Electrochemical Aspects of
6442
A n a ly t i c a l C h e m i s t r y / s e p t e m b e r 1 , 2 0 0 7
© 2007 American Chemical Societ y
Ionic-Liquid|Water Two-Phase Systems
Takashi Kakiuchi Kyoto University
An electrochemical perspective is crucial for understanding such two-phase systems.
R
oom-temperature ionic liquids (ILs) that are salts in the liquid state have attracted much attention in a variety of areas because of their unique features, such as extremely low volatility, high polarity, and reasonable electrical conductivity (1–11). An IL with sufficient hydrophobicity forms a two-phase system upon contact with an aqueous phase (W). Among IL applications, the use of IL|W two-phase systems is one of the most promising for extraction and electroanalytical chemistry. An IL can be used as a substitute for a conventional organic solvent because its properties are similar to (but somewhat different from) those of polar molecular solvents (12, 13). However, IL|W systems have one distinct feature—when the liquids are brought into contact, ions in the IL start to dissolve in W, and a potential difference develops across the interface between the two phases. The origin of this phase-boundary potential is the difference between the affinity that the anions in the IL have for W and the affinity that the cations in the IL have for W. This is where the fundamental difference between ILs and molecular solvents emerges (14, 15). The phase-boundary potential dramatically influences the partitioning of charged components between the IL and W phases. The electrical double layers formed at the IL|W interface should affect the surface properties of the systems used in chromatography, electrophoresis, and potentiometric ion sensing. For a better understanding of the properties of IL|W two-phase systems, an electrochemical viewpoint is useful and important. In this article, some examples illustrating the importance of the phase-boundary potential are presented together with basic electrochemical concepts required in dealing with IL|W two-phase systems. The mutual solubility of the IL and W vary considerably, depending on the ions that make up the IL. Most ILs used for extraction studies have relatively high solubility in W, on the order of a few tens of millimolar (items 1–3 in Table 1; 16, 19). Strongly hydrophobic ILs (items 8–10) are 3 orders of magnitude less soluble in water; moderately soluble ones (items 4–6) are in between, at a few millimolar. Although octanol–water partition coefficients
Tony Fernandez
Partition equilibria
s e p t e m b e r 1 , 2 0 0 7 / A n a ly t i c a l C h e m i s t r y
6443
d SIL/W (M)a SW/IL DWIL´ DCsClb DSrCl2b (g/cm3) (wt %) (V)
Item IL 1
[C2mim+][C1C1N−]
1.50
4.5 × 10−2
1.9b
0.09
589
10,700
2
[C2mim+][C2C2N−]
1.53
8.9 × 10−3
1.1b
0.05
1140
79,700
3
[C4mim+][C1C1N−]
1.4
1.9 × 10−2
1.4b
0.02
380
935
4
[C4mim+][C2C2N−]
1.5
3.4 × 10−3
5
[C8mim+][C1C1N−]
6
[C8mim+][C2C2N−]
7
[THxA+][BEHSS−]
0.968
8
[THxA+][C1C1N−]
9
[C18Iq+][TFPB−]
10
[TOMA+][TFPB−]
0.7b
0.05
567
3950
1.32
1.8 ×
10−3
0.9
–0.04
8.35
3.95
1.38
1.1 × 10−3
0.6
–0.01
25.7
6.89
10−4
4.5
–0.19
—
—
1.186
7.2 × 10−6
0.3
–0.17
—
1.302
4 × 10−5
0.31c
—
1.228
10−5
0.20c
—
2×
1×
+ GAIL–
�
W,0
W,0
RT
�
Na+ 0.354
Ph4B – 0.372
H+ 0.337 Ba2+ 0.320
0.3
NH4+ 0.277
K+ 0.242
0.2 C4F9BF3– 0.140
—
—
Choline+ 0.117
—
—
C1C1N– 0.133 C3F7BF3– 0.102
(1)
�
�
�
A n a ly t i c a l C h e m i s t r y / s e p t e m b e r 1 , 2 0 0 7
Mg 0.361
Cs+ 0.159
It might seem impracticable to have a universal scale of the hydrophobicity of ions for predicting the properties of ILs on the basis of GIL→W,0tr,i. Fortunately, some measures of polarity, such as ET(30) (a solvent polarity scale based on the solvatochromism of a betaine dye; 23), show that ILs of different types have surprisingly similar polarities that are comparable to those of lower aliphatic alcohols (23, 24). It would therefore be useful to have a general measure of the ionic hydrophobicity or hydrophilicity. The most appropriate quantity for this purpose is G Org→W,0tr,i, which is the standard 6444
0.4
2+
—
where K Wsp is the solubility product of [C +][A−]. The solubilW W 0 0 + � � + – IL→W,0 IL→W,0 C IL IL A W = ity, and hence the tr,C+ and G tr,A−, is � sum of G IL 2 energies cannot be evaluated measurable, but the latter Gibbs separately thermodynamically, because GIL→W,0tr,i is a single ionic property (22). Moreover, this value varies from one IL to another because each individual IL provides its own solvation environment for i (14, 15).
Hydrophobicity scale
Dipicrylaminate – 0 .407
C2C2N – 0.199
are a useful measure of hydrophobicity (20, 21), a constitutive approach can be taken instead, to separately evaluate the hydrophobicity of the anions and the cations that make up the IL. A quantitative measure of the solubility of an ionic species i is the standard Gibbs energy of transfer from the IL to W, GIL→W,0tr,i. With this quantity, the solubility product in water of an IL composed of the cation C + and the anion A− can be expressed as (15) G CIL+
Li + 0.395
Rb+ 0.201
d, density; SIL/W and S W/IL, solubility of IL in W and W in IL, respectively; WIL, the phaseboundary potential deduced from Figure 3 and Equation 2; DCsCl and DSrCl2, distribution ratios of CsCl and SrCl2; C 2mim +, 1-ethyl-3-methylimidazolium; C 4mim +, 1-butyl-3-methyl imidazolium; C 8mim +, 1-methyl-3-octylimidazolium; THxA +, tetrahexylammonium; C18Iq +, N -octadecylisoquinolinium; TOMA +, trioctylmethylammonium; C1C1N −, bis(trifluoromethyl sulfonyl)imide; C 2C 2N −, bis(pentafluoroethylsulfonyl)imide; BEHSS −, bis(2-ethylhexyl)sulfosuccinate; TFPB−, tetrakis[3,5-bis(trifluoromethyl)phenyl]borate. a In this column, items 1–4 at 21 °C, Ref. 16; items 5–10, Ref. 17. b 21 °C, Ref. 16. c 56 °C, Ref. 18.
W ln K sp =
Sr2+ 0.342
Table 1. Physicochemical properties of some hydrophobic ILs at 25 °C, except where otherwise noted.
Ca2+ 0.349
0.5
0.1
BEHSS– 0.095 Picrate– 0.069
AcCh+ 0.049
C2F5BF3– 0.064
TMA+ 0.035
PF6– 0.025
C2mim+ – 0.03 TEA+ – 0.055 BzTMA+ – 0.086 C4mim+ – 0.096
0 Dodecylsulfate– –0.043 IO4– – 0.072
– 0.1
ClO4– – 0.082
BF4– – 0.114 +
C6mim – 0.154
SCN– – 0.164
+
TPrA – 0.170
– 0.2
I – – 0.191
C8mim+ – 0.220 TBA+ – 0.275
C10mim+ – 0.303
NO3– – 0.261 ClO3– – 0.270
– 0.3
Br – – 0.288
Ph4As+ – 0.372 C12 mim+ – 0.375 TPnA+ – 0.408 CV+ – 0.410
Cl – – 0.396
– 0.4
THxA+ – 0.472
– 0.5 FIGURE 1. W NB0 i (V) for the NB|W two-phase system at 25 °C. AcCh +, acetylcholine; TMA +, tetramethylammonium; TEA +, tetraethylammonium; TPrA +, tetrapropylammonium; TBA +, tetrabutylammonium; TPnA +, tetrapentylammonium; THxA +, tetrahexylammonium; BzTMA +, benzyltrimethylammonium; Ph4 As +, tetraphenylarsonium; CV+, crystal violet; Ph4B −, tetraphenylborate; for others, see Table 1. (Adapted with permission from Ref. 64.)
Gibbs energy of transfer of an ion to W from a polar organic (Org) solvent that is immiscible with W—for example, nitrobenzene (NB) and 1,2-dichloroethane (DCE). Figure 1 shows the experimental values of the standard ion-transfer potentials for various ions between NB and W, W NB0i. This quantity has the dimension of electric potential and is defined as −G NB→W,0tr,i/(z iF), where z i is the charge on i with the signed unit of electronic charge and F is the Faraday constant. The rescaling of G NB→W,0tr,i by z iF is convenient and important in the present electrochemical context, as will be demonstrated shortly. Like GIL→W,0tr,i, G NB→W,0tr,i, and hence W NB0i, is not measurable thermodynamically. The values in Figure 1 are deduced on the basis of the tetraphen ylarsonium–tetraphenylborate (TATB) assumption, which is believed to be the most reliable extrathermodynamic assumption of this kind (25). The assumption is that the Gibbs energy of transfer, enthalpy, and entropy of TA are equal to those of TB (26). Table 2 compares the experimentally observed ratios of the solubility values of a pair of ILs that have a common cation or anion (Table 1, column 4) with those calculated by using the W NB0i values in Figure 1 and Equation 1. Reasonable agreements between the values in the second and third columns in Table 2 (within a factor of ~2) demonstrate the usefulness of the W NB0i values in Figure 1 and the constitutive approach for estimating W IL0i.
Phase-boundary potential across the interface between IL and W
When the absolute values of W IL0i [= −GIL→W,0i /(z iF)] for the IL-constituent cation and anion are different, the difference in the solubility gives rise to the distribution potential across the interface between the IL and W to compensate for the difference in solubility of the cation and the anion, as described earlier. The phase-boundary potential W IL is the difference between the inner potentials of the IL phase IL and the W phase W—that is,IL WW,0 W IL . This is not a IL − W,0 IL +=G G – + A C Wmeasurable quantity either. For [C +][A−], thermodynamically ln K sp = RT W IL is given by (15)
�
� C0+ +
�
�
W IL
�
�=
�
W IL
W IL
� A0 –
(2)
2
With the values of W NB0i in Figure 1, W IL for IL|W twophase systems can be approximated (Table 1). The midpoint between the chosen cation and anion in Figure 1 gives an estimate of W IL. Although these values have substantial uncertainties, semiquantitative knowledge of the sign and the magnitude of W IL can be helpful if treated with care. W IL is used to denote a value estimated from W NB0i. For example, ILs 1–4 can be predicted to have positive W IL values
