Ions and Hydrogen Bonding in a Hydrophobic Environment: CCl4

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J. Phys. Chem. A 2010, 114, 4051–4057

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Ions and Hydrogen Bonding in a Hydrophobic Environment: CCl4 Patrick Bisson, Han Xiao, Margaret Kuo, Noe Kamelamela, and Mary Jane Shultz* Pearson Lab, Tufts UniVersity, Medford, Massachusetts 02155 ReceiVed: NoVember 9, 2009; ReVised Manuscript ReceiVed: February 16, 2010

It is generally expected that ions in an aqueous ionic solution in contact with a hydrophobic phase enter the hydrophobic phase accompanied by a hydration shell. This expectation suggests that the ion mole fraction in the hydrophobic phase is less than, or at most, equal to that of water. Both gravimetric and spectroscopic evidence shows that for a model hydrophobic phase, carbon tetrachloride, this is not the case: In contact with a 1 M simple salt solution (sodium or potassium halide), the salt concentration in carbon tetrachloride ranges from 1.4 to nearly 3 times that of water. Infrared spectra of the OH stretch region support a model in which water associates with the cation, primarily as water monomers. Salts containing larger, more polarizable anions can form outer-sphere ion pairs that support water dimers, giving rise to a spectral signature at 3440 cm-1. In CCl4, the infrared spectral signature of the normally strongly ionized acid HCl clearly shows the presence of molecular HCl. Additionally, the presence of a Q branch for HCl indicates restricted rotational motion. The spectral and gravimetric data provide compelling evidence for ion clusters in the hydrophobic phase, which is a result that may have implications for hydrophobic matter in both biological and environmental systems. Introduction Aqueous solutions containing inorganic ions are prevalent in both environmental and biological systems: Sea water is 3.5% saline, corresponding to 35 g of salt per kg of water, and blood has a salt content of about 9 g per kg of blood (mainly water). Further, these relatively concentrated salt solutions are often in contact with a hydrophobic phase. How does the presence of salt modulate the concentration or dispersion of water in the hydrophobic material? The famous Hofmeister series indicates that water interacts with hydrophobic materials and that salts do modulate the interaction, although the mechanism of the modulation is less well understood.1 The current conjecture seems to be that salts alter the structure of water, thus modifying the interaction between water and hydrophobic groups. In this work, carbon tetrachloride is a model for the hydrophobic portion of biological material. Salts are found to alter water-water interactions. Similarly, the air side of a water-air interface is often described as a hydrophobic phase. Particularly for environmental chemistry, the issue of how the addition of salt alters the structure and dynamics of interfacial water is important for determining the fate of substances released into the atmosphere.2 Accordingly, there have been several recent studies directed at determining both the ion distribution3-9 and changes in the structure of water at the air-water interface for aqueous salt solutions.8,10-16 Specifically, sum frequency generation (SFG) studies11,17 have shown that addition of soluble inorganic ions significantly alters the vibrational spectrum of the water surface. Interpretation of the SFG data, however, is hampered by lack of a connection between the broad hydrogen-bonded resonances and the structure of water. Particularly germane to this paper, the addition of halide ions enhances the SFG resonance at 3440 cm-1; more polarizable ions, for example, I- vs Cl-, give a * To whom correspondence should be addressed. Phone: 617-627-3477. Fax: 617-627-3441. E-mail: [email protected].

greater enhancement of the 3440 cm-1 resonance. The mechanism for this enhancement, that is, how the structure of water is altered, is not well understood. The present paper provides insight into the origin of the 3440 cm-1 resonance. The developing molecular-level picture of the interface comes primarily from molecular dynamics (MD) simulations.18-20 MD simulations indicate that the ion distribution at the air-water interface is neither monotonic nor the same for anions and cations. This ion distribution picture supports the interpretation that the ion distribution is likely responsible for the differences among ions observed in the Hofmeister effect as well as changes in the water structure observed with surface spectroscopy. The present room-temperature, matrix-isolation FTIR results show differences among sodium and potassium halide salts, providing a model system for determining synergistic anion-cation effects on the water structure. Focusing on transport across the water-hydrophobic phase interface, simulations21,22 provide indirect evidence that small ions remain hydrated on transfer from water to an organic phase. Specifically, I- is less solvated with a less intact solvation shell than Cl- across the water-carbon tetrachloride interface.22 Unlike calculations, experimental systems contain both anions and cations. The results of the present study suggest that the same strong interactions that keep the solvation shell around a small ion in the simulation systems also produce ion clusters in the experimental systems. There is precedence for clustering even in aqueous solutions with an association constant that appears to depend on the cation (ion association constant, KA ) 0.60 for NaCl, 0.61 for NaBr, 0.52 for KCl, and 0.53 for KBr23). The present matrix-isolation FTIR results indicate that ions exist in CCl4 with some associated water, but the water: ion-pair ratio is less than one. The combination of a large, polarizable anion and a small cation is required to form water dimers such as would be present in an intact hydration shell. The salt concentration investigated, 1 M, is comparable to the salt concentration found in the ocean (1 M is ∼1.5 times that of seawater) and is about six times that of blood. Thus, the

10.1021/jp9106712  2010 American Chemical Society Published on Web 03/03/2010

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results of this study are directly relevant to aqueous-hydrophobic partitioning of salt and water in environmental and biological systems. This paper is organized as follows: After a description of the experimental procedure, results are presented for sodium chloride salt solution followed by those for NaBr, NaI, KBr, and KI. The variation in the spectra among these salts provides insight into interactions between water and the anions and cations of these salts. The water-salt spectra are contrasted with that of aqueous HCl. The ion concentration in the hydrophobic layer for all solutions is quantified with gravimetric analysis and these data are presented. An infrared spectrum of the carbon tetrachloride layer after gravimetric analysis provides further insight into the interactions between water and ions in the hydrophobic layer. The last section contains the conclusions. Experimental Section Sample preparation has been described previously,24 so only major features are cited here. Carbon tetrachloride (Sigma-Aldrich, > 99.5%, anhydrous) is removed from the stock bottle with a syringe, injected into a septum-cap sealed reservoir and kept over activated charcoal to remove hydrocarbon impurities. For the studies reported here, carbon tetrachloride was not dried further. Water used to prepare the salt solutions is 18 MΩ (Barnstead) and UV irradiated to remove organic contaminants. Removal of organic contaminants is confirmed by SFG spectra of the solution interface; upper limit to organic contamination is nanomolar. Salt solutions were made by preparing a 1 M stock aqueoussalt solution, shaking the stock solution with carbon tetrachloride in a sealed, silanized vessel and letting the resulting mixture phase separate until a stable FTIR spectrum is obtained: about a day. The resulting carbon tetrachloride with salt-water solution (or HCl-water) was stable (as reflected by reproducible IR absorption spectra) for weeks. The carbon tetrachloride layer is extracted into a silanized infrared cell equipped with silanized, IR quartz windows. Infrared spectra of the carbon tetrachloride phase at room temperature are obtained with a Nicolet Magna-IR 760 FTIR (32 scans, 1 cm-1 resolution) with a DTGS KBr detector. Samples for gravimetric analysis were similarly prepared, but placed in a sealed silanized separatory funnel. The first 20 mL CCl4 solution were drawn off and discarded; a 50 mL middle fraction was collected for analysis. The analyzed fraction was mixed with 50 mL of 5 M AgNO3/ 0.044 M HNO3, kept in the dark for 2 days, and cooled overnight to ensure complete precipitation. The resulting precipitate was collected in a funnel with a medium frit disk, washed with 0.042 M HNO3, and dried in a desiccator for 3 days until constant mass. Results The aim of the present study is to probe the effect of salt on partitioning of water and salt into a model hydrophobic solvent, carbon tetrachloride. The surprising result is that the hydrophobic phase supports a significant salt concentration; indeed, the salt concentration is greater than the water concentration. Additionally, relative to the neat water-carbon tetrachloride two-phase system, salts increase the water concentration and support additional water configurations. Chloride Salts: NaCl. Earlier work24 characterizing water in carbon tetrachloride determined that the saturated concentration of water in carbon tetrachloride is 7.5 mM and that water exists as monomers. The carbon tetrachloride-water system is

Figure 1. FTIR spectrum of the carbon tetrachloride layer saturated with 1 M NaCl. The water spectrum is magnified by a factor of 1.27 (b) and subtracted from the saturated salt spectrum ([). The difference spectrum (f) reveals enhanced and slightly red-shifted symmetric and asymmetric stretches

thus an excellent environment for probing interactions between limited water and a solute. Using this environment to investigate interactions between water and ions, several chloride salt solutions were prepared, including those of the singly charged cations Na+, K+, and Li+ and the doubly charged cation Mg2+. The anions consisted of the halides, Cl-, Br-, and I-. The MgCl2 solution spectrum is indistinguishable from neat water. The result with the singly charged cations is quite different. The spectral profile obtained with NaCl, shown in Figure 1, is typical of carbon tetrachloride saturated with an aqueous alkali chloride solution. The difference spectrum illustrates that the chloride salt has two perturbing effects on the profile of the water spectrum; (i) the overall amplitude is increased and (ii) the symmetric and asymmetric stretch resonances are slightly redshifted and enhanced, hence the positive peaks. The distinction among the chloride salts is the extent of the overall amplitude increase; all have a similar red shift. Both effects are discussed further below. Larger Halides. Several SFG studies11,25 have shown that larger, more polarizable anions have a substantial impact on the spectrum of water at the surface of salts solutions; specifically, the intensity is enhanced in the 3440 cm-1 region. The reason for the intensity increase has not been identified. With the motivation of unraveling the cause of the intensity increase, solutions of sodium and potassium with both bromide and iodide were examined. As shown in Figure 2, the spectra of the hydrophobic layer all show an amplitude increase. More importantly, some saltssNaBr, KI, and NaIsshow an additional resonance at 3440 cm-1. Interpretations of the additional feature as well as the amplitude increase are contained in the Discussion section. Acid: HCl. SFG spectra of water at the surface of aqueous acid solutions are unique in showing a significant enhancement in the strongest hydrogen-bonded region of the water spectrum.13,15,17,26-28 It is thus of interest to examine the effect of an acid on the spectrum of water confined in carbon tetrachloride. HCl was chosen due to the neutral effect of chloride on both carbon tetrachloride and water. The spectrum of carbon tetrachloride saturated with a concentrated HCl solution, shown in Figure 3, contains a clear spectral signature of molecular HCl centered around 2820 cm-1. The strong central Q band indicates restricted rotational motion. The hydrogen-bonded region

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Figure 2. FTIR spectra of saturated aqueous 1 M halide salt solutions in carbon tetrachloride: (a) NaBr, (b) KBr, (c) NaI, and (d) NaBr. [: bromide, [: iodide, b: water scaled, f: difference. All solutions show enhanced and slightly red-shifted symmetric and asymmetric stretches. Iodide salts and NaBr show a resonance at 3440 cm-1; KBr does not have the 3440 cm-1 resonance.

Figure 3. Spectrum of the carbon tetrachloride phase in contact with concentrated aqueous HCl; peaks around 2800 cm-1 clearly show the presence of molecular HCl. The presence of a Q band in the HCl spectrum indicates restricted rotational motion. The inset is a magnified view of the hydrogen-bonded region showing a broad featureless resonance in addition to the symmetric and asymmetric stretches of water.

(Figure 3 inset) shows a broad absorbance from 3000 to 3600 cm-1 and features due to water that are diminished compared with those of a saturated water-carbon tetrachloride spectrum.24 Gravimetric Analysis. The spectra shown in Figures 1-3 support the conclusion that the salt or acid concentrations in the carbon tetrachloride phases are significant. To quantify the ion concentration, the contents of the carbon tetrachloride layer were gravimetrically analyzed for halide. Table 1 lists the results of the halide analysis: ions are indeed present in the carbon

tetrachloride layer and at significantly higher molar concentrations than is water in a saturated water-carbon tetrachloride solution. These gravimetric analysis results are discussed along with the spectral interpretation in the Discussion section. Large Cation: Ag+. Finally, an infrared spectrum of the carbon tetrachloride layer was taken after gravimetric analysis of the HCl system. In addition to confirming removal of Cl-, the resulting spectrum in the OH region (Figure 4) reveals several notable features: (a) an enhanced water absorbance, (b)

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TABLE 1: Halide Gravimetric Analysis of Carbon Tetrachloride Saturated with 1 M Aqueous Salt or Concentrated Aqueous Hydrochloric Acida

a

salt

AgX mass (g/50 mL) (mg)

molarity (mM)

salt multiple of saturated water

Beer’s law scaling for water

salt/water ratio

NaCl NaBr NaI KCl KBr KI HCl

0.2524 ( 0.3 0.2396 ( 0.3 0.2607 ( 0.3 0.2311 ( 0.3 0.2302 ( 0.3 0.2248 ( 0.3 0.1687 ( 0.2

35.2 25.5 22.2 32.2 24.5 19.1 23.5

4.7 3.4 3.0 4.3 3.3 2.6 3.1

1.27 1.27 2.17

3.70 g2.68 g1.38

1.23 1.26

g2.68 g2.06

The salt concentration is compared with that of water in the neat water-carbon tetrachloride system.

Figure 4. Spectrum of the carbon tetrachloride layer following gravimetric analysis compared with a scaled water spectrum. Note the enhanced symmetric stretch intensity in the post-gravimetric analysis spectrum.

an increased symmetric stretch intensity relative to the asymmetric stretch, and (c) a broad hydrogen-bonded feature. The implications of each of these features are discussed in the next section. Discussion Model. There has been considerable recent progress in developing a picture of the local water structure and dynamics around ions in aqueous solution.29-35 Due to the large oscillator strength, the hydrogen-bond region of water is opaque in the infrared. So information about the local structure of water around ions in aqueous solution comes from Raman,30,31 neutron diffraction,32-34 K-edge X-ray,35 and fast dynamics29,36-38 spectroscopy. The result of these studies is that the effect of ions is quite localized, nearly confined to the first solvation shell.31,33,34 At the surface, Saykally39 reports that I- causes a polarization in the water that results in a significant nonresonant background in the second harmonic generation (SHG) spectrum. It is unclear whether this polarization is an accumulation of local effects or if it is longer range. These studies provide part of the motivation for the current work: to generate an environment containing limited water and ions so that signatures of the local structure are not overshadowed by the larger bulk signal. Raman studies suggest that the vibrational structure of water is primarily affected by the anions.31 The anion associates with the hydrogen atom, resulting in a red shift of the OH vibration from its gas-phase value, the extent of which depends on the identity of the anion. Since the cation associates with the oxygen atom perturbing the nonbonding lone pairs, the primary effect of the cation is to increase the oscillator strength, particularly of the symmetric stretch that is collinear with the ion-water dipole.40-42

Results of the above-mentioned studies will be used to analyze the vibrational spectra of the hydrophobic phase. Although it is challenging to precisely quantify the water concentration, it is possible to determine an upper limit. The oscillator strength for water in carbon tetrachloride is known.24 However, it is also well-known that increased polarization of the water electron cloud due to positive charges associated with the lone pairs increases the oscillator strength of water.40,41,43 Hence a Beer’s law scaling of the water spectrum to determine the water concentration results in an over estimation. In contrast, the salt concentration is determined rather precisely by gravimetric analysis. Thus, a lower limit to the salt:water ratio in the hydrophobic phase is determined by a ratio of the salt concentration to the Beer’s law scaled water concentration. The ratios are listed in Table 1. Note that in every case, the salt concentration is found to be greater than that of water. The conclusion is that the salt ions cannot be surrounded by a hydration shell. This result is consistent with an absence of hydrogen-bonded features in the infrared spectrum. Compelling evidence that substances normally ionized in aqueous solution exist as associated or molecular species in the hydrophobic phase is provided by the HCl data. The HCl spectrum (Figure 3) clearly shows the presence of molecular HCl in the hydrophobic phase. The spectrum contains the rotational envelop of the P and R branches centered at 2820 cm-1, indicating that the HCl rotational lifetime is shortened in carbon tetrachloride compared with that in the gas phase. In addition, the spectrum contains a sharp feature due to the Q branch. Due to selection rules, there is no Q branch for HCl in the gas phase. The presence of the Q branch in carbon tetrachloride is due to interactions pinning the end-over-end rotational motion, leaving a pure vibrational transition at the vibrational origin of 2820 cm-1, red-shifted from the gas phase due to the dielectric constant of carbon tetrachloride. There is precedence for appearance of the pure vibrational band when the rotational motion of HCl is quenched by bonding: The SFG spectrum of the surface of liquid HCl consists of a single resonance at 2800 cm-1 with a half-width of 40 cm-1.44 Thus, the resonance from 2700 to 3000 cm-1 is compelling evidence of molecular HCl. Like HCl, alkali halide salts are highly ionized in aqueous solution. Thus, the presence of molecular HCl supports the interpretation that the salts are also present as associated ion pairs or small clusters in carbon tetrachloride. Indeed, there is precedence for ion strings in comparably concentrated salt solutions, even in aqueous systems for lithium salts.1 There is a question of whether the salts simply shed water due to efflorescence resulting from the low water content of carbon tetrachloride.45 Calculation of the water concentration contradicts efflorescence: 100% RH, corresponds to a water concentration of 1.12 mM, nearly 7 times less than the 7.5 mM

Hydrogen Bonding in a Hydrophobic Environment saturation concentration of water in carbon tetrachloride.24 All salts investigated require a relative humidity of less than 100% to effloresce.46 Thus, the salts do not shed water due to simple efflorescence. Ion-Water Association. The evidence that water associates with the cation rather than the anion is as follows. Neat water in carbon tetrachloride is known to have restricted rotational motion except for rotation about the symmetry axis.43,47 Further, the rotational constant about the symmetry axis is nearly the same as for water in the gas phase,43 indicating that the symmetry axis rotation is essentially a free rotation. The dynamics of water donor-bonded to the anion would be modified in two ways: the symmetry axis rotational motion would be restricted due to interaction between a hydrogen atom and the anion. Second, the two OH stretches would decouple, leading to one free-OH oscillator and one hydrogen-bonded oscillator. The spectral signature would be an increased rotational constant for the symmetry axis rotation, appearance of a resonance due to the free-OH and a broad resonance due to the donor OHstretch. In addition, the vibrational resonance frequency for binding to an anion via H-donation is known to significantly red shift and to vary with the identity of the anion.48 These expectations are all contrary to observation. Alternately, interaction between water and the cation via the water lone pair is known to result in a rotational structure similar to that of water in carbon tetrachloride.43 Consistently, at low RH, water is known to preferentially interact with the cations,49 suggesting that deliquescence begins with water-cation interactions. Hence, the conclusion is that water associates with the ion pair or cluster via interaction with the cation. 3440 cm-1 Peak. The SFG spectra of aqueous halide solutions show enhanced intensity in the water 3440 cm-1 peak; the intensity enhancement is largest for the largest halide.11,15 MD simulations18,19 support the conclusion that halogen ions are in the surface region. What is unclear, however, is how the anions alter the hydrogen-bonding network to produce the intensity enhancement. The model that water is directly bonded to the anion is contradicted by the observation that the frequency lacks the expected increased red shift with anion size.31,33,34,48 Looking to the spectrum of the neat air-water interface provides little clarification given the controversy over interpretation of the H-bonded resonances for water.50-52 This work provides insight into the origin of the 3440 cm-1peak. Some spectra in Figure 2sthose of NaBr, NaI, and KIscontain a resonance at 3440 cm-1. Like the SFG spectra, this feature cannot be due to direct binding to the halide since it occurs at the same frequency for both NaBr and NaI. The only other H-bond acceptor in the system is water, so the resonance is assigned to a water-water hydrogen bond. Salt is essential here since water forms only monomerssdoes not form intermolecular hydrogen bondssin carbon tetrachloride without solutes. The larger anion of the salt plays a role, not by direct binding, but by enabling water to hydrogen bond to another water molecule. The cation-anion combination is similarly important. Ion pairing interactions are known to be critical for understanding solubility of simple inorganic ions, numerous thermodynamic properties of ion pairs, and structure-function in biological systems.53 A pair of small ions forms an inner-sphere or contaction pair due to the electrostatic ion-ion interaction overwhelming the ion-water interaction. A pair of large ions also forms an inner-sphere ion pair because the ion-water interaction is comparable to but slightly weaker than the water-water interaction. The largest salt solubility is found for a combination of a small and a large ion. The small ion-water interaction

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salt

classification

aqueous solubility (M)

NaCl NaBr NaI

inner outer outer

6.2 8.8 11.9

salt

classification

aqueous solubility (M)

KCl KBr KI

inner inner outer

4.8 7.6 8.7

provides the needed interaction energy to compete with the relatively weaker ion-ion electrostatic interaction. In the hydrophobic environment, the electrostatic ion-ion interaction is reflected in the salt solubility (Table 1): for both sodium and potassium the solubility decreases as the anion size increases. As indicated above, water interacts with the salt ion pair or cluster via the cation. In the case where the anion is small, the ion pair forms an inner-sphere ion pair, limiting water to the exterior of the cluster. Conversely, in the case where the anion is large, there is a balance between water and the anion for the inner-sphere position. With water in the inner-sphere position, the anion is relegated to the outer sphere position. The constancy of the donor OH stretch at 3440 cm-1 implies that the anion does not directly bond to the water molecule that screens it from the cation. Instead another water molecule accepts the hydrogen bond in the beginning of a clathrate-like structure. The dangling OH bonds of the acceptor water molecule do not bind to the anion. The data here presented indicates that sodium forms an innersphere ion pair with chloride, potassium with both chloride and bromide. This conclusion is consistent with the aqueous solubility53 (Table 2). Note that the hydrophobic environment provides a setting for finer grained distinctions than the previous classification of ions as kosmotropes and chaotropes that classifies all three halidesschloride, bromide, and iodidesas chaotropes.53 The differences observed here are not consistent with this broad classification. As shown by Arrhenius, for significant salt concentrations the oppositely charged ions do not behave independently and deviations from ideality are common. This work provides a close look at water in the vicinity of inner- vs outer-sphere ion pairs. Division into inner-sphere and outer-sphere complexes is consistent with the differential ion distribution calculated with MD simulations4,5 and the intensity variation seen in SFG measurements.11,15 The results of this work suggest that it is the structure of water sandwiched between the cation deeper in the interface and the anion closer to the surface that is responsible for the 3440 cm-1 intensity. Further, interaction with the cation alters the polarizability and orientation, leaving the anion accessible to interaction with impinging gas-phase molecules.2 Salts: Ag+. Compelling evidence for a direct interaction between water and the cation is shown by the spectrum in Figure 4. Following gravimetric analysis to precipitate chloride as AgCl, the carbon tetrachloride phase contains Ag+ ions due to addition of excess AgNO3, and NO3- plus H3O+ primarily from the original HCl. The spectrum in Figure 4 reflects the presence of both Ag+ and H3O+. Ag+, as a large, polarizable cation, enhances the oscillator strength of the symmetric stretch of water as is well-known for cations bonded to the lone pair.40,41 Figure 4 shows this oscillator strength increase as an enhanced symmetric stretch intensity supporting a picture in which water bonds to the large Ag+ ion. Acid. The spectra shown in Figures 3 and 4 contain a broad, featureless, absorbance from about 3000 cm-1 to nearly 4000

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cm-1. The broad absorbance reflects a dynamic species involved in hydrogen bonding. This must include water hydrating H3O+; H3O+ remains in the hydrophobic phase after chloride precipitation and is responsible for the broad tail in Figure 4. The broader, more intense tail for aqueous HCl, Figure 3, indicates that this tail must include hydrogen-bonded water in a hydration shell around molecular HCl. (The resonance is not assigned to H3O+ directly since H3O+ is expected to have weaker OH bonds than water and thus be red of 3000 cm-1.) The presence of molecular HCl is clearly shown by the resonance from 2700-3000 cm-1 including a Q band due to restricted rotational motion. (A similar Q band is observed on the surface of liquid HCl.44) Assignment to water incorporated in a hydration shell is consistent with the reduced free water absorption shown in Figure 3. The lack of a similar broad absorbance for the salt solutions further supports the interpretation that water associated with the salts is isolated. Salt Concentration. Table 1 lists the salt concentration in the hydrophobic phase. Note that the salt concentration depends on both the anion and the cation identity: within one cation the salt concentration decreases as the anion size increases, for a given anion the salt concentration is smaller for a larger cation. Thus a model for salt transport across the interface requires attention to both the anion and the cation. Generally, a stronger electrostatic interaction results in a greater salt concentration. Thus the salt concentrations are ordered: NaCl > KCl > NaBr > KBr > NaI > KI. In general, high charge density ions interact strongly with water in an aqueous phase; these same ions form ion pairs in the hydrophobic phase. Lower charge densities, as in the larger ions, result in a lower concentration in the hydrophobic phase. Summary/Conclusions A combined spectroscopic and gravimetric analysis of a hydrophobic solvent, carbon tetrachloride, saturated with 1 M aqueous salt solutions or concentrated aqueous HCl clearly shows the presence of associated ions in the hydrophobic phase. The spectra (Figures 1, 2, and 4) reflect interaction with the cation by the increased oscillator strength and slight red shift. The anion appears to not be directly associated with watersthe hydrogen bonded resonance frequency does not shift with anion. Rather, formation of an inner-sphere ion pair limits water to the outside of the salt cluster, resulting in only water monomers in the hydrophobic phase. Formation of an outer-sphere ion pair puts water between the anion and the cation. The sandwiched water molecule does not donate a hydrogen bond to the anion but rather to another water molecule resulting in an OH resonance at 3440 cm-1. For every salt investigated, the hydrophobic-phase ion concentration is greater than the water concentration. Hence, rather than isolated ions entering the hydrophobic phase with an intact hydration shell, the ion pairs or cluster appears to be associated with a limited number of water molecules. The salt concentration is determined by the electrostatic interaction between the ions: the smallest anion-cation combination has a hydrophobic-phase concentration nearly double that of the largest anion-cation combination. There is precedent for such ion cluster formation.1 Water associated with the HCl appears to be quite different. With HCl, water hydrates molecular HCl, forming hydrogen bonds with other water molecules or with the H+ of HCl. The presence of molecular HCl is confirmed by the vibrational signature around 2800 cm-1. The presence of a vibrational Q band shows that, like water, HCl has restricted rotational motion.

Bisson et al. The picture that emerges is one of ions in the hydrophobic phase not accompanied by a hydration shell. Instead, ions associate forming ion pairs or clusters. Water interacts with the associated ions via a water-cation interaction that supports additional water in the hydrophobic phase. Interaction between the cation and the anion controls the salt concentration. Water associated with ion pairs or clusters appears much as isolated water: rotational motion is restricted perpendicular to the symmetry axis, but essentially free around the symmetry axis. If the ion combination forms an outer-sphere ion pair, the cluster supports limited water-water interaction and H-bond formation resulting in a resonance at 3440 cm-1. Acknowledgment. Support from the American Chemical Society, Petroleum Research fund grant NO. 46671-AC6 and from the National Science Foundation grant number NSFCHE0613757 for partial support of this work is gratefully acknowledged. References and Notes (1) Thomas, A. S.; Elcock, A. H. J. Am. Chem. Soc. 2007, 129, 14887– 14898. (2) Finlayson-Pitts, B. J. Chem. ReV. 2003, 103, 4801–4822. (3) Ghosal, S.; Hemminger, J.; Bluhm, H.; Mun, B.; Hebenstreit, E.; Kettler, G.; Ogletree, F.; Requejo, F.; Salmeron, M. Science 2005, 307, 563–566. (4) Ghosal, S.; Brown, M.; Bluhm, H.; Krisch, M.; Salmeron, M.; Jungwirth, P.; Hemminger, J. J. Phys Chem. A 2008, 112, 12378–12384. (5) Brown, M.; D’Auria, R.; Kuo, W.; Krisch, M.; Starr, D.; Bluhm, H.; Tobias, D.; Hemminger, J. Phys. Chem. Chem. Phys. 2008, 10, 4778– 4784. (6) Salvador, P.; Curtis, J.; Tobias, D.; Jungwirth, P. Phys. Chem. Chem. Phys. 2003, 5, 3752–3757. (7) Thomas, J.; Roeselova, M.; Dang, L.; Tobias, D. J. Phys. Chem. A 2007, 111, 3091–3098. (8) Brown, M. A.; Winter, B.; Faubel, M.; Hemminger, J. C. J. Am. Chem. Soc. 2009, 131, 8354–8355. (9) Ghosal, S.; Hemminger, J. C.; Bluhm, H.; Mun, B. S.; Hebenstreit, E. L. D.; Ketteler, G.; Ogletree, F.; Requejo, F. G.; Salmeron, M. Science 2005, 108, 563–566. (10) Shultz, M. J. In AdVances in Multiphoton Processes and Spectroscopy; Lin, S. H., Villaeys, A. A., Fujimura, Y., Eds.; World Scientific: Singapore, Japan, 2008; Vol. 18, p 133-200. (11) Gopalakrishnan, S.; Liu, D.; Allen, H. C.; Kuo, M.; Shultz, M. J. Chem. ReV. 2006, 106, 1155–1175. (12) Shultz, M. J.; Baldelli, S.; Schnitzer, C.; Simonelli, D. J. Phys. Chem. B 2002, 106, 5315–5324. (13) Baldelli, S.; Schnitzer, C.; Shultz, M. J. Chem. Phys. Lett. 1999, 302, 157–163. (14) Levering, L. M.; Sierra-Hernandez, M. R.; Allen, H. C. J. Phys. Chem. C 2007, 111, 8814–8826. (15) Mucha, M.; Frigato, T.; Levering, L. M.; Allen, H. C.; Tobias, D. J.; Dang, L. X.; Jungwirth, P. J. Phys. Chem. B 2005, 109, 7617–7623. (16) Gopalakrishnan, S.; Jungwirth, P.; Tobias, D. J.; Allen, H. C. J. Phys. Chem. B 2005, 109, 8861–8872. (17) Baldelli, S.; Schnitzer, C.; Shultz, M. J.; Campbell, D. J. Phys. Chem. B 1997, 101, 10435–10441. (18) Jungwirth, P.; Tobias, D. J. Chem. ReV. 2006, 106, 1259–1281. (19) Jungwirth, P.; Winter, B. Annu. ReV. Phys. Chems. 2008, 2008.59.343.366, 343–366. (20) Kuo, I.-F. W.; Mundy, C. J.; Eggimann, B. L.; McGrath, M. J.; Siepmann, J. I.; Chen, B.; Vieceli, J.; Tobias, D. J. J. Phys. Chem. B 2006, 110, 3738–3746. (21) Benjamin, I. Science 1993, 261, 1558–1560. (22) Wick, C.; Dang, L. X. J. Chem. Phys. 2007, 126, 134702-1 to 134702-4. (23) Fennell, C.; Bizjak, A.; Vlachy, V.; Dill, K. J. Phys Chem. B 2009, 113, 6782–6791. (24) Kuo, M. H.; David, A.; Kamelamela, N.; White, M.; Shultz, M. J. J. Phys. Chem. C 2007, 111, 8827–8831. (25) Raymond, E. A.; Richmond, G. L. J. Phys. Chem. B 2004, 108, 5051–5059. (26) Schnitzer, C.; Baldelli, S.; Campbell, D. J.; Shultz, M. J. J. Phys. Chem. A 1999, 103, 6383–6386. (27) Baldelli, S.; Schnitzer, C.; Shultz, M. J.; Campbell, D. J. Chem. Phys. Lett. 1998, 287, 143–147.

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