117 [51] (25/ 75) (40
C, 3.5 MPa)

Water

0.997 [124]

0.892 [124]

0.122 [125] (365
C, 18.53 MPa) Glycerol Methanol

Ethanol

11.68 [78] (32.5/ 67.5) (25
C, 0.5 MPa) 2.43 [126]

1.85 [128]

1.12 [127] (400
C, 29.1 MPa)

1.258 [130]

80.00 [128]

0.280 [106] (28.8/ 71.2) (35
C, 6.72 MPa)

15.66 [129] (350
C, 30 MPa)

875.0 [130]

1.03 [131]

2.67 [132]

40.10 [130]

1.210 [107]

0.510 [107]

0.620 [107]

0.772 [73] (40 C, 0.1 MPa)

0.542 [51] 0.460 [51] (40
C, 0.1 MPa)

2.27 [126]

1.70 [128]

33.00 [128]

0.980 [107]

0.660 [107]

0.600 [107]

37.50 [123] (35
C, 10 MPa)

0.638 [78] (25
C, 0.5 MPa)

0.785 [71] (50
C, 11.83 MPa)

0.993 [133] (30
C, 0.1 MPa)

1.01 [126]

24.30 [128]

0.860 [107]




1.69 [128]

0.535 [78] (25
C, 0.5 MPa)

0.372 [78] (25
C, 0.5 MPa) 0.750 [107]

0.540 [107] 0.291 [78] (25
C, 0.5 MPa)

Propan-1-ol

Propan-2-ol

0.787 [73] (40
C, 0.1 MPa)

1.703 [133] (30
C, 0.1 MPa)

0.65 [126]

0.781 [134]

2.049 [134]

0.65 [126]

1.68 [128]

20.10 [128]

0.840 [107]

0.900 [107]

0.336 [78] (25
C, 0.5 MPa)

0.593 [78] (25
C, 0.5 MPa) 1.62 [135]

19.25 [136]

0.760 [107]

0.520 [107]

0.840 [107]

0.606[78] (25
C, 0.5 MPa)

0.480 [107] 0.346 [78] (25
C, 0.5 MPa)

0.793 [73] (40
C, 0.1 MPa)

2.237 [133] (30
C, 0.1 MPa)

0.50 [126]

Ethyl acetate

0.895 [134]

0.424 [134]



1.78 [128]

6.02 [128]

0.000 [107]

0.450 [107]

0.550 [107]

Ethyl lactate

1.028 [137]

2.398 [138]



2.40 [139]

15.70 [137]

0.690 [20]

0.520 [20]

0.820 [20]

Acetone

0.770 [71] (50
C, 11.83 MPa)

0.308 [52] (25
C, 0.1 MPa)

0.45 [140]

2.98 [128]

20.70 [128]

0.08 [107]

0.480 [107]

0.670 [107]

PEG (polyethylene glycol)

1.111 [141]

17.134 [141]

0.01 [142]

2.36 [139]

41.20 [143]

0.792 [144]

0.510 [145]

0.932 [146]

Acetonitrile

0.776 [147]

0.346 [147]

0.42 [140]

3.92 [128]

36.60 [128]

0.190 [107]

0.400 [107]

0.750 [107]

Toluene

0.862 [147]

0.556 [147]

0.23 [148]

0.38 [139]

2.38 [149]

0.000 [107]

0.110 [107]

0.540 [107]

Heptane

0.680 [150]

0.396 [150]

0.32 [151]

0.00

1.92 [149] (20
C)

0.000 [107]

0.000 [107]

0.000 [107]




0.200 [107]

0.100 [107]

0.580 [107]




Butan-1-ol

Chloroform

1.476 [152]

1.66 [128]

17.40 [128]

0.840 [107]

0.840 [107]

0.470 [107] 0.439 [78] (25
C, 0.5 MPa)

0.723 [78] (25
C, 0.5 MPa)

0.273 [52] (40
C, 0.1 MPa)

0.495 [152]

0.25 [140]

1.04 [139]

4.81 [149] (20 C)

Dichloromethane

1.316 [152]

0.380 [152]

0.35 [140]

1.60 [139]

9.08 [149] (20 C)

0.130 [107]

0.100 [107]

0.820 [107]

DMSO

1.106 [153]

2.213 [153]

0.07 [140]

3.96 [128]

47.20 [128]

0.000 [107]

0.340 [107]

1.000 [100]

THF

0.886 [153]

0.538 [153]

0.25 [154]

1.63 [128]

7.52 [128]

0.000 [107]

0.550 [107]

0.580 [107]




When temperature and pressure are not indicated, then the conditions are ambient (1 atm, 20 C). Source: Data from Refs. [20,34,36,51,52,61,71,73,78,106,107,112154].

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SECTION | II Green Solvents

describes the basics in phase equilibria for pure CO2 and binary mixtures of CO2 and organic solvents. Not only SCFs will be considered here, but also GXLs. When a system is in equilibrium, it does not present any variation with time or tendency to vary with time. The free energy of the system is zero. If more than one phase is in equilibrium, the system state is defined by one of the specified properties (composition, pressure or temperature). This is based on the Gibbs phase rule, which describes the degree of freedom (F), obtained by the number of phases (NP) and the number of compounds (NC) of the system. Applying Gibbs phase rule, the degree of freedom will indicate the number of independent variables or the number of thermodynamic parameters that are necessary, such as temperature and pressure. F 5 2 2 NP 1 NC

ð7:14Þ

For pure fluids, the number of compounds is equal to one (NC 5 1). If one phase is present in the system (NP 5 1), the degree of freedom is equal to two, which means that there are two independent variables, temperature, and pressure. If two phases are present in the system (NP 5 2), the degree of freedom is equal to one, which means that pressure and temperature are no longer independent variables. Check Fig. 7.1 for visualization of the phases for neat CO2. For more than one compound in the system (mixtures), the degree of freedom increases, because the composition also becomes an independent variable. Depending on the compounds in a closed system and the temperature and pressure, one or more phases can exist. If two phases are present in equilibrium, the bubble point is the temperature (at a given pressure) or the pressure (at a given temperature) where the first bubble (vapor drop) appears. The dew point is the temperature (at a given pressure) or the pressure (at a given temperature) where the fist dew (liquid drop) appears. The bubble point line and dew point line are visualized in Fig. 7.8 for a binary system containing CO2 and ethanol at 40
C, 60
C, and 80
C (plotted with data from [6370]). The figure also shows the region for SCF (shaded area above the CP) and GXL (two-phase GXL inside the VLE phase envelope and one-phase GXL in the liquid regime). Obviously, the higher the temperature of the mixture, the higher the pressure is needed to assure a one-phase liquid. This is an important fact to consider when using compressed fluid mixtures as solvents in separation processes.

7.4.2 PV and PT Phase Diagrams A common way to represent pure component or mixture states is by using a PV phase diagram. In such diagram, the variation of volume (V) with pressure (P) is given (Fig. 7.15). In the same diagram, it is also possible to visualize the roots of the cubic EOS in the model representation and visualize

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183

FIGURE 7.15 Pressure/volume (PV) phase diagram.

isotherms below and above the CP (T . Tc and T , Tc). It is important to highlight here that the examples refer to a closed system, where the concentrations of the compounds are constant. Using cubic EOS, for a given condition of pressure and temperature, it is possible to calculate three different molar volumes and compressibility factor for systems in the two phases (VLE). The small root (i.e., solutions of algebraic equations) of the cubic equation represents the liquid molar volume, while the bigger root represents the vapor molar volume. The intermediate root value does not have a physical meaning and it is not considered. The EOS curve in the two-phases region (in red color) happens when the EOS tries to model the phase boundary discontinuity. For the one-phase region, the calculated values from the EOS are the same as the experimental values (the line in red color would overwrite the black dashed line). In the PV diagram, an easier way to visualize and understand the changes in volume, pressure, and state of matter is to start in the vapor phase (right side of the PV diagram). Considering an isotherm below the CP (T , Tc), when a gas is compressed and its volume gets smaller, the variation in pressure due to the compressibility is small. If the compression continues, inside the two-phases region, the molar volume of the system will decrease because of the decrease in the vapor concentration. The vapor molar volume is higher than the liquid molar volume. The pressure inside the two-phases region will not vary until the end of the phase transition. If the volume continues to the decrease, one liquid phase will be formed and the pressure of the system will increase dramatically because of the incompressibility of the liquid.

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SECTION | II Green Solvents

FIGURE 7.16 Types of phase diagrams for high-pressure systems—Van Konynenburg and Scott classification for binary mixtures according to their changes near the CP. Reprinted from Van Konynenburg, P.H., Scott, R.L., 1980. Critical lines and phase equilibria in binary van der Waals mixtures, Philos. Trans. R. Soc. London Ser. A, 298, 495540 with permission from The Royal Society.

Increasing the temperature, the liquid molar volume increases and the vapor molar volume decreases (observe that the distance between the saturated liquid line and the saturated vapor line is smaller). At the top of the phase diagram, the three roots converge into one root, so the cubic EOS gives as result of the calculation one real root and two imaginary. One of the most used classifications of phase diagrams for binary mixtures at high pressure is the Van Konynenburg and Scott [155]. This classification is based on how binary mixtures change near the CP, and shows pressure versus temperature on the scales. According to Van Konynenburg and Scott [155], Types 1, 2, and 6 phase diagrams described below present CP values of the binary mixture continuous and between the CP of compound 1 and compound 2, as given in Fig. 7.16. Type 1 presents only one liquid phase in equilibrium with the vapor phase at any conditions of temperature and pressure, and it is exactly as described before; it contains a continuous critical line. It generally occurs in a mixture of two compounds with similar CPs. The other types of phase diagrams in this classification (Types 26) contain two liquid phases in equilibrium with the vapor phase (three phases in total), and at certain conditions of temperature and pressure, one of the phases in equilibrium will be miscible in the other phase and will be vanished, resulting in a system containing two phases from the intersection point.

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185

In Type 2, two liquid phases exist in equilibrium with vapor, but at certain conditions of temperature and pressure, it exists a point where only one liquid phase will be present in equilibrium with the vapor. This intersection point is called upper critical end point (UCEP), which is where the two liquid phases become one liquid phase with increasing temperature. In this way, there are two critical lines, one connecting the CP of compound 1 and compound 2, and another starting at the UCEP (where it has its maximum temperature) and continuing in the mixture at higher pressures. In Type 6, the continuous critical line is present along with the UCEP, however there is another intersection point in the phase diagram. In this intersection point called the lower critical end point (LCEP), the one liquid phase becomes two liquid phases with the increase in temperature. In this way, there are two critical lines, one connecting the CP of compound 1 and compound 2, and another starting at the LCEP and finishing at the UCEP (where it has its maximum temperature). Types 35 do not present a continuous critical line. In Type 3, there are two critical lines, one connecting the CP of compound 1 with the UCEP and another starting with the CP of compound 2 that continue at higher pressure and lower temperatures (Fig. 7.16). Type 4 presents three critical lines, one connecting the CP of compound 1 with the UCEP, one connecting the CP of compound 2 with the LCEP, and the last one starting at the UCEP (where it has its maximum temperature) and continuing in the mixture at higher pressures. Type 5, there are also two critical lines, one connecting the CP of compound 1 with the UCEP and the CP of compound 2 with the LCEP. Binary systems containing alcohols are highly nonideal because of the polarity and solvatochromic effects explained in Section 7.3.4 and the phase diagrams can be of Types 2, 3, or 4. However, in this chapter the main solvents of interest are the “green” ones, methanol, ethanol, and propan-1-ol, e.g. These solvents consist of small carbon chains (lower than 5) and mainly present phase diagrams of Types 1 and 2 [155157]. Mixtures containing CO2 and acetone present phase diagram Type 1 according to Roma´n-Ramı´rez et al. [156]. In this chapter, we will not include the systems and examples including azeotropes (positive or negative). Another important highlight is that cubic EOS can be used for Types 15, while other types of EOS are necessary for Type 6 (e.g., SAFT, Statistical Associating Fluid Theory, variations). VLE experimental data of systems containing CO2 and alcohols have been investigated by several authors [72,158,159]. Fig. 7.17AD shows VLE data for CO2 1 methanol [64,80,160,161], CO2 1 ethyl acetate [162165], CO2 1 acetone [166170], and CO2 1 ethyl lactate [171].

7.4.3 Thermodynamic Modeling of SCFs and GXLs Properties Fugacity (rather than only pressure, P) enables modeling of the behavior of real gases using thermodynamics relationships. EOS establishes a

186

SECTION | II Green Solvents

(A)

(B) FIGURES 7.17 VLE data for (A) CO2 1 methanol; (B) CO2 1 ethyl acetate; (C) CO2 1 acetone; and (D) CO2 1 ethyl lactate. ’ 25
C, K 40
C, ▲57
C, V 60
C, and ▬ 80
C. Plotted with data from Refs. [64,80,160171].

relationship between pressure, temperature, and volume (PVT). Since van der Waals EOS was described in 1873, many EOSs have been proposed in the literature. EOS is generally applied to systems at high pressure. On the contrary, for pure liquids a low pressure for which the volume change is low, activity coefficient models that are pressure dependent are mostly used (e.g., Excess Gibbs models). When thermodynamic models such as EOSs are used to correlate experimental data, the experimental values can be extrapolated inside the temperature or pressure interval considered in the parameter regression. For example, for a binary system containing CO2 plus an organic solvent at determined temperature and pressure, if experimental data are available for molar fraction of CO2 of 0.5 and 0.7, but the information of molar fraction of CO2 is not available at 0.6, this information can be obtained from an EOS after using the available experimental data to regress

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187

(C)

(D) FIGURES 7.17 (Continued).

the parameters. Moreover, predictive EOS available in the literature, which had the parameters regressed considering the available data in the literature are options to be used (carefully) to predict composition and/or temperature and/or pressure of systems in equilibrium. Since experimental work can be expensive and time consuming, EOS becomes an important tool for predictions of phase equilibrium and solubility of solutes. When both liquid and vapor phases are present at high pressure, it follows the isofugacity criterion: fiV 5 fiL

ð7:15Þ

where fiV is the fugacity of the compound in the vapor phase and fiL is the fugacity of the compound in the liquid phase. The fugacity of the compound in each phase is given by: fiV 5 yi φVi P

ð7:16Þ

188

SECTION | II Green Solvents

fiV 5 xi φVi P

ð7:17Þ

where yi and xi are the composition of the compound i in the vapor and in the liquid phases, respectively. φVi and φVi are the fugacity coefficients of the compound i in the vapor phase and in the liquid phases, respectively. P is the system pressure. The fugacity coefficients (φVi and φVi ) can be calculated using EOS, for example. When the fluid (e.g., scCO2) is in equilibrium with the solute (solid phase), it also follows the isofugacity criterion: f1S 5 f1F

ð7:18Þ

where fIS is the fugacity of the solute (compound 1) in the solid phase and f1F is the fugacity of the solute in the fluid phase. The fugacity of the solute in each phase is given by:   v1 ðP 2 PSat Sat 1 Þ f1S 5 φSat P exp ð7:19Þ 1 1 RT f1F 5 y1 φF1 P

ð7:20Þ

Sat where PSat 1 is the sublimation pressure, φ1 is the fugacity coefficient of the solute (compound 1) at the sublimation pressure, v1 is the molar volume of the solid at the system temperature (T) and φF1 is the fugacity coefficient of the solute (compound 1) at the fluid phase. The fugacity of the SCF is considered negligibly in the solid phase. One of the most used equations of state to calculate the fugacity coefficient for SCFs and GXLs is PengRobinson (PR-EOS) [29].

P5

RT aαðTÞ 2 ðv 2 bÞ v2 1 2bv 1 b2

ð7:21Þ

where P is the pressure, R is the gas constant, and v is the molar volume. For pure component, the variables a, b, and αðTÞ are written in function of the critical properties: a5

0:457235R2 Tc2 Pc

ð7:22Þ

b5

0:077796RTc Pc

ð7:23Þ

αðTÞ 5 ½11ð0:3746411:54226ω20:26992ω2 Þð12Tr0:5 Þ2

ð7:24Þ

where ω is the acentric factor and Tr is the relative temperature. The same equation can be written in terms of the compressibility factor: Z 3 2 ð1 2 BÞZ 2 1 ðA 2 2B 2 3B2 ÞZ 2 ðAB 2 B2 2 B3 Þ 5 0

ð7:25Þ

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where A5

aαðTÞP R2 T 2

ð7:26Þ

bP RT

ð7:27Þ

B5

The compressibility factor (Z) can be used to indicate the deviations of the thermodynamic properties from the ideal gas. For mixtures, the same equations can be used, however a mixing rule is needed. Different mixing rules can be used for PR-EOS and mixtures containing CO2. One of the most used mixing rules is the classical van der Waals with one binary interaction parameter (vdW-1) or two interaction parameters (vdW-2): a5

k X k X

xi xj aij

ð7:28Þ

i51 j51

b5

X

x i bi ;

for vdW-1

ð7:29Þ

i

b5

k X k X

xi xj bij ;

for vdW-2

ð7:30Þ

i51 j51

where aij 5 bij 5

pffiffiffiffiffiffiffiffi ai aj ð1 2 kij Þ

bi 1 bj ð1 2 lij Þ; 2

for vd-2

ð7:31Þ ð7:32Þ

Different EOSs were used for solid solubility correlation and prediction including GXLs. Laird et al. [172] have compared a methodology based on molecular simulation with PR-EOS in the representation of phase equilibrium data containing GXLs. Su et al. [173] have compared and improved the Predictive SoaveRedlichKwong equation of state (PSRK-EOS) for a selection of solvents used as GXLs. One of the challenges in using groupcontribution models to predict thermodynamic properties of mixtures containing GXLs is the lack of binary interaction parameters between the compounds.

7.5 SOLUBILITY OF COMPOUNDS IN SCFs AND GXLs Information about solubility is important in the extraction, separation and purification of compounds. Some examples of solubility in scCO2 in the literature are the work of Mendez-Santiago and Teja [174], which brings the information of solute solubility in scCO2 for different types of compounds,

190

SECTION | II Green Solvents

and the solubility of caffeine in scCO2 with and without the addition of a cosolvent by Kopcak and Mohamed [175]. The book by Gupta and Shim [2] has tabulated solubility data for a large number of compounds in scCO2 as found in the literature. The variation in solubility was analyzed by Aim and Fermeglia [176] for naphthalene in dense ethylene at two different temperatures (25
C and 45
C). The authors [176] showed that the increase in solubility of naphthalene in dense ethylene with temperature is significant at low pressures, and smaller for pressures over 20 MPa. Moreover, at low pressures, the solubility decreases with increasing temperature, while the opposite is true at higher pressure (i.e., the so-called crossover point). This can be explained by the larger negative change in density with increasing temperature at lower pressure, while at higher pressures density does not vary to the same extent with a change in temperature (Table 7.3). At higher pressure, the increase in temperature rather has a positive effect in terms of increasing vapor pressure. The solubility diagram of naphthalene in ethylene versus pressure shows that different temperatures present a point that the solubility curves cross, called the cross-over point. This point can be observed as a result of the competition of the solvent density and solute vapor pressure effects. The cross-over point can also be observed in enantiotropic systems, where different polymorphs are stable at different temperatures. Fig. 7.18 shows the cross-over point for the solubility of phenanthrene in neat scCO2 [2].

FIGURE 7.18 Crossover pressure demonstrated for the solubility (y, mole fraction) of phenanthrene in scCO2 versus pressure (bar), at 40
C, 50
C, and 60
C (313, 323, and 333 K). Reprinted from Gupta, R.B., Shim, J.-J., 2007. Solubility in supercritical carbon dioxide. CRC Press, Taylor & Francis Group, Boca Raton, FL with permission from CRC Press, Taylor & Francis Group.

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Shimizu and Abbott [177] have discussed how the addition of an entrainer (cosolvent) to scCO2 enhances the solubility. They consider three different hypotheses: (1) that there is a preferential entrainersolute interaction in comparison to CO2solute interaction; (2) that the bulk scCO2 becomes heterogeneous due to a self-aggregation of the entrainer molecules or by the aggregation of the CO2 molecules around the entrainer molecules; or (3) the density of the mixture increases, which increases the solubility. Using the KirkwoodBuff theory, the authors [177] showed that the first hypothesis (1) and part of the second hypothesis (entrainerentrainer interaction) are valid, while the increase in density by the entrainer and the CO2 clustering around entrainer do not explain the effect of all entrainers. Near the CP, the solubility of a solute at infinite dilutions changes significantly due to its partial volume variation and the large isothermal compressibility of the solvent [178]. Furthermore, if the scCO2 is mixed with an organic cosolvent, the solutesolvent interactions are higher than in homogeneous liquids, which causes local densities of scCO2 around the organic solvent molecules that is higher than the bulk density. For GXLs containing scCO2 and an organic solvent, depending on the polarity of the solute, solubility can increase or decrease with the scCO2 concentration (molar fraction). For example, if the solute of interest is polar, smaller concentrations of scCO2 (molar fraction) will result in an increase in the solubility. The same was observed for SCFs in terms of solubility for variation in pressure and temperature in the system is observed for GXLs. Solute solubility in SCFs and GXLs (binary mixtures) can be represented considering the following relationship: y1 5

  PSat v1 ðP 2 PSat 1 1 Þ exp RT φP1

ð7:33Þ

where P is the system pressure, PSat 1 is the saturated pressure, φ1 is the fugacity coefficient, and v1 is the molar volume for compound 1 (the solute/analyte). Also here, one of the most used equations of state to calculate the fugacity coefficient for SCFs and GXLs is PR-EOS. Solubility of solutes can also be correlated using density-based models. T¨urk et al. [179] have compared EOS with density-based models for pharmaceutical compounds. According to the authors [179], PR-EOS gives higher deviation than the empirical approaches, while Leonhard and Kraska EOS (LK-EOS) [180] (noncubic) gives similar deviation of the empirical approaches. As the authors [179] also highlight, the empirical approach only correlates density and solubility, while EOS present correlations in function of pressure. Sparks et al. [181] have also analyzed semi-empirical densitybased models for solute solubility in scCO2 and proposed a new model. Hansen solubility parameters (HSP) describe the solubility by the similarities between the solute and the solvent, such as polarity and hydrogen

192

SECTION | II Green Solvents

bonding, similar to what is described above using the KamletTaft solvatochromic method. Hansen improved Hildebrand’s work [182] considering the interaction between the compounds. The Hildebrand solubility parameter is based on the cohesive energy density or internal pressure of the liquid: sffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ΔHV 2 RT δ5 ð7:34Þ Vm where ΔHV is the enthalpy of vaporization and Vm is the molar volume. This means that the Hildebrand solubility parameter only takes into account the energy required to break the intermolecular interactions of the solvent molecules (in order to produce vapor from liquid), the space one mole of solvent molecule occupies and the temperature of the solvent. Lower molar volume is related to lower Gibbs free energy of mixing, which is the energy of mixing two substances, which in turn favors the solubility. For Hansen [183], the cohesive energy is based on the following three intermolecular interactions: 1. ED, dispersion interaction (nonpolar, induced dipoleinduced dipole interactions) 2. EP, polar interactions (dipoledipole and dipoleinduced dipole interactions) 3. EH, hydrogen bonding or other attractive interaction (e.g., Lewis acidbased) The solubility is related to the cohesive energy per volume, these three parameters then generate the total solubility parameter: E ED EP EH 5 1 1 V V V V

ð7:35Þ

δT 5 δD 1 δP 1 δH

ð7:36Þ

Experimental data were used in the correlation of the dispersion interaction parameter (δD) by Blanks and Prausnitz [184]. It is based on the energy of vaporization and molar volume as a function of temperature. The polar interaction (δP) is based on the dipole moment of the solvent molecule, DM: δP 5

37:4DM V 1=2

ð7:37Þ

Finally, the hydrogen bonding interaction and its parameter δH are obtained by subtracting the polar and dispersion energies of vaporization from the total energy of vaporization. Solubility parameters are important in chemical separation processes, where they can be used to compare different solvents in dissolving a specific solute. As an example, Srinivas et al. [185] have used HSP for the optimization of subcritical water extraction of bioactive compounds. Al-Hamimi

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193

et al. [186] have used HSP to better understand the interaction between the analyte and the solvent and compare the results with the analytes extraction. Seo et al. [187] used HSP for the supercritical antisolvent process. Li et al. [188] have compared theoretical solubility using HSP with experimental values in the extraction of volatile aroma compounds. Different authors [31,189196] have proposed a group contribution or EOS method to predict HSP. Li et al. [190] have showed different results obtained from available HSP models in the literature for scCO2. Nevertheless, three of the selected models [31,191,196] seem to agree between them. One of these models that is mostly used for scCO2 is the one by William et al. [191]. In this method, the correction of the HSP is as follows:
 δDref Vref 21:25 5 ð7:38Þ δD V
 δPref Vref 20:5 5 δP V "
 # δHref Vref 0:5 23 5 exp 21:32:10 ðTref 2 TÞ 2 ln δH V

ð7:39Þ

ð7:40Þ

where δD, δP, and δH are the HSP considering the temperature effect, δDref, δPref, and δHref are the HSP at the reference temperature (25
C), T is the new temperature, Tref is the reference temperature (25
C), V is the molar volume at the new temperature and pressure, and Vref is the molar volume at the reference temperature (25
C) and reference pressure (1 atm). The temperature dependence can be calculated by the Jayasri and Yaseen [197] method for liquids:

30:34 2 12 TTc ;i

5 δ 5 δref U4 ð7:41Þ T 12 Tref ;i c where δ is the HSP considering the temperature effect, δref is the HSP at the reference temperature (25
C), T is the new temperature, Tref is the reference temperature (25
C), and Tc is the critical temperature of the compound i. For a mixture of solvents, each of the HSP is linear with the composition of the solvents in the mixture: X δmixture 5 xi δT;i ð7:42Þ where xi is the composition of each of the compounds (CO2 and organic solvent, in percentage) and δi is the total HSP of each of the compounds i.

194

SECTION | II Green Solvents

7.6 EXPERIMENTAL PROCEDURES Some basic equipment is needed to conduct experiments with compressed fluids like SCFs and GXLs, as shown schematically in Fig. 7.19. First of all, a high-pressure pump is needed that can deliver compressed liquid CO2. Examples of pumps are single-piston syringe pumps, reciprocating dualpiston pumps, and pneumatic booster pumps. In SCF applications, compressed liquid CO2 chilled to around 4
C (to increase the viscosity as much as possible to enable easier pumping) is often mixed continuously with the organic cosolvent via a T-junction. The cosolvent is usually not chilled, and an ordinary HPLC pump can be used if a higher standard SCF pump is not available. It is possible but rare that premixed solvents are used in SCF applications. For GXLs it is more complicated. Regarding two-phase GXLs (inside the phase envelope), the solvent should be prepared in a batch vessel with stirring at constant temperature and pressure, see Fig. 7.19. From this highpressure vessel, only the liquid phase is pumped and used as the solvent in the process. In some cases, it may be needed to refill the vessel with more CO2 and/or organic solvent, and in this case, it is crucial not to disturb the temperature and pressure. It is less crucial to control the composition, as discussed earlier, as long as the liquid is in equilibrium with its vapor phase. For one-phase GXLs under conditions above the phase envelope, temperature, pressure and composition all need to be carefully controlled. In this case, premixed solvents are most likely preferred. For instance, in Susan Olesik’s group, GXLs are mixed inside single-piston syringe pumps and the mixture after equilibration is used as a mobile phase in chromatographic

FIGURE 7.19 Schematic drawing of equipment useful in separation processes using SCFs and GXLs as solvent.

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195

separations [198]. An alternative is to produce the GXL similarly as for SCF in continuous flow operation using a T-junction [186]. The limitation then is that the molar flow rate as well as the efficiency (kinetics) of the mixing of the two solvents must be precise and known at the mixing point. For a molar flow rate to be calculated, temperature and pressure along with the volumetric flow rate need to be known. This is unfortunately not usual for commercial equipment. In addition, the compressibility of the solvent may lead to systematic errors that are difficult to discover. Further about equipment, a treatment vessel is needed—usually made of stainless steel. There are many different stainless steel alloys available for vessels and tubings, and the lower the carbon content is the higher the corrosion resistance. The vessel itself can be of batch type with or without stirring, or in tubular format for continuous flow operations. About tubings, stainless steel is normally used, and depending on outer and inner diameters, the tubing is rated for different maximum pressure. Careful labeling of tubing in the lab is important for safe operations. Heating of the vessel is provided either by a heating jacket (tubular band heater), heating tape, cartridge heaters or by an oven. For continuous flow operation, a heat exchanger can be used to control the temperature of the CO2 or CO2/solvent mixture. Thermocouples of type K, J or T are used to monitor the temperature inside the treatment vessel as well as at the heating device (e.g., at the heating jacket). Temperature is then controlled using a proportional integral derivative (PID) controller, which uses a control loop feedback mechanism based on the data from the thermocouples. Pressure is monitored using pressure gauges inserted at appropriate points of the system. Pressure can also be controlled using a PID controller. Pressure of the system is maintained either in batch operations by closing the outlet tubing by a needle valve or ball valve, or in continuous flow operations by attaching a pressure restrictor, such as a narrow capillary, a nozzle, or a (spring-loaded) backpressure regulator (BPR). Pressure restrictors are usually heated to prevent freezing or clogging of precipitated solutes. Over-pressurization of the system is avoided by inserting burst disks (rupture disks) at appropriate highpressure locations of the system. It is important to make sure that correct disks are used based on the part of the system that can withstand the lowest pressure. The disk should burst at a pressure just above the maximum allowed for a specific system. Some experimental procedures regarding the study of SCFs and GXLs fundamentals as well as solubility studies are described here. In order to allow the study of phase equilibria, the key equipment to be used is a viewcell that can withstand high pressures. This is normally based on a stainless steel vessel of either fixed or variable volume, and it has one or more transparent sapphire windows mounted to allow for visual inspection. Such equipment can be used to study phase transfer and phase composition at different pressure, temperature, and composition. An example of a view-cell is shown

196

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FIGURE 7.20 Variable-volume view-cell. In one end, a sapphire window is mounted (righthand side), and on the left-hand side a piston is moved by applying external pressure of, for instance, compressed CO2. Thermocouples (not shown in the figure) are inserted to measure the temperature of both inside the vessel and at the heating point supplied by the heating jacket. Pressure is monitored using a pressure gauge mounted to the the vessel, and is adjusted by applying more or less external pressure on the moving piston. A burst disk is mounted on the same line as the pressure gauge. Several ports are available to access the inside of the vessel, as controlled by manual needle valves. Connected to the view-cell is one or two high-pressure syringe pumps to deliver compressed liquid CO2 as well as organic solvent (not shown in the figure). If a solute is to be used in a study, this is placed inside the vessel before closing the side with the sapphire window.

in Fig. 7.20. The photos in Fig. 7.2 are taken through the sapphire window of this kind of view-cell. One of the most common method for phase equilibria studies is the cloud point pressure method [199]. In this method, the solvent is added to the view-cell and the temperature is set at the heater (e.g., using a heating jacket or an oven). After the equilibration of the temperature, pressure is increased by decreasing the volume of the view-cell by moving the piston. This is done until the mixture inside of the view-cell (visualization through the sapphire window) is cloudy. At this point, only one phase is seen in the system (liquid or SCF, depending on the conditions of pressure and temperature) instead of the previous two phases (liquid and vapor). For some of the studied mixtures, a better visualization of the cloud point is enabled when starting with a higher pressure and then lowering the pressure of the system toward the phase transfer. For solubility measurements in SCFs or GXLs, the most common methods used can be classified as static [69], recirculation [200,201], and flow [202,203] methods. In the static method, after the temperature and pressure has equilibrated in a vessel containing the mixture, the composition of the single or more phases is analyzed. The analysis can be visual (through a

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sapphire window) or by sampling (by opening a valve and collect/trap a portion of the solvent containing dissolved solute). In the recirculation methods, the solvent inside the vessel is driven in an external loop using a recirculating pump until equilibrium is reached at determined temperature and pressure. A sample can be taken using a multiport valve or an injector loop (for online analysis). Alternatively, the concentration of the solute can be determined in situ, using an in-line detector, such as a spectrophotometric charged couple device detector. In the flow methods, the solvent/solute mixture is prepared and equilibrated at determined temperature and pressure, or alternatively the pure components are heated separately and mixed in a T-junction. After that, the mixture is pumped with controlled pressure. If more than one phase is present in the mixture in equilibrium, a vessel is used before the pump. When the equilibrium is reached at certain conditions of pressure and temperature, the light and the heavy phases flow from different lines using two different pumps. The flow of the solvent/solute mixture is afterward collected and the composition is analyzed. For any of the experimental setups described above, physicochemical properties of SCFs and GXLs can be determined. For instance, density can be measured using a vessel connected to a vibrating-tube densitometer [85,200,204208]. Other options for the density measurements are found in the literature, such as a laser okhim fiber optic densitometer [209], singlesinker densimeter [210], gamma densitometer [211], among others. Regarding the viscosity measurements, one of the most common equipment used is the falling-weight viscometer [49,51,52,212,213]. A vibratingtube was also used by different authors [207,208,210]. In many of the publications, the falling-weight or the vibrating-tube is designed by the research group. Dielectric properties can be experimentally determined using the KamletTaft solvatochromic method as described earlier. In practice, the same apparatus used for solubility measurements can be used to which a spectrophotometer (UV/Vis) is attached online or in-line. Ngo et al. [214] have used different spectroscopic techniques including UV/Vis for the measurement of the solubility of anthracene, 1,4-naphthoquinone, and 2-naphtol in scCO2/methanol mixtures (concentration of methanol between 0 and 2.37 mol%). Another example is the solubility measurement of racemic paroxetine intermediate in scCO2 by Bao et al. [215].

7.7 EXAMPLES OF APPLICATIONS Applications will only be briefly covered here, since other chapters describe separation techniques based on SCFs and to some extent also GXLs. First of all, SCFs have been extensively used for extraction of natural and pharmaceutical compounds. One of the benefits of using an SCF as extraction solvent is that the extracted product/analytes can be easily separated from the

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solvent by depressurization. Well-known applications of using liquid CO2 in extraction is the decaffeination of coffee beans and the dry-cleaning of textiles. More recently, SFE has been studied by King [216] at pressures of up to 70 MPa regarding different analytes/solutes such as lipids, pigments, and polymers with focus on the solubility. The same author has also demonstrated the use of SCFs in food applications [217] with focus on the solubility, extraction and fractionation. A review of the solubility of solutes in SCFs has been done by Skerget et al. [218]. The authors discuss different applications of using SCFs as a solvent, such as in drying, cleaning and reactions, among others. In terms of using GXLs as an extraction solvent, the number of publications is few [22,186,219,220]. For example, CO2-expanded ethanol (onephase GXL) was used in the extraction of α-pinene and cis-verbenol by AlHamimi et al. [186] with focus on the solubility and extraction kinetics. CO2-expanded ethanol was compared with SFE and solidliquid extraction, and extraction using CO2-expanded ethanol as a solvent showed faster extraction kinetics of the analytes. In Paudel et al. [22], CO2-expanded methanol (two-phase GXL) was used to extract lipids from microalgae. In chromatography, SCFs or GXLs are used as a mobile phase, and the technique is most commonly called SFC. This chromatographic technique is described in more detail in Chapter 16, but some general remarks and examples will be given. Only for the most hydrophobic (e.g., nonpolar lipids) SFC is being operated with an SCF such as scCO2 mixed with, for instance, a small volume fraction of methanol as a mobile phase. It is more common to increase the content of organic cosolvent so that the solvent is outside the regime of SCF, entering the liquid phase (or one-phase GXL as we call it here), see Fig. 7.8. In old literature, it was more common to enable gradient separation in SFC by changing temperature and pressure, while it nowadays is more common to simply add more and more of an organic cosolvent, just like it is common in HPLC. Examples of the modern literature are the separation of lipid classes and lipid species, e.g., in lipidomics studies [221], separation of vitamin D metabolites in blood plasma [92], carotenoids in microalgae [10] and in serum [222], and for the separation of pharmaceutical compounds [223]. Again, the benefits of fast mass transfer are obvious in chromatographic separations, along with the fact that scCO2 is a nonpolar molecular solvent that enables both normal-phase and reversed-phase separation mechanisms with a larger possibility to use stationary phases with different functional groups to enhance the selectivity of the separation. Attention has to be given to the fact that compressibility of the solvent varies a lot depending on the CO2 content, pressure, and temperature, as described above. In particle formation and processing of polymers, SCFs and GXLs have been used as solvents or antisolvents to promote crystallization/precipitation. For instance, Ventosa et al. [224] presented a process called depressurization of an expanded liquid organic solution (DELOS). It consists in first

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solubilize the solute in the organic solvent (such as acetone), add the SCF (e.g., scCO2) in order to obtain the GXL, and then start the homogeneous depressurization of the system to obtain microparticles with narrow particle sizes distribution. Finally, although a bit outside the scope of this chapter, SCFs and GXLs have been used as solvent in heterogeneous and homogeneous catalysis. For instance, Subramaniam [225,226] have described the benefits of using GXLs in reactions present in catalytic processes. The advantages of the use of such solvents include the waste minimization, reduced use of volatile organic solvents, process intensification, and efficient feedstock utilization. Akien and Poliakoff [227] bring a list of reactions including GXL with and without CO2 as a fluid. The authors [226] discuss the use of SCF and GXL as potential options to replace fossil-based raw materials in industry.

7.8 GREENNESS OF SCFs AND GXLs In this chapter, fundamentals of SCFs and GXLs have been described and compared to conventional liquid organic solvents. Table 7.5 shows a list of the most common conventional organic solvents used in separation processes, along with a selection of SCFs and “green” GXLs at certain temperature, pressure, and composition. Some of these solvents’ KamletTaft solvatochromic parameters are plotted in Fig. 7.14. In order to discuss “greenness” of a solvent, the functionality of a solvent should be considered, i.e., to dissolve solutes/analytes. In addition, the entire life cycle of a solvent should be considered [228]; from production to usage and disposal/recycling as well as transportation. Some of the aspects of greenness of solvents are as discussed by Refs. [229231] and from these references, a green solvent is one that demonstrate one or several of the following: G G G G

G

G G

Low toxicity to humans and other organisms Easily biodegradable in the environment without adverse effects Is naturally occurring, i.e., does not have to be produced Produced from renewable sources, i.e., does not contribute to our oil dependency Produced as a byproduct, i.e., shares the environmental impact with other products Has a low vapor pressure, i.e., is not readily released to air A traditional evaporation step can be avoided, i.e., a solvent change is energy efficient.

Carbon dioxide is suggested as being a green solvent. CO2 has limited toxicity, and is not added to the environment as a new product. It is rather “borrowed” from the environment and then returned. In fact, when CO2 is

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used as a solvent, it is not contributing as a green house gas during its usage, i.e., it has a positive environmental effect. CO2 is naturally occurring, and it is also produced as a byproduct from chemical processes as mentioned in Section 7.1. It does not however have a low vapor pressure, hence gas-tight systems should be used with appropriate ventilation and recycling systems. Evaporation of CO2 is easy, since it is naturally occurring as a gas. Hence, an energy-demanding evaporation step can be avoided. In summary, scCO2 is a green solvent considering the aspects above. However, high-pressure equipment such as pumps and heat-controllers do require electricity. Hence, depending on the application, how high pressure is needed and whether the selected SCF solvent mixture is a good choice in terms of solubility of the target compounds in a certain application, will determine whether the scCO2 is the greenest as compared to other green solvent options. Similarly, how green a GXL is depends on which organic solvent is selected, and how good of a solvent the GXL is in a certain application. The higher the solubility, the less amount of solvent is needed. There are several articles describing definitions of green solvents and suggested lists of solvents that could be considered green. For instance, if environment, health, and safety (EHS) is used as a “green” indicator, methyl acetate, ethanol, methanol, ethyl acetate, propanol, butyl acetate, and butan-1-ol are among the greenest organic solvent options [228]. What is considered as part of EHS is environmental effects (water hazards, air hazards, persistency), health effects (chronic toxicity, irritation, acute toxicity), and safety aspects (reactivity/decomposition, fire/explosive, release potential). If instead a life cycle assessment (LCA) is performed for organic solvents, focusing on the cumulative energy demand (CED) during production, transportation, and recycling/incineration of the solvent, then hexane, heptane, diethyl ether, pentane, methanol, and ethanol appear to be green solvents. This is due to the fact that some organic solvents like straightchain alkanes are easily produced directly without any synthesis steps (just distillation) from the oil refinery. As discussed in Ref. [228], combining the EHS and LCA results, clearly methanol, ethanol, and ethyl acetate are very good options of organic solvents. Moreover, in a GXL the organic solvent is “diluted” with compressed liquid CO2, meaning that less amount of organic solvent is needed. How large this effect is, how much less organic solvent that is used because of adding CO2, should be investigated—preferably by conducting comparative LCA studies. It should be taken into consideration the electricity used for compressing the fluid. GXL require pressurization, although not as high pressure as is used for producing an SCF. In conclusion, SCFs and GXLs based on the use of compressed liquid CO2 mixed with a “green” organic solvent giving tunable solubility as well as fast mass transfer as compared to conventional solvents are potentially interesting green solvents that should be further explored in separation science.

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7.9 CONCLUSIONS With this chapter we have gathered information about fundamental physicochemical properties of SCFs and GXLs, including both solubility (dielectric) properties and mass transfer properties. Clearly, scCO2 shows potential as a green solvent not only neat in terms of dissolving hydrophobic compounds, but also in mixtures with green organic solvents either as an SCF or as a GXL for dissolution of more polar compounds. In general, information about GXLs is scattered in the literature, partly because there is no consensus around the definition of a GXL. At pressures, temperatures, and compositions above the binary phase envelope, it is unclear what the binary fluid mixture is called. As a consequence, the literature shows a range of different terminologies, which complicates the search for reference data. It is also evident that physicochemical property data are missing for many GXLs. We hope that this chapter is a good starting point for both the novice user as well as for the more experienced SCF user who would like to discover the field of GXL in separation science. The ambition is also that students in natural science subjects without an engineering background may find this chapter useful.

LIST OF ABBREVIATIONS CP EOS GXL HPLC HSP LCEP PR-EOS SAFT scCO2 SCF SFC SFE UCEP VLE

critical point equation of state gas-expanded liquid High-performance liquid chromatography Hansen solubility parameter lower critical end point PengRobinson equation of state statistical associating fluid theory supercritical carbon dioxide supercritical fluid supercritical fluid chromatography supercritical fluid extraction upper critical end point vapor liquid equilibrium

ACKNOWLEDGMENTS The Swedish Research Council (VR, 621-2014-4052, 622-2010-333) and the Swedish Research Council Formas (239-2013-971) are acknowledged for financial support. Victor Abrahamsson, Irene Rodriguez-Meizoso, and Margareta Sandahl are thanked for their suggestions on how to improve this chapter.

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