J. Phys. Chem. B 2007, 111, 4901-4909
4901
Vibrational Spectroscopy and Dynamics of Azide Ion in Ionic Liquid and Dimethyl Sulfoxide Water Mixtures† Gerald M. Sando,‡ Kevin Dahl,§ and Jeffrey C. Owrutsky* Code 6111, U.S. NaVal Research Laboratory, Washington, District of Columbia 20375-5342 ReceiVed: October 30, 2006; In Final Form: January 11, 2007
Steady-state and time-resolved infrared spectroscopy of the azide (N3-) anion has been used to characterize aqueous mixtures both with the ionic liquid (IL) 1-butyl-3-methylimidazolium tetrafluoroborate ([BMIM][BF4]) and with dimethyl sulfoxide (DMSO). In the DMSO-water mixtures, two anion vibrational bands are observed for low water mole fractions (0 > Xw > 0.25), which indicates a heterogeneous ion solvation environment. The band at 2000 cm-1 observed for neat DMSO does not shift but decreases in amplitude as the amount of water is increased. Another band appears at slightly higher frequency at low Xw ()0.05). As the amount of water is increased, this band shifts to higher frequency and becomes stronger and is attributed to azide with an increasing degree of hydration. At intermediate and high Xw, a single band is observed that shifts almost linearly with water mole fraction toward the bulk water value. The heterogeneity is evident from the infrared pump-probe studies in which the decay times depend on probe frequency at low mole fraction. For the azide spectra in IL-water mixtures, a single azide band is observed for each mole fraction mixture. The azide band shifts almost linearly with mole fraction, indicating nearly ideal mixing behavior. As with the DMSO-water mixtures, the time-resolved IR decay times are probe-frequency-dependent at low mole fraction, again indicating heterogeneous solvation. In both the DMSO and IL mixtures with water, the relaxation times are slower than would be expected from ideal mixing, suggesting that vibrational relaxation of azide is more sensitive than its vibrational frequency to the solvent structure. The results are discussed in terms of preferential solvation and the degree to which the azide shift and vibrational relaxation depend on the degree of water association in the mixtures.
I. Introduction There is current interest in exploring ionic liquids (ILs) as solvents, because of their potential to be “green” and their inherent tunability.1,2 ILs are molten salts that melt at low temperatures (4 µJ of tunable IR with a spectral width of >150 cm-1 (∼250 fs). Most (90%) of the resulting IR is chopped at 500 Hz and used to pump the sample, while the remaining 10% traverses a delay stage and is used to probe the pump-induced transmission change of the sample. The probe beam is transmitted through wire-grid polarizers before and after the sample, then passes through a monochromator (resolution ∼5 cm-1), and is detected by a HgCdTe detector. The signals are processed with a pair of boxcar integrators and lock-in amplifiers to determine the timedependent absorbance changes [∆A(t) ) log (I0/I(t))]. The second wire-grid polarizer is rotated such that the difference in polarization angles between pump and probe beams is the magic angle, 54.7°, to remove rotational contributions to the transient signal. Time-resolved decay curves collected at the magic angle were fit to the sum of an exponential decay convoluted with a Gaussian pulse. Sometimes an additional instrument-limited Gaussian function was used to account for multiphoton or coherence effects.
Azide Ion in Ionic Liquid and DMSO-Water Mixtures
J. Phys. Chem. B, Vol. 111, No. 18, 2007 4903 TABLE 1: Spectroscopic Parametersa of N3- in DMSO-H2O Mixtures
Xw
υdryb (cm-1)
0.00 0.02 0.05 0.10 0.15 0.20 0.25 0.50 0.75 0.90 0.95 0.98 1.00
2000.2 2000.0 2000.0 1999.8 1999.5 1999.2
υwetb (cm-1)
relaxation timec (ps)
range of relaxation times (ps)
10.7 ( 1.0 2008.4 2008.8 2009.0 2009.4 2010.9 2012.4 2021.8 2033.5 2040.6 2043.7 2045.6 2047.5
9.0 ( 1.0 7.5 6.0 4.7 ( 0.5 4.0 ( 0.4 2.4 ( 0.2 1.4 ( 0.1
5.5-9.2 4.2-7.8
0.9 ( 0.1 0.8 ( 0.1
a
Determined from steady-state FTIR and time-resolved IR pumpIR probe studies. b Uncertainty of (1 cm-1. c Relaxation times for Xw ) 0.0 and 1.0 are those for neat DMSO and neat water, respectively, and agree with those reported in refs 65-67. The relaxation time for Xw ) 0.05 is attributed the dry band, and those for X > 0.15 are attributed to the wet band. A range of relaxation times that depend on the probe frequency is observed for X ) 0.10 and 0.15 as described in the text.
Figure 1. Steady-state FTIR spectra of TBAN3 in various mole fraction mixtures of H2O and DMSO. The solid red lines are the experimental spectra, and the dashed black line is the fit to the function described in the text. The solid blue lines are the individual components for the multiple band fits.
III. Results and Analysis A. Steady-State Spectroscopy. 1. DMSO-Water Mixtures. Steady-state IR spectra for the υ3 band of N3- near 2000 cm-1 for TBAN3 (0.05 M) in mixtures of DMSO-H2O with 0.00 e Xw e 1.00 are presented in Figure 1. Using the TBA rather the Na salt eliminates the ion-pair band that appears at 2024 cm-1 for NaN370 (see Figure S1, Supporting Information) and greatly simplifies fitting of the spectra. In neat DMSO, the strongest band is the fundamental antisymmetric stretching (υ3) band at 2000 cm-1. The weak band that appears as a shoulder on the low-frequency side of the fundamental is the bending hot band72 at ∼1987 cm-1, due to thermal population of the bending state (υ2), that is, the υ3 + υ2 - υ2 band. The anharmonicity, the shift of the hot band relative to the fundamental band, was measured in the gas phase to be 14.1 cm-1,72 and the position observed in our spectra is consistent with this value. (The bending frequency73 is about 645 cm-1.) At the other end of the mole fraction range is the TBAN3 spectrum in neat water, in which the azide band is centered at 2047.5 cm-1 with a 25.6 cm-1 width. The band for TBAN3 in water appears at slightly lower frequency than for NaN3 in water (2049 cm-1).74
When water is added to the neat DMSO solution (i.e., for Xw ) 0.02 and 0.05), a new, partially resolved band appears slightly to the high-frequency side of the band originally observed in DMSO. As Xw is increased, the frequency of the DMSO band does not shift much but its amplitude decreases, while the new band grows in and shifts blue. In the Xw ) 0.10 spectrum, the new band has shifted to higher frequency and appears to have an amplitude comparable to that of the DMSO band. This new band that appears when water is added and gets stronger and shifts blue with increasing water concentration will be referred to as the “wet band” in order to differentiate it from the weakly shifting band near 2000 cm-1 (hereafter called the “dry band”). When enough water is added to reach Xw ) 0.25, only a single band is observed at 2012.4 cm-1. The band continues to shift blue with increasing Xw until the bulk H2O value is reached. The steady-state FTIR spectra for TBAN3 in DMSO-water mixtures shown in Figure 1 were fit using the appropriate number of mixed Lorentzian/Gaussian functions (as in the GRAMS32 software), and the results are presented in Table 1. The neat DMSO spectrum was fit to two bands, for the fundamental and hot bands. The low mole fraction spectra (Xw ) 0.02, 0.05, and 0.10) were fit to three bands, the wet, dry, and hot bands. For Xw ) 0.15 and 0.20, the fit included the wet and dry bands, and for Xw ) 0.25 and higher mole fractions, a single band provided an acceptable fit. The individual band contributions to the spectra are shown for the lower mole fractions in Figure 1. The entire set of parameters determined from the analysis is available as Supporting Information. The frequencies of the wet and dry bands determined from the fitting are shown as a function of mole fraction in Figure 2. An initial series of spectra for NaN3 in DMSO-water mixtures was obtained in which the ion-pair band is observed. Spectra were also measured for NaN3 in DMSO-D2O and DMF-H2O mixtures, as well as for TBAN3 in DMF-H2O. These spectra are qualitatively similar to those presented here for TBAN3 in DMSO-H2O and are shown and described in the Supporting Information. 2. IL-Water Mixtures. Steady-state IR spectra for the ν3 band of N3- for TBAN3 (0.03 M) in mixtures of [BMIM][BF4] and
4904 J. Phys. Chem. B, Vol. 111, No. 18, 2007
Figure 2. Vibrational frequencies of the “wet” (3) and “dry” (b) bands of the antisymmetric stretch of TBAN3 in mixtures of H2O and DMSO as a function of mole fraction.
Sando et al.
Figure 4. Vibrational frequency of the antisymmetric stretch of azide as a function of water mole fraction in mixtures of H2O and [BMIM][BF4].
TABLE 2: Spectroscopic Parametersa of N3- in [BMIM][BF4]-H2O Mixtures
Xw
υmaxb (cm-1)
FWHMb (cm-1)
relaxation timec (ps)
0.00 0.05 0.1 0.25 0.5 0.75 1
2008.3 2011.0 2014.5 2020.2 2027.0 2034.5 2047.1
18.0 21.8 24.7 27.0 28.3 28.5 25.6
9.1 ( 0.8 7.1 5.1 3.4 1.7 ( 0.1 1.2 ( 0.1 0.8 ( 0.1
range of relaxation times (ps) 5.8-8.2 4.0-6.1 2.7-4.0
a Determined from steady-state FTIR and time-resolved IR pumpIR probe studies. b Uncertainty of (1 cm-1. c Relaxation times for Xw ) 0.0 and 1.0 are those for neat [BMIM][BF4] and neat water, respectively, and agree with those reported in refs 65-67. Relaxation times for 0.05 e Xw e 0.25 exhibit relaxation times that depend on the probe frequency, as indicated in the last column and as described in the text.
Figure 3. Steady-state FTIR spectra of TBAN3 in mixtures of H2O and [BMIM][BF4] for several mole fractions. The solid red line is the experimental spectrum, and the dashed black line is the fit to the Gaussian/Lorenztian function described in the text.
H2O are shown in Figure 3 as a function of Xw. In neat [BMIM][BF4],65 the antisymmetric stretch appears as a single band near 2008 cm-1. The band is broader and less shifted from water in the IL than in DMSO. Unlike in DMSO, no hot band is evident in the spectrum. Upon addition of water to the neat IL solution, the antisymmetric stretch frequency undergoes a blue shift that increases with increasing water concentration and approaches the frequency in neat water (2047.1 cm-1). Only a single band is apparent for all mole fractions. The FTIR spectra were fit to a single Gaussian/Lorenztian band, and the results are shown in Figure 3. The single band provides an acceptable fit in all cases. However, small deviations are apparent at low mole fractions (Xw ) 0.05 and 0.1). Attempts to account for these deviations with an additional band were not successful. The deviations indicate that the band has some asymmetry for these mole fractions and suggest a possible additional unresolved band. The mole fraction dependence of the vibrational frequency as determined from the spectral fitting is shown in Figure 4
and in Table 2. The entire set of parameters is included as Supporting Information. Spectra taken with NaN3 are nearly indistinguishable from those with TBAN3; however, use of the TBA salt improves the solubility in the neat ionic liquid. B. Vibrational Dynamics. 1. DMSO-Water Mixtures. Transient IR spectroscopy was used to measure VER times for azide in DMSO-H2O mixtures over the entire range of mole fractions. The results for the neat solutions of DMSO (10.7 ( 1 ps) and water (0.8 ( 0.1 ps) agree with the values we have reported previously.65-67 Figure 5 shows representative magicangle population decay curves (i.e., transient bleach and transient absorption pairs) and fits to the decays for N3- in Xw ) 0.25 (T1 ) 4.0 ( 0.4 ps) and Xw ) 0.75 DMSO-H2O mixtures (T1 ) 1.4 ( 0.1 ps). For mixtures with Xw g 0.20, the observed decay times were the same for the excited-state decay and ground-state bleach recovery times within experimental error, and did not depend on the probe frequency. The results for VER (T1) times are given for neat DMSO, H2O, and for several intermediate Xw values in Table 1. For the low Xw mixtures, especially 0.10 and 0.15, there was a marked dependence of the observed decay times on the probe frequency. In the static spectra for these mole fractions, the wet and dry bands are separated by less than the bandwidth, so it is difficult with our probe frequency resolution (of about 5 cm-1) to clearly resolve the T1 times for each of the bands separately. The vibrational dynamics were characterized by concentrating on the transient bleach region on the blue side of the transient signal, corresponding to the frequencies for the static absorption bands and the transient bleaches. In this region, the relative positions of the dry and wet bands are most reliable since they
Azide Ion in Ionic Liquid and DMSO-Water Mixtures
Figure 5. Representative time-resolved IR pump-IR probe transients (polarizer at 54.7°) for N3- in mixtures of DMSO/H2O.Open circles represent population dynamics of the transient bleach, gray circles represent population dynamics of the transient absorption, and solid lines represent fits to the decay data. (a) N3- in Xw ) 0.25 DMSO/H2O, absorption probe frequency ) 1989 cm-1, bleach probe frequency ) 2031 cm-1. (b) N3- in Xw ) 0.75 DMSO/H2O, absorption probe frequency ) 2008 cm-1, bleach probe frequency ) 2046 cm-1.
are determined in the static spectra. The anharmonicities and excited-state band widths are hard to know a priori, so it is difficult to predict or determine the relative contributions from the wet and dry bands on the long-wavelength side of the transient spectrum where the excited-state absorptions appear. The times were consistently longer for lower probe frequencies for a given Xw, which is consistent with the expectation that the transients observed with a lower-frequency probe correspond to a greater relative contribution from the dry band. For Xw ) 0.10, the decay times varied between 5.5 and 9.2 ps for probe frequencies in the range 2015-2000 cm-1. The band-averaged estimate, that is, including contributions from the dry and wet bands, is 7.5 ps. For Xw ) 0.15, times in the range 4.2-7.8 ps were observed over the same frequency interval, and the average time in this case is estimated to be 6.0 ps. The VER times for the intermediate Xw solutions vary monotonically from that for neat DMSO (10.7 ps) to that for neat H2O (0.8 ps) over the range of DMSO-H2O mixtures. The band-averaged VER rates are shown in Figure 6 along with the composite averaged vibrational frequency. 2. IL-Water Mixtures. Transient IR spectroscopy was also used to measure VER times in the [BMIM][BF4] mixtures. Figure 7 shows representative scans at the magic angle and single-exponential fits used to determine the VER times. The VER times were measured for the entire range of mole fractions. For neat IL and for Xw g 0.5, fits of the transient bleach and transient absorption decays agreed within experimental uncertainty and showed no dependence on the probe frequency. In the neat IL, the measured T1 time was 9.1 ( 0.8 ps, which is
J. Phys. Chem. B, Vol. 111, No. 18, 2007 4905
Figure 6. Vibrational energy relaxation (VER) rates (a) and vibrational frequencies (b) for azide in mixtures of H2O and DMSO as a function of water mole fraction. For Xw ) 0.10 and 0.15, the VER rates are averaged over the probe-frequency-dependent results. At low mole fraction (Xw ) 0.02-0.20) the frequencies represent a weighted average of the wet and dry bands (see text).
Figure 7. Decays of the transient bleach of azide at different probe frequencies for Xw ) 0.10 mixtures of H2O and [BMIM][BF4]. The probe frequencies are 2010 cm-1 (b) and 2025 cm-1 (4). The solid lines are fits to single-exponential decays with decay times of 6.5 and 4.0 ps, respectively.
within the uncertainty of our previously reported time (9.8 ( 0.8 ps).65 The lower mole fraction mixtures clearly exhibit decays that depend on the probe frequency. This is evident in Figure 7, where the two decay curves are for the same sample at different probe frequencies. The transient bleaches of the low mole fraction samples exhibit faster decays for higher probe frequencies. The transient absorption decays, however, show little frequency dependence and decay with a time constant near that of the longest decay time for the transient bleach. The decays were well fit by a single exponential in all cases. The results for the VER times are listed in Table 2. For the frequency-dependent samples, both the range of measured T1 times and an average T1 time are listed. The average T1 time was determined by averaging the T1 times at all the probe frequencies for that mole fraction. For Xw ) 0.05, the decay times varied from 5.8 to 8.2 ps for bleach probe frequencies in the range of 2010-2025 cm-1, while the transient absorption decays were slightly slower. For Xw ) 0.1 in the same frequency range, the transient bleach decays times ranged from 4.0 to
4906 J. Phys. Chem. B, Vol. 111, No. 18, 2007
Figure 8. VER rate of azide as a function of water mole fraction for mixtures of H2O and [BMIM][BF4]. The rate is averaged over the probe frequencies for Xw ) 0.05-0.25, where the VER rate depends on the frequency of the probe.
6.1 ps, while the transient absorption decays time were near 6 ps. For Xw ) 0.25, the transient bleach decays varied from 2.7 to 4.0 ps in the 2015-2030 cm-1 range, while the transient absorption decays ranged from 3.1 to 4.1 ps in the 1980-1995 cm-1 frequency range. The uncertainties are larger for smaller Xw, partly due to small errors in the mole fraction that result from adding a small amount of water to a much larger volume of ionic liquid and any residual water in the neat IL. Small errors in the mole fraction are more important in the low mole fraction regime. The mole fraction dependence of the averaged VER rates is shown in Figure 8. The average VER rates increase monotonically toward that of neat H2O as Xw increases. The same overall trend is seen when the range of T1 times for the frequency-dependent samples is included, even though there is some overlap between different mole fractions. IV. Discussion A. Overview. ILs are being studied by various methods to determine if and how they differ from molecular solvents. Fundamental questions are being explored, including how polar,75-77 how ionic,78 or even how liquid they appear and whether they are well described by models for conventional liquids. For example, even though they are ionic, they have relatively low refractive indices.10,76 They may or may not be miscible with water, depending on the anion, which also mediates whether water in ILs is weakly associating or has substantial hydrogen bonding.9,10,59 It is interesting that in ILs the water bands do not shift as the mole fraction is increased, and the water vibrational frequencies compared to those in the gas phase provide a convenient way to measure the IL refractive index.10 The solvation properties of ILs, such as polarity or Kamlet-Taft parameters from solvatochromatic studies,6,12,79,80 have been described as somewhere between those of acetonitrile and short-chain alcohols.77,79 Expanding the variety of techniques employed provides a more comprehensive evaluation of ILs. Most spectroscopic studies have used solvation dynamics and solvatochromatic studies. The work presented here provides a complementary perspective of an anionic solute probe by use of vibrational spectroscopy. While there are several studies that point to heterogeneity of neat ILs,56-58 we find no evidence of it from the azide vibrational studies until water is added. Previous studies have demonstrated that azide is effective as an anionic solute probe for investigating solvation in bulk solvents65 and more complex systems.71,81 The present experiments combine steady-state vibrational spectroscopy and timeresolved dynamics to probe the azide solvation in aqueous DMSO and [BMIM][BF4] mixtures. We were initially interested in exploring the effects of water contamination on our results
Sando et al. for studies of azide ion pairs in DMSO with Na+, Ca2+, and Mg2+.66 There is a much higher vibrational frequency and faster VER dynamics of azide in water compared to DMSO. Also, ion pairs of azide with alkali and alkali earth metal cations are observed in DMSO, indicating weaker solvation than in water. We therefore expected a stronger affinity of water for azide might result in its preferentially solvating the anion in aqueous DMSO mixtures. Similarly, the vibrational spectra and dynamics for azide in ILs are similar to those for DMSO, so the same preference for water might be expected in IL-water mixtures. On the other hand, results from our previous studies show that small amounts of water do not have a strong effect on the vibrational spectra and dynamics of anions in ILs.65 The vibrational shifts and T1 times in ILs that are difficult to dry, such as for NCS- in [BM2IM][NCS] (with 1.5 wt % water or Xw ) 0.15) and N(CN)2- in [BM2IM] [N(CN)2] (with 0.05 wt % water), were very similar to those with the drier [BM2IM][BF4] IL. B. Inhomogeneity from the Vibrational Spectra. The results of this study indicate that both IL and DMSO mixtures with water are inhomogeneous for low mole fractions. The result is most obvious for DMSO since two bands are resolved in the spectra. The band originally observed in neat DMSO close to 2000 cm-1 does not shift appreciably upon adding water. The wet band appears to the blue side of the DMSO frequency when water is added, and as the amount of water is increased, it shifts blue and gains strength at the expense of the wet band. This indicates that as more water is added, a higher fraction of the anions have water penetrating into the first solvent shell. The mixed environment is also indicated by the probe frequency dependence observed for vibrational dynamics in the DMSO as well as in the IL-water mixtures. The results for the DMSO-water mixtures reinforce our interpretation that the probe frequency dependence observed in the IL-water mixtures is a signature of mixed solvation environments in that case also. The mixed environment is most obvious in the azide spectra of low Xw DMSO-water mixtures. The IL-water mixture spectra consist of single bands for all mole fractions. It is more difficult to resolve two bands in this case because the bandwidth is broader for azide in the IL than in DMSO and because there is less shift between the neat IL and water bands. The spectra in the IL-water mixtures are slightly asymmetric, but not enough to identify distinct solvation environments. Therefore, evidence of heterogeneity in these solutions depends solely (or at least primarily) on the probe frequency dependence of the vibrational dynamics, which is reinforced by the similar observation in the DMSO mixtures. In both the DMSO and IL mixtures, for a particular sample (mole fraction), probing with higher frequencies results in shorter relaxation times. This correlation between spectral frequency and relaxation rate is similar to the trend for different mole fractions as well as the trend for the bulk solvent dependence for azide. The inhomogeneity observed in the IL-water mixtures at low Xw suggests that these mixtures deviate further from ideal mixing than water-rich solutions. This observation agrees with studies of solvation parameters in [BMIM][BF4]-water mixtures in which the solvation excess, the fractional deviation from ideal behavior, was found to be larger for small Xw than for the waterrich mixtures.13 Also, there have been numerous reports of heterogeneity in neat IL solutions, from Raman, fluorescence, and OKE experiments as well as numerous MD simulations.56-58,60-64 Most of these are attributed to the nonpolar regions of the solutions, which will not be probed the by
Azide Ion in Ionic Liquid and DMSO-Water Mixtures hydrophilic and polar azide solute used in this study. Furthermore, there may be mixed polar regions of the IL, but it appears, if there are, that they are not populated by azide or that the inhomogeneity from adding water is greater in the anionsolvating regions than for the neat ILs. C. Preferential Solvation. The azide spectra can be evaluated for evidence of preferential solvation. One way is to consider how the azide vibrational frequency varies with Xw. Preferential solvation is typically considered for homogeneous solutions, in which a single value for the property is observed and nonideal behavior is indicated when the property deviates from the molefraction-weighted average of the values for the neat components. For this purpose, an average azide frequency of the dry and wet bands can be determined by an integrated band intensityweighted average azide frequency. If the bands (at low mole fraction) were not resolved and a single unresolved band were observed, it would be centered at the calculated composite frequency, υc ) FAw υw + FAdυd, where FAw and FAd are the fractional integrated band intensities for the wet and dry bands, respectively, and υw and υd are the azide center frequencies for the wet and dry bands, respectively. A plot of the composite frequency as a function of Xw is shown in Figure 6b. The plot indicates that the azide frequency exhibits a mole fraction dependence that is close to that expected for ideal mixing. There is a sigmoidal dependence in which there is a small preference toward the water value for low Xw and then a shift to a slight DMSO preference for higher mole fractions. This behavior is similar to what has been observed for the solvatochromatic behavior of an organic dye.31 It indicates that, based on the azide vibrational frequency, there is not much preferential solvation for water in DMSO-water mixtures. The mole fraction dependence of the azide band frequency observed in the ILwater mixtures (Figure 4) looks almost the same as for the composite frequency of the DMSO-water mixtures. The DMSO-water mixtures can also be analyzed for preferential solvation by use of the (integrated) band intensities of the wet and dry bands. The integrated intensities of the bands reflect the anion populations with and without water, respectively, in the solvent shell, which can be compared to statistical expectations. For example, for Xw ) 0.05, the integrated intensities of the wet and dry bands are the same within 10%. If we assume that there are 5-6 water molecules in a full first shell (Nw) of azide in neat water, as predicted by MD simulations by Klein and co-workers,82 then the statistically anticipated fraction of azides with water in the solvent shell can be estimated as NwXw ) 6 × 0.05 ) 0.3. This suggests that at Xw ) 0.05, the fraction of azide ions with water in the solvent shell (∼0.5) is slightly higher than the statistical amount, so there is a small degree of preference for water at this mole fraction. This is consistent with the result based on the mole fraction dependence of the frequency at low Xw, in which a small preference for water was identified. We have determined how the azide frequency varies with water mole fraction in both mixtures, but it is not clear how the frequency will depend on the local concentration of water. If it were possible to assign a specific solvation structure, such as the number of water molecules in the solvent shell, based on the frequency, we could attribute those structures to the mole fraction solutions. The lower frequencies and slow rates for azide in DMSO and ILs imply that these properties are more strongly affected by water solvation, so the shift may reflect the number of waters associated with the anion. Pertinent to this are density functional theory (DFT) calculations of N3- solvated in (H2O)n clusters (n ) 0-5) reported by Yang et al.83 in connection with
J. Phys. Chem. B, Vol. 111, No. 18, 2007 4907 their studies on the effect of hydration on electron binding energy of hydrated azide clusters (for n ) 0-16) by use of photoelectron spectroscopy. They show that for a few waters (2-5) associated with azide, there are several isomers that are similar in energy. The azide frequencies for the isomers for a given number of waters varies enough to undermine the notion of simply correlating the frequency to the degree of hydration.84 This approach might be suspicious anyway, since the behavior of an isolated, gas-phase cluster would probably be quite different than a few waters penetrating the solvent shell of a DMSO- or IL-solvated anion. (Spectra of neutral clusters tend to be rotationally resolved, indicating they have much longer lifetimes than solution species.) D. Vibrational Relaxation. The mole fraction dependence of the vibrational relaxation is different than for the vibrational frequencies, which probably reflects that these properties depend differently on details of the solvent structure, more than being an indication of how DMSO and IL differ. The VER rates for azide as a function of Xw in the DMSO and IL mixtures are shown in Figures 6a and 8, respectively. In both cases, the rates increase monotonically upon adding water to both of the nonaqueous solvents. However, the rates increase more slowly than would be expected for ideal mixing; that is, the observed rates are slower than expected for the ideal mixing line and demonstrate a solvent preference for the nonaqueous solvents. There is some ambiguity in the low mole fraction results because they represent an average of the probe-frequency-dependent values. The result is probably better demonstrated by the results for the intermediate Xw mixtures, which have single azide bands. The rates for Xw ) 0.25, 0.50, and 0.75 are clearly slower than the ideal mixing line for both DMSO (Figure 6a) and [BMIM][BF4] (Figure 8). There is a preference for the nonaqueous solvent in the vibrational relaxation times. At higher Xw, the shift and VER rate increase almost linearly with water content, suggesting little or no preferential water solvation. E. Comparison between Spectra and Dynamics. It is surprising that the spectra and dynamics for azide behave differently in the mixtures. There are several solvatochromatic studies of aqueous DMSO mixtures,16-19 and as for bulk solvents, the polarity or solvation behavior depends on the dye used to measure it. The interactions with various solutes are the basis for the large number of polarity scales.85 But in the present case, different results with respect to the mole fraction dependence are observed for different properties of the same solute probe species, azide. The vibrational frequencies and relaxation rates of azide are generally correlated in bulk solvents, unlike solvatochromatic shifts and solvation dynamics of dyes. This demonstrates that the vibrational spectral shifts and relaxation times are sensitive to different characteristics of the solvation environment. A possible explanation is that the relaxation mechanism depends more on the degree of association and extent of hydrogen bonding, which may not develop until more water is present in the higher water mole fraction mixtures. For the azide vibrational relaxation of the antisymmetric stretch, there are predictions that intramolecular vibration relaxation (IVR) mechanisms via the symmetric stretching mode contribute to the azide VER, which could be solvent-assisted, such as by the librational mode of water.86,87 If so, the rate may increase nonlinearly with mole fraction. A similar result was reported in studies of the vibrational relaxation dynamics of the OH stretching bands of HOD in its mixtures with acetonitrile by Cringus et al.88 They demonstrated that the relaxation rate increased exponentially with the HOD mole fraction and that the mechanism changed from dissipation to the solvent, at low
4908 J. Phys. Chem. B, Vol. 111, No. 18, 2007 HOD content, to solvent-assisted intramolecular vibration relaxation (IVR) to the bend overtone as the degree of water association increased. The situation is somewhat different in our case; we are monitoring a solute in the mixture rather than one of the major solvent components, and the inhomogeneity in HOD results in the spectra being broader. But there may be a similar effect of the solvent modes contributing to the relaxation, which results in slower than linear onset of the relaxation rate due to the libration bands developing nonlinearly with water content. There is a puzzling aspect of this notion that the degree of water association in the mixtures mediates the relaxation rates. The water in [BMIM][BF4] appears to be less associated than in DMSO even though the VER rates for a given mole fraction are slower in the latter. For example, the relaxation time for Xw ) 0.50 is 1.7 ( 0.1 ps for [BMIM][BF4] and 2.4 ( 0.2 ps for DMSO. But the water is more associated in the DMSO mixtures, as indicated by the broader and more red-shifted IR spectra of the water OH bands compared to those in IL-water mixtures, where the latter suggest that the water is not strongly associated.9,10,59 This may result from the stronger bond between water and DMSO than between two water molecules,19,39,89 which could prevent water from solvating the anion as well as in the IL mixtures. This explanation is not supported by the similarity between the azide frequency dependence on mole fraction for the DMSO- and IL-water mixtures. It would be instructive if MD simulations could reproduce the results we have observed in terms of the azide spectra and dynamics to assist us in explaining them. Also, the results, especially the rates, indicate a solvent preference for the nonaqueous solvents in both cases, which is opposite to our expectations that water would have a strong affinity to solvate azide. For the DMSO mixtures, it may be a consequence of the strong interaction between water and DMSO. There are several theoretical studies on ion solvation in DMSOwater mixtures. These demonstrate that ions significantly impact the solvent properties and that the complex interplay between solvent-solvent and solute-solvent interactions results in complicated behavior. As with some of the early experimental studies on preferential ion solvation,33,34 no consistent trends emerge from the calculations regarding the importance of charge, polarity, size, and shape to identify ion-solvent affinities. Molecular dynamics simulations of single-particle ions in DMSO-H2O mixtures by Laria and Skaf41 have predicted that anions will be preferentially solvated by H2O and that cations will be solvated by DMSO. Day and Patey40 and Chowdhuri and Chandra,42 in similar MD simulations, found that H2O preferentially solvates both anions and cations. Simulations by Das and Tembe49 have found, at least for mixtures with Xw e 0.10, that H2O preferentially solvates the Na+ cation in the presence of Cl- anion. Our results suggest that water is not strongly preferred in either DMSO or IL mixtures with water. It would be interesting to see if these similarities are uncovered in MD simulations that model these systems and determine the effect of ions not only on the solvent properties but also on the properties of the ions themselves. V. Summary and Conclusions We have characterized water mixtures of [BMIM][BF4] and DMSO using vibrational spectroscopy and dynamics of the antisymmetric stretching band of azide ion. This expands on our earlier studies of anion spectra and vibrational dynamics in neat ILs65 and IL reverse micelles.90 In the DMSO-water mixtures, two anion vibrational bands are observed for small
Sando et al. amounts of water (0 > Xw > 0.25), which indicates a heterogeneous ion solvation environment. The heterogeneity is evident from the infrared pump-probe studies in which the decay times depend on probe frequency at low mole fraction. In the azide spectra in IL-water mixtures, a single azide band is observed for each mole fraction mixture. The azide band shifts almost linearly with mole fraction, indicating nearly ideal mixing behavior. As with the DMSO-water mixtures, the time-resolved decay times are probe-frequency-dependent at low mole fraction, again indicating heterogeneous solvation. In both the DMSO and IL mixtures with water, the relaxation times are slower than would be expected from ideal mixing, suggesting that vibrational relaxation is more sensitive to the solvent structure than the vibrational frequency. Although there are several studies, both experimental and theoretical, that indicate inhomogeneity in neat ILs, the polar regions in which anions are solvated become more heterogeneous when water is added. From the perspective of small anion vibrational spectra and dynamics, the differences between molecular and ionic liquids are less than those between either of the neat solvents and the corresponding mixture with water. Acknowledgment. Support for this work was provided by the Office of Naval Research through the Naval Research Laboratory. K.D. acknowledges the Naval Research Laboratory-National Research Council research associateship. G.M.S. acknowledges the Naval Research Laboratory-American Society for Engineering Education postdoctoral fellowship program. We gratefully acknowledge Qun Zhong for early stages of data collection and Andrew Baronavski for assistance with the laser system. Supporting Information Available: Steady-state IR spectra as a function of water mole fraction for NaN3 in DMSO-H2O, DMSO-D2O, and DMF-H2O mixtures (Figure S1); IR spectra of TBAN3 in DMF-H2O mixtures (Figure S2); and parameters determined from fitting the IR spectra shown in Figure 1 (Table S1) and Figure 3 (Table S2). This material is available free of charge via the Internet at http://pubs.acs.org. References and Notes (1) Welton, T. Chem. ReV. 1999, 99, 2071. (2) Seddon, K. R.; Stark, A.; Torres, M. J. Pure Appl. Chem. 2000, 72, 2275. (3) Earle, M. J.; Esperanca, J.; Gilea, M. A.; Lopes, J. N. C.; Rebelo, L. P. N.; Magee, J. W.; Seddon, K. R.; Widegren, J. A. Nature 2006, 439, 831. (4) Pretti, C.; Chiappe, C.; Pieraccini, D.; Gregori, M.; Abramo, F.; Monni, G.; Intorre, L. Green Chem. 2006, 8, 238. (5) Fletcher, K. A.; Pandey, S. J. Phys. Chem. B 2003, 107, 13532. (6) Fletcher, K. A.; Baker, S. N.; Baker, G. A.; Pandey, S. New J. Chem. 2003, 27, 1706. (7) Schroder, U.; Wadhawan, J. D.; Compton, R. G.; Marken, F.; Suarez, P. A. Z.; Consorti, C. S.; de Souza, R. F.; Dupont, J. New J. Chem. 2000, 24, 1009. (8) Fitchett, B. D.; Rollins, J. B.; Conboy, J. C. J. Electrochem. Soc. 2005, 152, E251. (9) Cammarata, L.; Kazarian, S. G.; Salter, P. A.; Welton, T. Phys. Chem. Chem. Phys. 2001, 3, 5192. (10) Koddermann, T.; Wertz, C.; Heintz, A.; Ludwig, R. Angew. Chem., Int. Ed. 2006, 45, 3697. (11) Tran, C. D.; Lacerda, S. H. D.; Oliveira, D. Appl. Spectrosc. 2003, 57, 152. (12) Fletcher, K. A.; Pandey, S. Appl. Spectrosc. 2002, 56, 266. (13) Harifi-Mood, A. R.; Habibi-Yangjeh, A.; Gholami, M. R. J. Phys. Chem. B 2006, 110, 7073. (14) Bosch, E.; Roses, M. J. Chem. Soc., Faraday Trans. 1992, 88, 3541. (15) Catalan, J.; Diaz, C.; Garcia-Blanco, F. J. Org. Chem. 2001, 66, 5846. (16) Higashigaki, Y.; Christensen, D. H.; Wang, C. H. J. Phys. Chem. 1981, 85, 2531.
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