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Ionic liquid salt bridge based on N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide for low ionic strength aqueous solutions
Journal of Electroanalytical Chemistry 651 (2011) 61–66
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Ionic liquid salt bridge based on N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide for low ionic strength aqueous solutions Yousuke Fujino, Takashi Kakiuchi ⇑ Department of Energy and Hydrocarbon Chemistry, Graduate School of Engineering, Kyoto University, Kyoto 615-8510, Japan
a r t i c l e
i n f o
Article history: Received 29 July 2010 Received in revised form 17 September 2010 Accepted 30 October 2010 Available online 18 November 2010 Keywords: Ionic liquids Ionic liquid salt bridge Low ionic strength samples Liquid junction potential Distribution potential N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide
a b s t r a c t The phase-boundary potential between the moderately hydrophobic ionic liquid and a low ionic strength aqueous solution is demonstrated to be stable and constant with the standard deviation of 0.4 mV down to 20 lmol kg1 HBr, LiBr, and KBr solutions, for three ionic liquids that consist of either N-methyl-Noctylpyrrolidinium, N-heptyl-N-methylpyrrolidinium, or N-hexyl-N-methylpyrrolidinium and a common anion species, bis(pentafluoroethanesulfonyl)amide. This stability is promising for accurate measurements of pH of low ionic strength samples and reliable estimates of single ion activities in general. The phase-boundary potential deviates from the value determined by the partition of the ionic liquid in further dilute aqueous solutions. The magnitude of the deviation ranges from 3 to 11 mV at 5 lmol kg1 MBr (M is H+, Li+, or K+). The solubility of these ionic liquids in water is 0.2 mmol dm3 at most at 25 °C, which is another advantage of ionic liquid salt bridge in electroanalytical chemistry. Ó 2010 Elsevier B.V. All rights reserved.
1. Introduction An ionic liquid (IL) that is moderately hydrophobic and immiscible with water can work as a salt bridge having several advantages over traditional KCl-type salt bridges [1,2]. A notable property of ionic liquid salt bridge (ILSB) is their stability of the phaseboundary potential when it is in contact with a low ionic strength aqueous solution. In the case of 1-methyl-3-octylimidazolium bis(trifluoromethanesulfonyl)amide (C8mimC1C1N) whose solubility in water is 1.6 mmol dm3 at 25 °C [3], the phase-boundary potential stays constant within ±1 mV in a 100 lmol dm3 aqueous KCl solution [4]. However, when the ionic strength of a sample solution is lower than this concentration range, the deviation from the constant value becomes greater and amounts to 20 mV at 20 lmol dm3 KCl [4]. This deviation was ascribed, mainly, to the diffusion potential formed in the aqueous phase accompanied with the dissolution of C8mimC1C1N, where the mobility of C8mim+ is less than that of C1C1N. Although the ionic strength on the order of 100 lmol dm3 may be seen dilute in typical potentiometric measurements [5,6], the ionic strength of actual samples can be less than this level. For example, rainwater in nonpolluted regions usually has the ionic strength of 50 lmol dm3 or lower [7–9].
⇑ Corresponding author. Tel.: +81 75 383 2489; fax: +81 75 383 2490. E-mail address: [email protected] (T. Kakiuchi). 1572-6657/$ - see front matter Ó 2010 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2010.10.028
In such low ionic strength solutions, the liquid junction potential between a KCl-type salt bridge and a sample solution is the main source of inaccuracy in determining pH [9–12], because the potential is not stable and can amount to more than 10 mV. One more serious problem is the elevation of the ionic strength of sample solutions due to the dissolution of KCl, which changes the activity of analyte ions, for example, H+. Given the stability of ILSB based on C8mimC1C1N in relatively low ionic strength solutions [2,4,13], well-designed ILSBs that have smaller solubility than C8mimC1C1N and exhibit a negligibly small diffusion potential in low ionic strength samples are promising in accurate electroanalytical chemistry, in particular, in the potentiometric determination of pH of low ionic strength samples. For such samples, the lower solubility of ILSB-constituent ions in W is preferable, because the contribution of the diffusion potential due to the dissolution becomes smaller [4], and the degree of sample contamination is less. Although the lower solubility means the higher electrochemical polarizability of the ILSBjW interface [14], the concentration level of interfering ions, if any, would also be small in such low ionic strength samples. Toward this goal, we have studied several moderately hydrophobic ILs. We herein report the use of N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide, where the alkyl group is either octyl-, heptyl-, or hexyl moiety, for ILSB. Aside from the solubility, it is important in choosing ions for ILSB to balance the hydrophobicity of the IL-constituting cation and that of the anion, so that
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the expected phase-boundary potential is close to 0 [15]. Judging from the standard ion transfer potentials of N-methyl-Noctylpyrrolidinium and bis(pentafluoroethanesulfonyl)amide in the nitrobenzene–water two-phase system, 0.23 and 0.2 V [16], respectively, it is expected, first, that the phase-boundary potential between the ionic liquid and water (W) would be close to zero [1]. Second, the solubility of this ionic liquid in water, whose measure is the difference in the standard ion transfer potentials of IL-constituting cation and anion [1], should be smaller than that of C8mimC1C1N, where the constituent cation and anion have the standard ion transfer potentials of 0.24 and 0.13 V in the same two-phase system, respectively [16]. We will show that the ionic liquids that consist of N-alkyl-N-methylpyrrolidinium and bis(pentafluoroethanesulfonyl)amide exhibit a stable liquid junction potential down to 20 lmol kg1 solutions of HBr, LiBr, and KBr and are hence suitable to a salt bridge for pH measurements of low ionic strength samples. 2. Experimental 2.1. Materials N-methylpyrrolidine (Tokyo Chem. Ind., TCI), 1-bromooctane (TCI), 1-bromoheptane (Aldrich), 1-bromohexane (TCI), N-methylN-octylpyrrolidinium chloride (Merck), and hydrogen bis(pentafluoroethanesufonyl)amide (Central Glass, 70% aqueous solution) were used without further purification. Other chemicals used were of reagent grade. N-alkyl-N-methylpyrrolidinum bromides were prepared from methylpyrrolidine and 1-alkylbromide by stirring the mixture of 1.05 (methylpyrrolidine):1.00 (1-alkylbromide) at about 80 °C for at least 6 h. The mixture was washed three times with ethylacetate and then volatile components were removed by drying the mixture first with an evaporator and then with a vacuum pump. N-alkyl-N-methylpyrrolidinium bromides thus obtained were identified with 1H NMR. Three salts, N-methylN-octylpyrrolidinium bis(pentafluoroethanesufonyl)amide, Nheptyl-N-methylpyrrolidinium bis(pentafluoroethanesufonyl)amide, and N-hexyl-N-methylpyrrolidinium bis(pentafluoroethanesufonyl)amide, were prepared by metathesis reactions between Nalkyl-N-methylpyrrolidinium bromides and bis(pentafluoroethanesufonyl)amide acid. These salts will hereafter be abbreviated as ½C18 pyrrþ ½C2 C2 N ; ½C17 pyrrþ ½C2 C2 N , and ½C16 pyrrþ ½C2 C 2 N , respectively. These salts were washed repeatedly with copious water until no precipitation of AgBr was detected when a few drops of a AgNO3 solution were added in washing water.
sample solution is non-negligible when the initial concentration, as prepared, of chloride in a sample solution is comparable to the solubility of AgCl. Cell (I) was constructed in a glass tube with upper and lower compartments separated by a disk of glass frit, as described elsewhere [2]. The cell was housed in a box for light shielding. The temperature of the cell was maintained at 25 ± 0.5 °C by circulating water through the jacket of the glass cell [2]. The cell voltage, i.e., the potential of the right-hand-side terminal with respect to the that of the left, E, was measured with an electrometer (ADC, ADCMT-8252), which was connected to a computer through GPIB, as described elsewhere [2]. The sampling rate was typically 0.2 Hz. Ag/AgBr electrodes were prepared by anodization of silver wires of 0.5 mm diameter in an aqueous solution of ca. 80 mmol dm3 KBr at 10 mA for 15 min. Before the AgBr coating, the silver wire was first polished with an emery paper and then treated by ultrasonic cleaning successively with Milli-Q water, 2% aqueous ammonia, 2% nitric acid, and Milli-Q water, for 30 min in each step. The electrodes were kept in Milli-Q water in the dark. þ The viscosity of water-saturated C1;n pyrrl ½C2 C2 N was measured with a viscometer (Toki Sangyo, TVE-33, cone-plate type) at 1, 2, 5, 10, 20, and 50 rpm at 25 ± 0.5 °C by circulating water in the cell. No significant dependence of the measured viscosity on the share rate was observed. The density of ½C1;n pyrrþ ½C2 C2 N was measured with an Ostwald pycnometer. The electrical conductivity of water-saturated ½C1;n pyrrþ ½C2 C2 N was measured with a conductivity meter (Kyoto Electronics, CM-117). The electrical conductivity of electrolyte solutions to evaluate the molar conductivity of ions in water was measured with a bridge-type conductivity meter (Husou, HECS362D with HECS363D) and a dip-in-type cell (DKK-TOA, CT-5710B) at four concentrations of C1,npyrrolidinium bromides and bis(pentafluoroethanesufonyl)amide acid in a water bath at 25 ± 0.02 °C. The cell constant was determined using a 0.1 mol dm3 KCl solution. The solubility of ½C1;n pyrrþ ½C2 C2 N in water was determined by precipitation titration of C1,npyrrl+ with an aqueous solution of 10 mmol dm3 sodium tetraphenylborate solution using an automatic titrator (Hiranuma, COM1600). The solubility of water in ½C1;n pyrrþ ½C2 C2 N was measured with a Karl Fischer moisture meter (Mitsubishi Chemical Analytech, CA-21).
3. Results and discussion 3.1. Physicochemical properties of ½C1;n pyrrþ ½C2 C2 N
Potentiometric measurements were made using the following cell,
þ The physicochemical properties of water-saturated C1;n pyrrl ½C2 C2 N are listed in Table 1. In comparison with C8mimC1C1N, the ILs, ½C1;n pyrrþ ½C2 C2 N (n = 6, 7, 8), have low mutual solubility with water; the solubility of these ILs in W is (1.5 ± 0.5) 104 mol dm3, at most, which is one order of magnitude less than that of C8mimC1C1N. The large uncertainty is due to the
where C1,npyrrBr is N-alkyl-N-methylpyrrolidinium bromide, n denotes the number of carbon atoms in the alkyl moiety, and MBr is either HBr, LiBr, or KBr. Instead of Ag/AgCl electrodes, we employed Ag/AgBr electrodes, because the solubility of silver chloride in water is 13 lmol dm3 at 25 °C and its dissolution into the
difficulty in accurately determining the end point in the titration. The solubility of water in the ILs, 0.07–0.45%, is significantly lower in comparison with the case of C8mimC1C1N, 0.96% at 25 °C [3]. The lower mutual solubility is preferable in view of the stability of ILSBs.
2.2. Methods
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Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66 Table 1 Physicochemical properties of water-saturated N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide at 25 °C. Cation
Density a mmol dm3
Viscosity g cm3
Conductivity m Pa s
solubility of IL in W mS cm1
solubility of W in IL wt%
C1,8Pyrr+ C1,7Pyrr+ C1,6Pyrr+
1.37 1.40 1.42
266 241 226
0.342 0.436 0.510
0.2 0.1 0.2
0.45 0.07 0.08
a
At 25 ± 2 °C.
3.1.1. Solubility of AgX in dilute alkali halide solutions Generally, when a sufficient amount of silver halide, AgX, is added to an aqueous solution of alkali halide, KX, whose concentration is c0KX , AgX dissolves in the solution depending on c0KX . At equilibrium when residual solid AgX is coexistent with the solution, the equilibrium concentration of X ; cX , can be calculated as follows. The conditions of the mass balance with respect to K+ and the charge balance in W are combined to the definition of the solubility product of AgX in W to give the equilibrium concentration of X in W, cX ,
c X ¼
qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 0 2 1 0 cKX þ 4K SP cKX þ 2
ð1Þ
where c0KX is the initial concentration of KX, a strong 1–1 electrolyte, in W and KSP is the solubility product of AgX. Here we have neglected the presence of neutral AgX dissolved in the solution and assumed that the activities of the relevant ionic species were equal to their concentrations. For the values of solubility products for AgCl, AgBr, and AgI, 1.80 1010, 5.00 1013, and 8.30 1017 mol2 dm6 at 25 °C, [17]. Some representative values of the equilibrium halide concentration are listed for several values of c0KX in Table 2. It can be seen that when c0KX = 10 lmol dm3, cX in the presence of AgCl becomes 19 lmol dm3, that is, almost twice as large as c0KX . This corresponds to the difference of 0.27 in the common logarithm of the concentration. But, in the case of AgBr, the difference between c0KX and cKX is negligible when c0KX = 10 lmol dm3 and is only 2% when c0KX = 5 lmol dm3. Cell (I) with Ag/AgBr electrodes can hence be used in such a lower concentration range to examine the constancy of the phase-boundary potential between the IL and the aqueous solution. 3.2. Time dependence of the phase-boundary potential between the ionic liquid salt bridge and a dilute aqueous solution of HBr, LiBr, and KBr The values of E of Cell (I) were generally very stable for all three liquids in aqueous solutions of HBr, LiBr, and KBr. Fig. 1a–c illustrates the time courses of E for three ILSBs, ½C1;8 pyrrþ ½C2 C2 N (Fig. 1a), ½C1;7 pyrrþ ½C2 C2 N (Fig. 1b), and ½C1;6 pyrrþ ½C2 C2 N (Fig. 1c) at different concentrations of HBr (Fig. 1a), LiBr (Fig. 1b) and KBr (Fig. 1c) between 5 lmol kg1 and 500 lmol kg1. In these
Table 2 Calculated values of equilibrium concentration of halides solutions in the presence of AgX precipitates at 25 °C. c0KX is the feed concentration of KX in molarity and cX denotes the equilibrium concentration of X. cKX0 1.00 l04 5.00 105 2.00 105 1.00 l05 5.00 106 2.00 106 1.00 106
cCl
cBr
cI
1.02 104 5.34 105 2.67 105 1.93 105 1.61 105 1.44 105 1.39 105
1.00 l04 5.00 105 2.00 105 1.00 l05 5.10 106 2.22 106 1.37 106
1.00 l04 5.00 105 2.00 105 1.00 105 1.00 106 2.00 106 1.00 106
Fig. 1. Time course of the cell voltage of Cell (I) for three ILSBs, ½C1;8 pyrrþ ½C2 C2 N (a), ½C1;7 pyrrþ ½C2 C2 N (b), and ½C1;6 pyrrþ ½C2 C2 N (c) at different concentrations of HBr (a), LiBr (b) and KBr (c) between 5 lmol kg1 and 500 lmol kg1 at 25 °C.
measurements, time was set to zero after the reading of E was stabilized, typically, 2–3 min after the setting of the cell in the dark box. For all the ILSBs, the excursion of E from the average value for 15 min was within ±1 mV for all cases examined. The standard deviation in a single run was typically 0.2 mV except at the lowest concentration, 5 lmol kg1 KBr, at which the standard deviation was 0.7 mV. The pooled standard deviation of all data in Fig. 1b was 0.1 mV. When the measurements were made under the illumination from a fluorescent ceiling light, the E was drifted away from the value under dark, considerably, e.g., more than 1 mV in 15 min. The E value was in fact stable over more than 15 min shown in Fig. 1. To illustrate this stability over a longer time span, Fig. 2 illustrates a few examples of the stability of E for ½C1;6 pyrrþ ½C2 C2 N in
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Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66
Fig. 2. Long-term stability of the cell voltage of Cell (I) for ILSB made of ½C1;6 pyrrþ ½C2 C2 N in contact with 5 (1), 10 (2), 100 (3), or 500 (4) lmol kg1 HBr at 25 °C.
contact with 5 (1), 10 (2), 100 (3), or 500 (4) lmol kg1 HBr for 1 h. The data 5 min after the setting of Cell (I) were particularly stable: the excursion is typically within ±0.5 mV. 3.3. Constancy of the phase-boundary potential between the ionic liquid salt bridge and aqueous solutions of different compositions Fig. 3 shows E values for ½C1;8 pyrrþ ½C2 C2 N (red marks, upper line), ½C1;7 pyrrþ ½C2 C2 N (black marks, middle line), and ½C1;6 pyrrþ ½C2 C2 N (blue marks, lower line) as a function of the mean activity of MBr, a MBr , where MBr is either KBr (filled circles with error bars), LiBr (triangle), or HBr (filled squares). The concentration range shown in Fig. 3 is about 5 and 500 lmol kg1, where the actual molal concentrations of MBr varied on preparation of the solutions. For the mean activity coefficients above m = 1 mmol kg1, the literature values [18] were used and for those below this concen€ ckel tration the values were calculated using the Debye–Hu limiting law, assuming that only M+(M+is K+, Li+, or H+) and Br contribute to the ionic strength. The error bars show the standard deviations for triplicate measurements for KBr solutions with renewed settings of Cell (I). The straight lines have the slope of 59.2 mV/decade, which is expected for the Nernstian response of the Ag/AgBr electrode on the
left-hand-side of Cell (I). One can see that all experimental points in the concentration range between 20 and 500 lmol kg1 ð4:8 < log10 a MBr < 3:3Þ are on the straight line having a slope of 59.2 mV/decade for all ILSBs. Since this Nernstian response of E is due to the Ag/AgBr electrode on the right-hand-side of Cell (I), the results in Fig. 3 demonstrate that DW IL / is stable down to at least 20 lmol kg1, and possibly 10 lmol kg1, KBr, LiBr, and HBr solutions. The points at 5 lmol kg1 outlie the straight lines for all three data sets. The points at 10 lmol kg1 also seem to be slightly off the straight lines but the lines are within the errors bars of the data points. The pooled standard deviation of the deviations from the straight lines calculated for three data sets at five higher concentrations of KBr was 0.4 mV. The results in Fig. 3 assure that the phase-boundary potential between the ILSB and the aqueous solution (Phase V), DW IL /, remains constant over the change in the concentration of KCl, LiBr, and HBr, from 20 to 500 lmol kg1 with the standard deviation of 0.4 mV. The constancy of DW IL / was thus confirmed between 20 and 500 lmol kg1 KBr, LiBr, and HBr for all the ILs examined. It is interesting to see such a stability at lower concentrations of MBr, because the solubility of the ILs on the order of 200 lmol kg1 can affect the mean activity of MBr in Phase V and causes the change in the potential of the Ag/AgBr electrode on the righthand-side of Cell (I). In fact, this effect is too small, if any, to be € ckel limdetectable in the plots of the type in Fig. 2. The Debye–Hu iting law estimates the activity coefficient of 200 lmol kg1 being 0.9835 at 25 °C. For example, the use of this value to calculate the mean activity of MBr at 5 lmol kg1 changes the log10 a MBr only from 5.302 to 5.308. 3.4. Dependence of E on the type of ILs Fig. 3 shows the dependence of E at a given value of a MBr on the alkyl-chain length of C1,npyrr+. We consider the major factor that gives rise to this height difference in the plots. The E may be decomposed into several contributions.
E ¼ /VII /I VII ¼ / /V þ /V /IV þ /IV /III þ /III /I
ð2Þ
where /a is the inner potential of the phase a, where the superscript a designates the phase I, III, IV, V, or VII in Cell (I). For the two Ag/ AgBr/Br electrodes in Cell (I), we may write
/VII /V ¼ D/0Ag=AgBr=Br
RT ln aVBr F
ð3Þ
and
/III /I ¼ D/0Ag=AgBr=Br þ
RT ln aIII Br F
ð4Þ
where D/0Ag=AgBr=Br is the standard potential of the Ag/AgBr/Br electrode, R is the gas constant, T is the absolute temperature, F is the Faraday constant, and aaBr is the activity of Br in Phase a (a = III or V). Assuming that the partition equilibrium is established between Phases IV and V, we may write [2]
/V /IV ¼ + + + Fig. 3. Plots of E as a function of the mean activity of MBr, a MBr (M = K , Li , and H for ½C1;8 pyrrþ ½C2 C2 N (red marks, upper line), ½C1;7 pyrrþ ½C2 C2 N (black marks, middle line), and ½C1;6 pyrrþ ½C2 C2 N (blue marks, lower line) at 25 °C. Error bars in (a) are from triplicate measurements of KBr solutions. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.)
RT cV cIVþ 1 W 0 0 ln A C D / þ þ DW IL /A þ 2 IL C 2F cVCþ cIV A
ð5Þ
0 where DW IL /i is the standard ion transfer potentials of the IL-constituting ion i (i = C+ or A) between Phases IV and V and cai is the activity coefficient of i in Phase a (a = IV or V) [2]. Since C+ is common to III and IV, that is, the potential determining ion, we can write [1]
Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66 0 /IV /III ¼ DW IL /Cþ
RT ln aIII Cþ F
65
ð6Þ
By adding Eqs. (3)–(6), we obtain
E¼
RT aIII 1 1 RT ln Br ln aIII DW /0þ þ DW /0 Cþ F 2 IL A F aVBr 2 IL C
ð7Þ
where we have neglected the activity coefficient term in Eq. (5). At given concentrations of C+Br in Phase III and of M+Br (M = H+, Li+, or K+ in the present case) in Phase V in Cell (I), the effect of the difference in the ILSB-constituent cation, C1,npyrr+, on E is
DE ¼ EðC1;n pyrrþ Þ EðC1;n1 pyrrþ Þ
1 0 ¼ DW /0 DW IL /C1;n1pyrrþ 2 IL C1;npyrrþ
ð8Þ
where E(C1,npyrr+) and E(C1,n1pyrr+) are the cell voltages obtained using the IL ½C1;n pyrrþ ½C2 C2 N and ½C1;n1 pyrrþ ½C2 C2 N , respectively. The assumptions made in deriving Eq. (8) are that the activities of Br in HBr, LiBr, and KBr, are the same at a given concentration of MBr and that the activities of C1,npyrr+ in Phase III are the same for different n. The vertical distance between the straight lines in Fig. 3 can then be ascribed to the difference in 0 DW IL /C1;npyrrþ , i.e, the difference in hydrophobicity, of the IL-constituW 0 0 ent cations. The average values for DW IL /C1;8pyrrþ DIL /C1;7pyrrþ and W 0 0 DW IL /C1;7pyrrþ DIL /C1;6pyrrþ estimated from the plots in Fig. 3 are 15
and 12 mV, respectively. 0 The effect of the elongation of one methylene unit on DW IL /i is 30 mV at 25 °C in the case of 1-alkyl-3-methylimidazolium C1C1N-W two-phase systems [15]. The half of this value, 15 mV, is comparable to the observed difference in the height of the straight lines above and confirms the prediction of Eq. (8). 3.5. Contribution of diffusion potential due to the dissolution of the ILSB constituting ions We examine if the deviation from the straight line at 5 lmol kg1 is ascribed to the diffusion potential created by the dissolution of ILSB-constituent ions in W. To estimate the magnitude of this diffusion potential, we first measured the limiting molar conductivity of C1,npyrrBr and lithium bis(pentafluoroethanesulfonyl)amide at 25 °C. The values obtained are 115, 119 and 122 S cm2 mol 1 for n = 8, 7, and 6, and 67.1 S cm2 mol 1 for lithium bis(pentafluoroethanesulfonyl)amide. From literature values of the limiting molar ionic conductivity of Li+ and Br, we estimated the limiting molar ionic conductivity values as 33(C1,8pyrr+), 38(C1,7pyrr+), 41(C1,6pyrr+) and 28(C2C2N) S cm2 mol1, respectively. Since the mobility of C2C2N is smaller than those of C1,npyrr+, the diffusion potential referred to the inner potential in the bulk phase of Phase V should be negative and its magnitude depends on the concentration of indifferent ions in V [4]. For example, the values of the phase-boundary potential between Phases IV and V when the concentration of MBr in Phase V is 5 lmol kg1 are 2.9 (n = 8), 4.9 (n = 7), and 6.3 (n = 6) mV. As E is measured as /VII /I, this diffusion potential is opposite to the experimentally observed deviation of E at this concentration of MBr in Phase V. According to Eq. (1) and Table 1 the dissolution of AgBr in Phase V is small at cVMBr = 5 lmol kg1 and cannot explain the downward deviation from the straight line. The change in the value of a MBr is too small in the logarithmic scale as discussed above. At this moment, we have no explanation for this discrepancy. The presence of a certain negative bias of about 10 mV in E should be assumed at this concentration. More precise measurements of the molar or molal ion conductivity of relevant ions seem to be required, be left the applicability of the Henderson equation.
+ Fig. 4. Plots of E as a function of the mean activity of MBr, a MBr (M = K (red filled squares) and H+ (open circles)) for ½C1;6 pyrrþ ½C2 C2 N at 25 °C. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.)
3.6. Constancy of the liquid junction potential in a wider range of the MBr concentration The constancy of the DW IL / between the ILSB and an aqueous solution extends to higher concentration range than that shown in Fig. 3. As an example, Fig. 4 shows the plots of E against þ log10 a MBr for ½C1;7 pyrr ½C2 C2 N in the wider concentration range. W The constancy of the DIL / at the ILSB j aqueous solution is maintained up to 0.5 mol kg1 KBr and HBr solutions. In other words, the DW IL / stays constant over four orders of magnitude change in the concentration of MBr in W. At higher concentrations, the upward deviation took place in HBr solutions, as has been seen in the ILSB made of C8mimC1C1N due to the penetration of H+ in the ILSB [2]. 4. Conclusions Three ionic liquids that consist of N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide, where alkyl- is either hexyl-, heptyl- or octyl-, are suitable to an ionic liquid salt bridge for electrochemical measurements, particularly, potentiometry, of low ionic strength aqueous solutions down to the ionic strength of 20 lmol kg1. The remarkable stability of the phase-boundary potential with the standard deviation of 0.4 mV at such low ionic € ckel limiting law is relistrength solutions, where the Debye–Hu able for calculating the activity of ions, opens the way to accurately estimate single ion activities, including pH. Acknowledgements This work was partly supported by Japan Science and Technology Agency under the program, ‘‘Development of Systems and Technology for Advanced Measurement and Analysis’’ and by Grant-in-Aid for Scientific Research (No. 21245021) from the Ministry of Educations, Sports, Science, and Technology, Japan. Support by the Global COE Program, International Center for Integrated Research and Advanced Education in Materials Science (No. B-09) from the Ministry of Education, Culture, Sports, Science and Technology of Japan is highly appreciated. References [1] T. Kakiuchi, N. Tsujioka, S. Kurita, Y. Iwami, Electrochem. Commun. 5 (2003) 159–164.
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Journal of Electroanalytical Chemistry journal homepage: www.elsevier.com/locate/jelechem
Ionic liquid salt bridge based on N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide for low ionic strength aqueous solutions Yousuke Fujino, Takashi Kakiuchi ⇑ Department of Energy and Hydrocarbon Chemistry, Graduate School of Engineering, Kyoto University, Kyoto 615-8510, Japan
a r t i c l e
i n f o
Article history: Received 29 July 2010 Received in revised form 17 September 2010 Accepted 30 October 2010 Available online 18 November 2010 Keywords: Ionic liquids Ionic liquid salt bridge Low ionic strength samples Liquid junction potential Distribution potential N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide
a b s t r a c t The phase-boundary potential between the moderately hydrophobic ionic liquid and a low ionic strength aqueous solution is demonstrated to be stable and constant with the standard deviation of 0.4 mV down to 20 lmol kg1 HBr, LiBr, and KBr solutions, for three ionic liquids that consist of either N-methyl-Noctylpyrrolidinium, N-heptyl-N-methylpyrrolidinium, or N-hexyl-N-methylpyrrolidinium and a common anion species, bis(pentafluoroethanesulfonyl)amide. This stability is promising for accurate measurements of pH of low ionic strength samples and reliable estimates of single ion activities in general. The phase-boundary potential deviates from the value determined by the partition of the ionic liquid in further dilute aqueous solutions. The magnitude of the deviation ranges from 3 to 11 mV at 5 lmol kg1 MBr (M is H+, Li+, or K+). The solubility of these ionic liquids in water is 0.2 mmol dm3 at most at 25 °C, which is another advantage of ionic liquid salt bridge in electroanalytical chemistry. Ó 2010 Elsevier B.V. All rights reserved.
1. Introduction An ionic liquid (IL) that is moderately hydrophobic and immiscible with water can work as a salt bridge having several advantages over traditional KCl-type salt bridges [1,2]. A notable property of ionic liquid salt bridge (ILSB) is their stability of the phaseboundary potential when it is in contact with a low ionic strength aqueous solution. In the case of 1-methyl-3-octylimidazolium bis(trifluoromethanesulfonyl)amide (C8mimC1C1N) whose solubility in water is 1.6 mmol dm3 at 25 °C [3], the phase-boundary potential stays constant within ±1 mV in a 100 lmol dm3 aqueous KCl solution [4]. However, when the ionic strength of a sample solution is lower than this concentration range, the deviation from the constant value becomes greater and amounts to 20 mV at 20 lmol dm3 KCl [4]. This deviation was ascribed, mainly, to the diffusion potential formed in the aqueous phase accompanied with the dissolution of C8mimC1C1N, where the mobility of C8mim+ is less than that of C1C1N. Although the ionic strength on the order of 100 lmol dm3 may be seen dilute in typical potentiometric measurements [5,6], the ionic strength of actual samples can be less than this level. For example, rainwater in nonpolluted regions usually has the ionic strength of 50 lmol dm3 or lower [7–9].
⇑ Corresponding author. Tel.: +81 75 383 2489; fax: +81 75 383 2490. E-mail address: [email protected] (T. Kakiuchi). 1572-6657/$ - see front matter Ó 2010 Elsevier B.V. All rights reserved. doi:10.1016/j.jelechem.2010.10.028
In such low ionic strength solutions, the liquid junction potential between a KCl-type salt bridge and a sample solution is the main source of inaccuracy in determining pH [9–12], because the potential is not stable and can amount to more than 10 mV. One more serious problem is the elevation of the ionic strength of sample solutions due to the dissolution of KCl, which changes the activity of analyte ions, for example, H+. Given the stability of ILSB based on C8mimC1C1N in relatively low ionic strength solutions [2,4,13], well-designed ILSBs that have smaller solubility than C8mimC1C1N and exhibit a negligibly small diffusion potential in low ionic strength samples are promising in accurate electroanalytical chemistry, in particular, in the potentiometric determination of pH of low ionic strength samples. For such samples, the lower solubility of ILSB-constituent ions in W is preferable, because the contribution of the diffusion potential due to the dissolution becomes smaller [4], and the degree of sample contamination is less. Although the lower solubility means the higher electrochemical polarizability of the ILSBjW interface [14], the concentration level of interfering ions, if any, would also be small in such low ionic strength samples. Toward this goal, we have studied several moderately hydrophobic ILs. We herein report the use of N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide, where the alkyl group is either octyl-, heptyl-, or hexyl moiety, for ILSB. Aside from the solubility, it is important in choosing ions for ILSB to balance the hydrophobicity of the IL-constituting cation and that of the anion, so that
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the expected phase-boundary potential is close to 0 [15]. Judging from the standard ion transfer potentials of N-methyl-Noctylpyrrolidinium and bis(pentafluoroethanesulfonyl)amide in the nitrobenzene–water two-phase system, 0.23 and 0.2 V [16], respectively, it is expected, first, that the phase-boundary potential between the ionic liquid and water (W) would be close to zero [1]. Second, the solubility of this ionic liquid in water, whose measure is the difference in the standard ion transfer potentials of IL-constituting cation and anion [1], should be smaller than that of C8mimC1C1N, where the constituent cation and anion have the standard ion transfer potentials of 0.24 and 0.13 V in the same two-phase system, respectively [16]. We will show that the ionic liquids that consist of N-alkyl-N-methylpyrrolidinium and bis(pentafluoroethanesulfonyl)amide exhibit a stable liquid junction potential down to 20 lmol kg1 solutions of HBr, LiBr, and KBr and are hence suitable to a salt bridge for pH measurements of low ionic strength samples. 2. Experimental 2.1. Materials N-methylpyrrolidine (Tokyo Chem. Ind., TCI), 1-bromooctane (TCI), 1-bromoheptane (Aldrich), 1-bromohexane (TCI), N-methylN-octylpyrrolidinium chloride (Merck), and hydrogen bis(pentafluoroethanesufonyl)amide (Central Glass, 70% aqueous solution) were used without further purification. Other chemicals used were of reagent grade. N-alkyl-N-methylpyrrolidinum bromides were prepared from methylpyrrolidine and 1-alkylbromide by stirring the mixture of 1.05 (methylpyrrolidine):1.00 (1-alkylbromide) at about 80 °C for at least 6 h. The mixture was washed three times with ethylacetate and then volatile components were removed by drying the mixture first with an evaporator and then with a vacuum pump. N-alkyl-N-methylpyrrolidinium bromides thus obtained were identified with 1H NMR. Three salts, N-methylN-octylpyrrolidinium bis(pentafluoroethanesufonyl)amide, Nheptyl-N-methylpyrrolidinium bis(pentafluoroethanesufonyl)amide, and N-hexyl-N-methylpyrrolidinium bis(pentafluoroethanesufonyl)amide, were prepared by metathesis reactions between Nalkyl-N-methylpyrrolidinium bromides and bis(pentafluoroethanesufonyl)amide acid. These salts will hereafter be abbreviated as ½C18 pyrrþ ½C2 C2 N ; ½C17 pyrrþ ½C2 C2 N , and ½C16 pyrrþ ½C2 C 2 N , respectively. These salts were washed repeatedly with copious water until no precipitation of AgBr was detected when a few drops of a AgNO3 solution were added in washing water.
sample solution is non-negligible when the initial concentration, as prepared, of chloride in a sample solution is comparable to the solubility of AgCl. Cell (I) was constructed in a glass tube with upper and lower compartments separated by a disk of glass frit, as described elsewhere [2]. The cell was housed in a box for light shielding. The temperature of the cell was maintained at 25 ± 0.5 °C by circulating water through the jacket of the glass cell [2]. The cell voltage, i.e., the potential of the right-hand-side terminal with respect to the that of the left, E, was measured with an electrometer (ADC, ADCMT-8252), which was connected to a computer through GPIB, as described elsewhere [2]. The sampling rate was typically 0.2 Hz. Ag/AgBr electrodes were prepared by anodization of silver wires of 0.5 mm diameter in an aqueous solution of ca. 80 mmol dm3 KBr at 10 mA for 15 min. Before the AgBr coating, the silver wire was first polished with an emery paper and then treated by ultrasonic cleaning successively with Milli-Q water, 2% aqueous ammonia, 2% nitric acid, and Milli-Q water, for 30 min in each step. The electrodes were kept in Milli-Q water in the dark. þ The viscosity of water-saturated C1;n pyrrl ½C2 C2 N was measured with a viscometer (Toki Sangyo, TVE-33, cone-plate type) at 1, 2, 5, 10, 20, and 50 rpm at 25 ± 0.5 °C by circulating water in the cell. No significant dependence of the measured viscosity on the share rate was observed. The density of ½C1;n pyrrþ ½C2 C2 N was measured with an Ostwald pycnometer. The electrical conductivity of water-saturated ½C1;n pyrrþ ½C2 C2 N was measured with a conductivity meter (Kyoto Electronics, CM-117). The electrical conductivity of electrolyte solutions to evaluate the molar conductivity of ions in water was measured with a bridge-type conductivity meter (Husou, HECS362D with HECS363D) and a dip-in-type cell (DKK-TOA, CT-5710B) at four concentrations of C1,npyrrolidinium bromides and bis(pentafluoroethanesufonyl)amide acid in a water bath at 25 ± 0.02 °C. The cell constant was determined using a 0.1 mol dm3 KCl solution. The solubility of ½C1;n pyrrþ ½C2 C2 N in water was determined by precipitation titration of C1,npyrrl+ with an aqueous solution of 10 mmol dm3 sodium tetraphenylborate solution using an automatic titrator (Hiranuma, COM1600). The solubility of water in ½C1;n pyrrþ ½C2 C2 N was measured with a Karl Fischer moisture meter (Mitsubishi Chemical Analytech, CA-21).
3. Results and discussion 3.1. Physicochemical properties of ½C1;n pyrrþ ½C2 C2 N
Potentiometric measurements were made using the following cell,
þ The physicochemical properties of water-saturated C1;n pyrrl ½C2 C2 N are listed in Table 1. In comparison with C8mimC1C1N, the ILs, ½C1;n pyrrþ ½C2 C2 N (n = 6, 7, 8), have low mutual solubility with water; the solubility of these ILs in W is (1.5 ± 0.5) 104 mol dm3, at most, which is one order of magnitude less than that of C8mimC1C1N. The large uncertainty is due to the
where C1,npyrrBr is N-alkyl-N-methylpyrrolidinium bromide, n denotes the number of carbon atoms in the alkyl moiety, and MBr is either HBr, LiBr, or KBr. Instead of Ag/AgCl electrodes, we employed Ag/AgBr electrodes, because the solubility of silver chloride in water is 13 lmol dm3 at 25 °C and its dissolution into the
difficulty in accurately determining the end point in the titration. The solubility of water in the ILs, 0.07–0.45%, is significantly lower in comparison with the case of C8mimC1C1N, 0.96% at 25 °C [3]. The lower mutual solubility is preferable in view of the stability of ILSBs.
2.2. Methods
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Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66 Table 1 Physicochemical properties of water-saturated N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide at 25 °C. Cation
Density a mmol dm3
Viscosity g cm3
Conductivity m Pa s
solubility of IL in W mS cm1
solubility of W in IL wt%
C1,8Pyrr+ C1,7Pyrr+ C1,6Pyrr+
1.37 1.40 1.42
266 241 226
0.342 0.436 0.510
0.2 0.1 0.2
0.45 0.07 0.08
a
At 25 ± 2 °C.
3.1.1. Solubility of AgX in dilute alkali halide solutions Generally, when a sufficient amount of silver halide, AgX, is added to an aqueous solution of alkali halide, KX, whose concentration is c0KX , AgX dissolves in the solution depending on c0KX . At equilibrium when residual solid AgX is coexistent with the solution, the equilibrium concentration of X ; cX , can be calculated as follows. The conditions of the mass balance with respect to K+ and the charge balance in W are combined to the definition of the solubility product of AgX in W to give the equilibrium concentration of X in W, cX ,
c X ¼
qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 0 2 1 0 cKX þ 4K SP cKX þ 2
ð1Þ
where c0KX is the initial concentration of KX, a strong 1–1 electrolyte, in W and KSP is the solubility product of AgX. Here we have neglected the presence of neutral AgX dissolved in the solution and assumed that the activities of the relevant ionic species were equal to their concentrations. For the values of solubility products for AgCl, AgBr, and AgI, 1.80 1010, 5.00 1013, and 8.30 1017 mol2 dm6 at 25 °C, [17]. Some representative values of the equilibrium halide concentration are listed for several values of c0KX in Table 2. It can be seen that when c0KX = 10 lmol dm3, cX in the presence of AgCl becomes 19 lmol dm3, that is, almost twice as large as c0KX . This corresponds to the difference of 0.27 in the common logarithm of the concentration. But, in the case of AgBr, the difference between c0KX and cKX is negligible when c0KX = 10 lmol dm3 and is only 2% when c0KX = 5 lmol dm3. Cell (I) with Ag/AgBr electrodes can hence be used in such a lower concentration range to examine the constancy of the phase-boundary potential between the IL and the aqueous solution. 3.2. Time dependence of the phase-boundary potential between the ionic liquid salt bridge and a dilute aqueous solution of HBr, LiBr, and KBr The values of E of Cell (I) were generally very stable for all three liquids in aqueous solutions of HBr, LiBr, and KBr. Fig. 1a–c illustrates the time courses of E for three ILSBs, ½C1;8 pyrrþ ½C2 C2 N (Fig. 1a), ½C1;7 pyrrþ ½C2 C2 N (Fig. 1b), and ½C1;6 pyrrþ ½C2 C2 N (Fig. 1c) at different concentrations of HBr (Fig. 1a), LiBr (Fig. 1b) and KBr (Fig. 1c) between 5 lmol kg1 and 500 lmol kg1. In these
Table 2 Calculated values of equilibrium concentration of halides solutions in the presence of AgX precipitates at 25 °C. c0KX is the feed concentration of KX in molarity and cX denotes the equilibrium concentration of X. cKX0 1.00 l04 5.00 105 2.00 105 1.00 l05 5.00 106 2.00 106 1.00 106
cCl
cBr
cI
1.02 104 5.34 105 2.67 105 1.93 105 1.61 105 1.44 105 1.39 105
1.00 l04 5.00 105 2.00 105 1.00 l05 5.10 106 2.22 106 1.37 106
1.00 l04 5.00 105 2.00 105 1.00 105 1.00 106 2.00 106 1.00 106
Fig. 1. Time course of the cell voltage of Cell (I) for three ILSBs, ½C1;8 pyrrþ ½C2 C2 N (a), ½C1;7 pyrrþ ½C2 C2 N (b), and ½C1;6 pyrrþ ½C2 C2 N (c) at different concentrations of HBr (a), LiBr (b) and KBr (c) between 5 lmol kg1 and 500 lmol kg1 at 25 °C.
measurements, time was set to zero after the reading of E was stabilized, typically, 2–3 min after the setting of the cell in the dark box. For all the ILSBs, the excursion of E from the average value for 15 min was within ±1 mV for all cases examined. The standard deviation in a single run was typically 0.2 mV except at the lowest concentration, 5 lmol kg1 KBr, at which the standard deviation was 0.7 mV. The pooled standard deviation of all data in Fig. 1b was 0.1 mV. When the measurements were made under the illumination from a fluorescent ceiling light, the E was drifted away from the value under dark, considerably, e.g., more than 1 mV in 15 min. The E value was in fact stable over more than 15 min shown in Fig. 1. To illustrate this stability over a longer time span, Fig. 2 illustrates a few examples of the stability of E for ½C1;6 pyrrþ ½C2 C2 N in
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Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66
Fig. 2. Long-term stability of the cell voltage of Cell (I) for ILSB made of ½C1;6 pyrrþ ½C2 C2 N in contact with 5 (1), 10 (2), 100 (3), or 500 (4) lmol kg1 HBr at 25 °C.
contact with 5 (1), 10 (2), 100 (3), or 500 (4) lmol kg1 HBr for 1 h. The data 5 min after the setting of Cell (I) were particularly stable: the excursion is typically within ±0.5 mV. 3.3. Constancy of the phase-boundary potential between the ionic liquid salt bridge and aqueous solutions of different compositions Fig. 3 shows E values for ½C1;8 pyrrþ ½C2 C2 N (red marks, upper line), ½C1;7 pyrrþ ½C2 C2 N (black marks, middle line), and ½C1;6 pyrrþ ½C2 C2 N (blue marks, lower line) as a function of the mean activity of MBr, a MBr , where MBr is either KBr (filled circles with error bars), LiBr (triangle), or HBr (filled squares). The concentration range shown in Fig. 3 is about 5 and 500 lmol kg1, where the actual molal concentrations of MBr varied on preparation of the solutions. For the mean activity coefficients above m = 1 mmol kg1, the literature values [18] were used and for those below this concen€ ckel tration the values were calculated using the Debye–Hu limiting law, assuming that only M+(M+is K+, Li+, or H+) and Br contribute to the ionic strength. The error bars show the standard deviations for triplicate measurements for KBr solutions with renewed settings of Cell (I). The straight lines have the slope of 59.2 mV/decade, which is expected for the Nernstian response of the Ag/AgBr electrode on the
left-hand-side of Cell (I). One can see that all experimental points in the concentration range between 20 and 500 lmol kg1 ð4:8 < log10 a MBr < 3:3Þ are on the straight line having a slope of 59.2 mV/decade for all ILSBs. Since this Nernstian response of E is due to the Ag/AgBr electrode on the right-hand-side of Cell (I), the results in Fig. 3 demonstrate that DW IL / is stable down to at least 20 lmol kg1, and possibly 10 lmol kg1, KBr, LiBr, and HBr solutions. The points at 5 lmol kg1 outlie the straight lines for all three data sets. The points at 10 lmol kg1 also seem to be slightly off the straight lines but the lines are within the errors bars of the data points. The pooled standard deviation of the deviations from the straight lines calculated for three data sets at five higher concentrations of KBr was 0.4 mV. The results in Fig. 3 assure that the phase-boundary potential between the ILSB and the aqueous solution (Phase V), DW IL /, remains constant over the change in the concentration of KCl, LiBr, and HBr, from 20 to 500 lmol kg1 with the standard deviation of 0.4 mV. The constancy of DW IL / was thus confirmed between 20 and 500 lmol kg1 KBr, LiBr, and HBr for all the ILs examined. It is interesting to see such a stability at lower concentrations of MBr, because the solubility of the ILs on the order of 200 lmol kg1 can affect the mean activity of MBr in Phase V and causes the change in the potential of the Ag/AgBr electrode on the righthand-side of Cell (I). In fact, this effect is too small, if any, to be € ckel limdetectable in the plots of the type in Fig. 2. The Debye–Hu iting law estimates the activity coefficient of 200 lmol kg1 being 0.9835 at 25 °C. For example, the use of this value to calculate the mean activity of MBr at 5 lmol kg1 changes the log10 a MBr only from 5.302 to 5.308. 3.4. Dependence of E on the type of ILs Fig. 3 shows the dependence of E at a given value of a MBr on the alkyl-chain length of C1,npyrr+. We consider the major factor that gives rise to this height difference in the plots. The E may be decomposed into several contributions.
E ¼ /VII /I VII ¼ / /V þ /V /IV þ /IV /III þ /III /I
ð2Þ
where /a is the inner potential of the phase a, where the superscript a designates the phase I, III, IV, V, or VII in Cell (I). For the two Ag/ AgBr/Br electrodes in Cell (I), we may write
/VII /V ¼ D/0Ag=AgBr=Br
RT ln aVBr F
ð3Þ
and
/III /I ¼ D/0Ag=AgBr=Br þ
RT ln aIII Br F
ð4Þ
where D/0Ag=AgBr=Br is the standard potential of the Ag/AgBr/Br electrode, R is the gas constant, T is the absolute temperature, F is the Faraday constant, and aaBr is the activity of Br in Phase a (a = III or V). Assuming that the partition equilibrium is established between Phases IV and V, we may write [2]
/V /IV ¼ + + + Fig. 3. Plots of E as a function of the mean activity of MBr, a MBr (M = K , Li , and H for ½C1;8 pyrrþ ½C2 C2 N (red marks, upper line), ½C1;7 pyrrþ ½C2 C2 N (black marks, middle line), and ½C1;6 pyrrþ ½C2 C2 N (blue marks, lower line) at 25 °C. Error bars in (a) are from triplicate measurements of KBr solutions. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.)
RT cV cIVþ 1 W 0 0 ln A C D / þ þ DW IL /A þ 2 IL C 2F cVCþ cIV A
ð5Þ
0 where DW IL /i is the standard ion transfer potentials of the IL-constituting ion i (i = C+ or A) between Phases IV and V and cai is the activity coefficient of i in Phase a (a = IV or V) [2]. Since C+ is common to III and IV, that is, the potential determining ion, we can write [1]
Y. Fujino, T. Kakiuchi / Journal of Electroanalytical Chemistry 651 (2011) 61–66 0 /IV /III ¼ DW IL /Cþ
RT ln aIII Cþ F
65
ð6Þ
By adding Eqs. (3)–(6), we obtain
E¼
RT aIII 1 1 RT ln Br ln aIII DW /0þ þ DW /0 Cþ F 2 IL A F aVBr 2 IL C
ð7Þ
where we have neglected the activity coefficient term in Eq. (5). At given concentrations of C+Br in Phase III and of M+Br (M = H+, Li+, or K+ in the present case) in Phase V in Cell (I), the effect of the difference in the ILSB-constituent cation, C1,npyrr+, on E is
DE ¼ EðC1;n pyrrþ Þ EðC1;n1 pyrrþ Þ
1 0 ¼ DW /0 DW IL /C1;n1pyrrþ 2 IL C1;npyrrþ
ð8Þ
where E(C1,npyrr+) and E(C1,n1pyrr+) are the cell voltages obtained using the IL ½C1;n pyrrþ ½C2 C2 N and ½C1;n1 pyrrþ ½C2 C2 N , respectively. The assumptions made in deriving Eq. (8) are that the activities of Br in HBr, LiBr, and KBr, are the same at a given concentration of MBr and that the activities of C1,npyrr+ in Phase III are the same for different n. The vertical distance between the straight lines in Fig. 3 can then be ascribed to the difference in 0 DW IL /C1;npyrrþ , i.e, the difference in hydrophobicity, of the IL-constituW 0 0 ent cations. The average values for DW IL /C1;8pyrrþ DIL /C1;7pyrrþ and W 0 0 DW IL /C1;7pyrrþ DIL /C1;6pyrrþ estimated from the plots in Fig. 3 are 15
and 12 mV, respectively. 0 The effect of the elongation of one methylene unit on DW IL /i is 30 mV at 25 °C in the case of 1-alkyl-3-methylimidazolium C1C1N-W two-phase systems [15]. The half of this value, 15 mV, is comparable to the observed difference in the height of the straight lines above and confirms the prediction of Eq. (8). 3.5. Contribution of diffusion potential due to the dissolution of the ILSB constituting ions We examine if the deviation from the straight line at 5 lmol kg1 is ascribed to the diffusion potential created by the dissolution of ILSB-constituent ions in W. To estimate the magnitude of this diffusion potential, we first measured the limiting molar conductivity of C1,npyrrBr and lithium bis(pentafluoroethanesulfonyl)amide at 25 °C. The values obtained are 115, 119 and 122 S cm2 mol 1 for n = 8, 7, and 6, and 67.1 S cm2 mol 1 for lithium bis(pentafluoroethanesulfonyl)amide. From literature values of the limiting molar ionic conductivity of Li+ and Br, we estimated the limiting molar ionic conductivity values as 33(C1,8pyrr+), 38(C1,7pyrr+), 41(C1,6pyrr+) and 28(C2C2N) S cm2 mol1, respectively. Since the mobility of C2C2N is smaller than those of C1,npyrr+, the diffusion potential referred to the inner potential in the bulk phase of Phase V should be negative and its magnitude depends on the concentration of indifferent ions in V [4]. For example, the values of the phase-boundary potential between Phases IV and V when the concentration of MBr in Phase V is 5 lmol kg1 are 2.9 (n = 8), 4.9 (n = 7), and 6.3 (n = 6) mV. As E is measured as /VII /I, this diffusion potential is opposite to the experimentally observed deviation of E at this concentration of MBr in Phase V. According to Eq. (1) and Table 1 the dissolution of AgBr in Phase V is small at cVMBr = 5 lmol kg1 and cannot explain the downward deviation from the straight line. The change in the value of a MBr is too small in the logarithmic scale as discussed above. At this moment, we have no explanation for this discrepancy. The presence of a certain negative bias of about 10 mV in E should be assumed at this concentration. More precise measurements of the molar or molal ion conductivity of relevant ions seem to be required, be left the applicability of the Henderson equation.
+ Fig. 4. Plots of E as a function of the mean activity of MBr, a MBr (M = K (red filled squares) and H+ (open circles)) for ½C1;6 pyrrþ ½C2 C2 N at 25 °C. (For interpretation of the references to color in this figure legend, the reader is referred to the web version of this article.)
3.6. Constancy of the liquid junction potential in a wider range of the MBr concentration The constancy of the DW IL / between the ILSB and an aqueous solution extends to higher concentration range than that shown in Fig. 3. As an example, Fig. 4 shows the plots of E against þ log10 a MBr for ½C1;7 pyrr ½C2 C2 N in the wider concentration range. W The constancy of the DIL / at the ILSB j aqueous solution is maintained up to 0.5 mol kg1 KBr and HBr solutions. In other words, the DW IL / stays constant over four orders of magnitude change in the concentration of MBr in W. At higher concentrations, the upward deviation took place in HBr solutions, as has been seen in the ILSB made of C8mimC1C1N due to the penetration of H+ in the ILSB [2]. 4. Conclusions Three ionic liquids that consist of N-alkyl-N-methylpyrrolidinium bis(pentafluoroethanesulfonyl)amide, where alkyl- is either hexyl-, heptyl- or octyl-, are suitable to an ionic liquid salt bridge for electrochemical measurements, particularly, potentiometry, of low ionic strength aqueous solutions down to the ionic strength of 20 lmol kg1. The remarkable stability of the phase-boundary potential with the standard deviation of 0.4 mV at such low ionic € ckel limiting law is relistrength solutions, where the Debye–Hu able for calculating the activity of ions, opens the way to accurately estimate single ion activities, including pH. Acknowledgements This work was partly supported by Japan Science and Technology Agency under the program, ‘‘Development of Systems and Technology for Advanced Measurement and Analysis’’ and by Grant-in-Aid for Scientific Research (No. 21245021) from the Ministry of Educations, Sports, Science, and Technology, Japan. Support by the Global COE Program, International Center for Integrated Research and Advanced Education in Materials Science (No. B-09) from the Ministry of Education, Culture, Sports, Science and Technology of Japan is highly appreciated. References [1] T. Kakiuchi, N. Tsujioka, S. Kurita, Y. Iwami, Electrochem. Commun. 5 (2003) 159–164.
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