Contact Ion Pair Formation between Hard Acids and Soft Bases in

The interaction of charged species in aqueous solution has important implications for chemical, biological, and environmental processes. We have used ...
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Contact Ion Pair Formation between Hard Acids and Soft Bases in Aqueous Solutions Observed with 2DIR Spectroscopy Zheng Sun, Wenkai Zhang, Minbiao Ji,† Robert Hartsock, and Kelly J. Gaffney* PULSE Institute, SLAC National Accelerator Laboratory, Stanford University, Stanford, California 94305, United States S Supporting Information *

ABSTRACT: The interaction of charged species in aqueous solution has important implications for chemical, biological, and environmental processes. We have used 2DIR spectroscopy to study the equilibrium dynamics of thiocyanate chemical exchange between free ion (NCS−) and contact ion pair configurations (MNCS+), where M2+ = Mg2+ or Ca2+. Detailed studies of the influence of anion concentration and anion speciation show that the chemical exchange observed with the 2DIR measurements results from NCS− exchanging with other anion species in the first solvation shell surrounding Mg2+ or Ca2+. The presence of chemical exchange in the 2DIR spectra provides an indirect, but robust, determinant of contact ion pair formation. We observe preferential contact ion pair formation between soft Lewis base anions and hard Lewis acid cations. This observation cannot be easily reconciled with Pearson’s acid−base concept or Collins’ Law of Matching Water Affinities. The anions that form contact ion pairs also correspond to the ions with an affinity for water and protein surfaces, so similar physical and chemical properties may control these distinct phenomena.

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INTRODUCTION How water solvates ions and mitigates the interaction between charged species in solution has wide-ranging implications in chemistry, biology, and environmental science. The propensity for ions to form contact ion pairs in solution, in particular, the propensity for ions to bind to charged side chains at protein surfaces, has been hypothesized to explain the influence of ions on the solubility of proteins, an effect categorized by the Höfmeister series.1−5 Collins has postulated a Law of Matching Water Affinities to explain the influence of ions on protein solvation.3,4 The proposal of Collins states that only ions with similar charge densities form contact ion pairs. The Law of Matching Water Affinities conforms to the theory of hard−soft Lewis acid−base chemistry, often termed the Pearson acid− base concept.6,7 While Collins has demonstrated that many thermodynamic properties of aqueous ionic solutions can be correlated with his Law of Matching Water Affinities, more direct conformation has been lacking. Determining the validity of the Law of Matching Water Affinities requires the robust determination of contact ion pair (CIP) formation in aqueous solution, but clear identification of CIP formation has proven to be challenging. A variety of spectroscopic methods have been utilized to study the equilibrium and dynamical properties of aqueous ionic solutions.8−19 To date, dielectric relaxation spectroscopy has been the most useful technique for studying the interaction of ions in solution, providing information about the association and rotation of CIPs, solvent-separated ion pairs (SIPs), and doubly solvent-separated ion pairs (2SIP).8−14 This sensitivity also makes analysis challenging because the contribution of each ion pair configuration exhibits complex, overlapping lineshapes. Time-resolved IR spectroscopy of molecular anions provides an alternative approach to characterizing CIP equilibria in © XXXX American Chemical Society

aqueous solution for molecular ions. The IR spectra of a variety of anions shift measurably when they form a CIP with alkali and alkaline earth metal cations in polar solvents.20−27 Of these molecular anions, thiocyanate (SCN¯), isocyanate (OCN¯), and azide (N3¯) provide the most tractable IR spectroscopy, particularly for time-resolved measurements.19,22,24,25,28−32 We have used the NCS− anion to study contact ion pairing in water because of the long excited-state lifetime of the CN-stretch vibration.28,33 The sensitivity of the nitrile stretch frequency to variations in the nitrile chemical and electrostatic environment has been used in a variety of prior experiments.34,35 Under certain circumstances, the CN-stretch frequency provides a probe of the local electrostatic environment.34,36 Stronger intermolecular interactions, such as CIP formation or hydrogen-bond formation to the nitrile group, lead to shifts in the CN-stretch frequency that cannot be accounted for with the Stark effect.37,38 For the specific case of CIP formation between Mg2+ or Ca2+ and NCS−, bonding of the M2+ cation to the nitrogen atom in the NCS− anion induces an electronic polarization that shifts electron density from the CS bond to the CN bond, leading to an increase in the CN-stretch frequency.24 This electronic polarization induced by CIP stabilizes the CN triple-bond resonance structure of thiocyanate, an alternative means of explaining the increase in the CN-stretch frequency. In aqueous solution, the water− thiocyanate interaction also perturbs the CN-stretch frequency and generates an inhomogeneous broadened spectrum. Accurate description of the CN-stretch absorption spectrum requires modeling the distribution of cation and water solvation Special Issue: Michael D. Fayer Festschrift Received: April 5, 2013 Revised: July 24, 2013

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focus three mid-IR pulses with an off-axis parabolic mirror (focal length of 150 mm) onto the sample in a boxcar geometry and collimate the beams after the sample with another off-axis parabolic mirror (focal length of 150 mm). The spot size of the IR beams at the sample position is ∼50 μm in diameter. We control the relative time of these three pulses with computer controlled translational stages (ANT-50L, Aerotech). The sample emits the signal in a unique phase-matched direction, which we overlap with a local oscillator pulse for heterodyne detection. A grating in a spectrometer (iHR320, Horiba) disperses the heterodyned signal onto the top stripe of a dual 32 × 2 element mercury−cadmium−telluride (MCT) array detector with high speed data acquisition electronics (FPAS6416-D, Infrared Systems Development). We purchased the salts and deuterium oxide from SigmaAldrich and used them as received. The experiments used a sample cell with two CaF2 windows (3 mm thick) and a 6 μm thick Teflon spacer to reduce the D2O background absorption.

structures surrounding thiocyanate in solution and how these structures influence the CN-stretch frequency. Despite these complications, the MNCS+ CIP, where M2+ = Mg2+ or Ca2+, shifts the CN stretch frequency sufficiently to distinguish the influence CIP formation from the influence of water-SCN¯ Hbonding on the absorption spectrum. These distinct vibrational transition energies for the CIP and the free anion configurations provide the opportunity to use 2DIR spectroscopy to investigate the equilibrium dynamics of ion association and dissociation in aqueous ionic solutions.24,25,32,39,40 We present a multidimensional vibrational correlation spectroscopy, or two-dimensional infrared (2DIR) spectroscopy, study where we use the observation of chemical exchange between CIP and free anion configurations for the thiocyanate anion as a means of determining which anions form CIP with Mg2+ and Ca2+ in aqueous solution. 2DIR provides a robust experimental method for investigating equilibrium structural and conformational dynamics on the picosecond time scale.41−44 2DIR has been used to study solute−solvent complexation,45 C−C single bond rotation,46 pseudorotation in iron pentacarbonyl, 47 ion pairing and assembly in solution,24,25,32,40 and hydrogen-bond switching in aqueous ionic solutions.48−51 The present study will demonstrate that the chemical exchange observed in the 2DIR measurements results from anion exchange in the first solvation shell of the alkali earth cation: MNCS+ + A− ⇌ MA+ + NCS− 2+

2+

2+

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RESULTS 1. FTIR Spectroscopy. The linear triatomic thiocyanate anion (NCS−) has three IR-active vibrations: the CS stretching mode (747 cm−1), the CN stretching mode (2068 cm−1), and the doubly degenerate bending mode (470 cm−1).25−27 When dissolved in water, Na+ and NCS− do not form CIPs, yielding a FTIR spectrum with a single CN stretching absorption peak at 2068 cm−1 associated with water solvated NCS− anions. The situation changes for the larger charge density alkali earth cations, Mg2+ and Ca2+. Aqueous solutions containing NCS− and Mg2+ form N-bound 1:1 CIPs, as well as water solvated anions. This leads to an additional CN stretch absorption peak at 2115 cm−1 for MgNCS+ CIP, as shown in Figure 1. The N-

(1)

−

where M = Mg or Ca and A corresponds to a variety of different monovalent anions. This confirms the hypothesis proposed in Park et al.25 and extends that work by investigating a variety of different anions. Anion exchange happens when two distinct CIP species exist in solution, the MNCS+ CIP we observe directly and the MA+ CIP we infer from the existence of chemical exchange in the 2DIR spectroscopy measurement. When chemical exchange occurs we can conclude the MA+ CIP exists in solution. Significantly and surprisingly, we observe that CIP formation occurs preferentially with the soft bases, even though Mg2+ and Ca2+ cations are hard acids.

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EXPERIMENTAL METHODOLOGY The laser system employed in the experiments was built based on a design that has been described in detail elsewhere.52 We generate 800 nm pulses with a Ti:sapphire oscillator (KM Laboratories) and regenerative amplifier (Spitfire, SpectraPhysics) laser system at 1 kHz. The 800 nm pulses with 45 fs duration and ∼1 mJ per pulse were used to pump an optical parametric amplifier (OPA800CF, Spectra-Physics) to produce near-IR pulses at ∼1.4 and ∼1.9 μm, which were utilized to generate mid-IR pulses at 2050 cm−1 in a 0.5 mm thick AgGaS2 crystal by difference frequency generation. The power spectrum of the mid-IR pulses had a Gaussian envelope with a ∼250 cm−1 bandwidth (full width at half-maximum). We measured the pulse chirp with frequency-resolved optical gating measurements in a transient grating geometry.52 We used CaF2 plates with different thicknesses to compensate for the linear dispersion introduced by other dielectric materials in the setup, particularly a Ge Brewster plate. This setup produced transform-limited mid-IR pulses with pulse durations of ∼65 fs at the sample. The experimental details and principles of multidimensional vibrational correlation spectroscopy, generally termed 2DIR spectroscopy, have been described in detail elsewhere.52,53 We

Figure 1. (A) FTIR spectra of 0.12, 0.24, 0.4, and 1.2 M NaNCS dissolved in a 3 M MgI2 deuterated water solution. The spectrum for the 1.2 M NaNCS solution is drawn in at one-third of the real amplitude. (B) FTIR spectra of 0.24 and 1.2 M NaNCS in 1.5 M MgI2 deuterated water solution. (C) FTIR spectra of 0.24 M NaNCS in a 3 M MgA2 deuterated water solution, where A− = Cl− or Br−.

bound 1:1 CIPs between Ca2+ and NCS− have a CN-stretch absorption peak at 2093 cm−1, as shown in Figure 2. The watersolvated NCS− anions with a CN stretch absorption peak at 2068 cm−1 will be referred to as free NCS− anions, although the SIPs that exist in solution will also have a CN-stretch absorption peak that cannot be distinguished from free NCS− anions. Prior dielectric relaxation spectroscopy studies by Buchner et al. on aqueous Na2SO4 solutions show that SIPs do not represent an appreciable fraction of the ions for concentrated ionic solutions.11,12 This result indicates that anions and cations that do not form CIP move independently of one another in concentrated ionic solutions. The relative intensity of the two CN-stretch absorption peaks in the FTIR spectrum reflects the equilibrium constant between the CIP and free configurations B

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Figure 3. 2DIR spectra for ionic solutions with (A) 1.5 M MgI2 and 0.24 M NaNCS in deuterated water and (B) 1.5 M MgI2 and 1.2 M NaNCS in deuterated water. All spectra collected with TW = 40 ps.

Figure 2. FTIR spectra for 0.35 M NaNCS with or without 3 M CaI2 dissolved in deuterated water.

K=

[MNCS+] [M2 +][NCS−]

(2)

where M2+ = Mg2+ or Ca2+. The solutions presented in Figures 1A and 2 have a much higher [M2+] concentration than the [NCS−] concentration, making [M2+] roughly constant for all thiocyanate concentrations. As clearly shown by the equilibrium constant, using an excess of M2+ allows the ratio of the free anion and CIP conformations to remain roughly constant over a large range of NaNCS concentration. Quantitative determination of the concentrations of NCS− in the free and CIP configurations must account for the impact of CIP formation on the absorption cross section but does not influence our qualitative analysis of the 2DIR spectra. 2. 2DIR Spectroscopy. 2DIR spectroscopy monitors thermal equilibrium dynamics occurring on the picosecond time scale by vibrationally labeling molecules with their initial frequencies (ωτ) and then recording the final frequencies (ωm) of the initially labeled molecules after an experimentally controlled waiting time (Tw).41,42,45 This spectroscopic technique can observe chemical exchange events as slow as a few times the vibrational lifetime and as fast as 1/Δω, where Δω equals the vibrational frequency difference between the interconverting species. 2DIR spectra, S(ωτ,ωm,Tw), are displayed by correlating the initial frequency (ωτ) and final frequency (ωm) as a function of waiting times (Tw). The presence or absence of chemical exchange and the mechanistic origin of the exchanges represent the focus of the present manuscript. A detailed numerical analysis of the 2DIR spectra focused on the extraction of chemical exchange rates has been performed for a subset of the solutions presented in this study and will be addressed in a later manuscript. 2A. 2DIR Spectroscopy of Aqueous Solutions Containing Mg2+ and NCS−. Figures 3−5 show 2DIR spectra obtained for a variety of solutions containing Mg2+ and NCS− at waiting times of 40 ps in Figures 3 and 4 and 30 ps in Figure 5. For the spectral range corresponding to the CN stretch of thiocyanate, the 2DIR spectra can possess up to eight distinct vibrational transitions, although spectral overlap and slow chemical exchange only lead to the appearance of four distinct peaks and two shoulders for TW = 30 or 40 ps. Two positive peaks along the diagonal (ωτ = ωm) result from the fundamental vibrational transitions (υ = 0 → 1), with the higher frequency diagonal peak at ωτ = ωm = 2115 cm−1 resulting from the SCN¯ CIP and the lower frequency diagonal peak at ωτ = ωm = 2068 cm−1 resulting from the free SCN¯. For the TW = 30 or 40 ps 2DIR spectra, the relative intensities for the diagonal peaks do not match the relative peak heights in the FTIR spectra. This

Figure 4. 2DIR spectra for ionic solution with 3 M MgI2 and (A) 0.4 M NaNCS, (B) 0.24 M NaNCS, and (C) 0.12 M NaNCS dissolved in deuterated water. All spectra collected with TW = 40 ps.

Figure 5. 2DIR spectra for ionic solutions with 0.24 M NaNCS and 3 M concentrations of (A) MgCl2, (B) MgBr2, and (C) MgI2 dissolved in deuterated water. All spectra collected with TW = 30 ps.

results primarily from the slower anisotropy decay and longer CN-stretch lifetime for the CIP conformation.25 Polarizationresolved pump−probe data can be found in the Supporting Information. Two negative peaks below the diagonal arise from the υ = 1 → 2 transitions. The υ = 1 → 2 transitions have ωm values red-shifted relative to ωτ by the vibrational anharmonicity. Chemical exchange, in principle, leads to an additional four peaks, although we only resolve a positive shoulder at (ωτ = 2068 cm−1, ωm = 2115 cm−1) and a negative shoulder at (ωτ = 2115 cm−1, ωm = 2032 cm−1). As seen in Figure 4, the peak separation of the υ = 0 → 1 transitions resembles the vibrational anharmonicities. This leads to significant overlap in the free anion υ = 0 → 1 absorption and the υ = 1 → 2 free anion to CIP chemical exchange cross-peak as well as the CIP υ = 1 → 2 absorption and the υ = 0 → 1 CIP to free anion chemical exchange cross-peak. As can be clearly seen in Figures 3−5, the presence or absence of chemical exchange between MgNCS+ CIPs and NCS− free anions depends on the concentration and anion speciation in solution. The concentration and speciation dependence of the chemical exchange rate has allowed us to definitively determine the nature of the chemical exchange occurring on the tens of picoseconds time scale in these C

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observations that chemical exchange on the tens of picoseconds time scale for [NCS−] ≤ 0.4 M occurs when chemically distinct anions exchange between the first and second solvation shells surrounding Mg2+. This conclusion should not be viewed as ruling out water−anion exchange in aqueous Mg2+ solutions but rather that water−anion exchange occurs too slowly to be observed for TW ≤ 50 ps. This conclusion agrees with a large number of prior measurements that show a water residence time of a few microseconds even for high Mg2+ concentrations.56 The absence of anion exchange in the MgCl2 solution may reflect the absence of Mg2+-Cl− CIP formation, consistent with prior solubility studies of aqueous MgCl2 solutions.57 2B. 2DIR Spectroscopy of Aqueous Solutions Containing Ca2+ and NCS−. We have performed a similar series of measurements for aqueous solutions containing NCS− and Ca2+ ions as we performed for aqueous NCS− and Mg2+ solutions. While the spectral overlap of the CaNCS+ CIP and free NCS− make the measurements more challenging to analyze, we clearly see the impact of chemical exchange in the 2DIR spectra. By using [NCS−] ≤ 0.4 M, we could rule out (A) energy transfer and (B) NCS− “self” exchange as the origin of cross-peaks in the 2DIR spectra. To distinguish between (C) anion exchange and (D) water−anion exchange, we varied the secondary anion speciation, as already discussed for Mg2+ solutions. Figure 7 shows the 2DIR spectra for TW = 30 ps

solutions. In principle, the exchange dynamics we observe in our 2DIR measurements could originate from four distinct processes: (A) energy transfer from a vibrationally excited NCS− in one configuration to a vibrational ground-state molecule in the other configuration,30,54,55 (B) the “self” exchange of a vibrationally excited NCS− in either the free or the CIP conformation with a ground-state NCS− in the opposite conformation, (C) anion exchange between thiocyanate and the secondary anion in the first solvation shell of Mg2+, and (D) anion−water exchange where the thiocyanate anion replaces or is replaced by a water molecule in the first solvation shell of Mg2+. We show these distinct phenomena schematically in Figure 6.

Figure 6. Schematic depiction of the four distinct phenomena that could lead to cross peaks in the 2DIR spectra with M2+ = Mg2+ or Ca2+, *NCS− representing vibrationally excited thiocyanate, and A− representing the secondary anion in solution: (A) vibrational energy transfer, (B) anion “self” exchange, (C) anion exchange, and (D) water−anion exchange. For appropriately low NCS− concentrations, the cross peaks result from anion exchange. Please see the main text for a discussion of how these different processes can be distinguished experimentally.

The first two potential sources of chemical exchange depend critically on the NCS− concentration. Figure 3A shows that for TW = 40 ps no chemical exchange occurs for an aqueous solution containing 0.24 M NaNCS and 1.5 M MgI2, but Figure 3B clearly shows that chemical exchange does occur for a TW = 40 ps for a 1.2 M NaNCS and 1.5 M MgI2 solution. This exchange could occur because of (A) energy transfer or (B) “self” exchange. For the present discussion, Figure 3A,B allows us to rule out these sources of exchange on the tens of picoseconds time scale in the 2DIR spectra as long as we utilize [NCS−] ≤ 0.4 M in our experiments. Figure 4 shows 2DIR spectra at TW = 40 ps for NCS− concentrations of 0.4, 0.24, and 0.12 M NaNCS dissolved in a 3 M MgI2 aqueous solutions. All three spectra show the clear indication of chemical exchange for NCS− concentrations too low to observe either (A) energy transfer or (B) “self” exchange within 40 ps. The observed exchange could reflect either (C) anion exchange or (D) water−anion exchange. To distinguish between these two chemical exchange processes, we have varied the speciation of the secondary anion in solution. Figure 5 shows the 2DIR spectra for TW = 30 ps for 0.24 M NaNCS dissolved in aqueous Mg2+ solutions with three distinct anions: (A) 3 M MgCl2, (B) 3 M MgBr2, and (C) 3 M MgI2. The water−anion exchange mechanism should be insensitive to the nature of the secondary anion in solution because the secondary anion does not participate directly in the chemical reaction. The data in Figure 5 clearly show that the exchange process does depend on the nature of the secondary anion. Figure 5A shows no evidence of chemical exchange for TW = 30 ps for 3 M MgCl2, but Figure 5B,C does show exchange for TW = 30 ps for 3 M MgBr2 and 3 M MgI2. We conclude from these

Figure 7. 2DIR spectra for ionic solutions with 0.2 M Ca(SCN)2 and 3 M concentrations of (A) CaCl2, (B) CaBr2, (C) CaI2, and (D) Ca(ClO4)2 dissolved in deuterated water. All spectra collected with TW = 30 ps.

for aqueous solutions containing 0.2 M Ca(NCS)2 and 3 M (A) CaCl2, (B) CaBr2, (C) CaI2, and (D) Ca(ClO4)2. The 2DIR spectrum for the CaCl2 solution shown in Figure 7A has a signal that is clearly extended along the diagonal of the spectrum but does not have appreciable off-diagonal intensity. We conclude that no chemical exchange occurs for TW = 30 ps for this solution. We use this observation to conclude that the exchange between water and NCS− in the first solvation shell of Ca2+ occurs too slowly to observe for TW ≤ 50 ps. The 2DIR spectra for the CaBr2, CaI2, and Ca(ClO4)2 solutions shown in Figure 7B−D exhibit distinct spectra with intensity apparent in the (ωτ = 2068 cm−1, ωm = 2093 cm−1) cross-peak region D

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consistent with NCS − exchange from the free anion configuration to a CaNCS+ CIP configuration. Unlike Mg2+, The slow rate of water-NCS− exchange cannot be corroborated with prior measurements because the residence times for water in the first solvation shell of Ca2+ extracted from experiment17,58 and molecular dynamics simulations59 range from 15 ps to roughly 100 ns.

the equilibrium dynamics of NCS− chemical exchange between free ion and CIP configurations MNCS+, where M = Mg2+ or Ca2+. Our studies clearly show that the chemical exchange observed with the 2DIR measurements results from anion exchange in the first solvation shell surrounding Mg2+ and Ca2+ cations. The presence of chemical exchange provides an indirect, but robust, determinant of CIP formation between ion pairs that lack alternative spectral signatures. Surprisingly, we observe soft Lewis bases such as Br−, I−, CIO4−, and NCS− forming contact ions pairs with the hard Lewis acid cations Mg2+ and Ca2+. Consistent with prior thermodynamic and spectroscopic measurements,65,66 we do not observe CIP formation between Mg2+ and Cl− or NO3−, even though Cl− and NO3− are harder bases than Br−, I−, CIO4−, or NCS−. This series of observations cannot be easily reconciled with Pearson’s acid−base concept,6,7 Collins’ Law of Matching Water Affinities,3,4 or the molecular dynamics simulations of Fennell et al.67,68 We believe the inability of Pearson’s acid−base concept to explain our findings indicate the ion−solvent interaction has a level of importance similar to that of the ion−ion interaction. Emphasis of the ion−water interaction may highlight the importance of different ion properties, such as ion charge density and surface area. Interestingly, the anions that form CIPs also correspond to the ions with an affinity for water surfaces and protein surfaces,1,2,60−64 so similar physical and chemical properties may control all of the observed phenomena. The molecular dynamics study of Geissler et al.60 indicates that anion segregation to the surface of water results in part from an unfavorable anion hydration energy for anions with low charge density. For the interior of a solution, surface segregation cannot be achieved, but CIP formation will reduce the anion surface area that must be hydrated directly by coordinating water molecules and reduce the free-energy cost of cavity formation. Investigating the veracity of this hypothesis will likely require a simulation methodology including polarizability or quantum chemical effects to accurately model the energetics for CIP formation.62,69,70

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DISCUSSION The Results section presents a detailed investigation of the concentration and secondary anion speciation dependence of the free thiocyanate, MNCS+, where M2+ = Mg2+ or Ca2+, CIP chemical exchange dynamics measured with 2DIR spectroscopy. We have determined from these measurements that anion exchange in the first solvation shell surrounding Mg2+ or Ca2+ cations dominates the reaction dynamics on the tens of picoseconds time scale observed with 2DIR spectroscopy. While alternative exchange mechanisms can occur in solution, for sufficiently low concentrations of NCS−, none of the alternative mechanisms occur on such a fast time scale. The CN-stretch excited-state lifetime sets the time window of TW ≤ 50 ps over which we can observe chemical exchange. We observe chemical exchange between thiocyanate and A− when A− = Br−, I−, CIO4− but not when A− = Cl−. The absence of chemical exchange for secondary Cl− anions appears consistent with the results of solubility studies of aqueous MgCl2 and CaCl2 solutions that indicate that MCl+ contact ion pairing does not occur for the concentrations we have investigated.57 The presence of chemical exchange in the 2DIR spectroscopy measurements robustly signals the presence of CIP formation between Mg2+ and Ca2+ and a variety of anions in solution. These measurements have confirmed that CIP formation occurs for MgBr+, MgI+, MgClO4+, CaBr+, CaI+, and CaClO4+. Surprisingly, we observe CIP formation occurring between the hard Lewis acid cations Mg2+ and Ca2+ and the soft Lewis base anions Br−, I−, NCS−, and CIO4−. The observation of CIP formation between hard acids and soft bases does not follow the expected trend of Pearson’s acid− base concept6,7 or Collins’ Law of Matching Water Affinities.3,4 We presently lack a clear explanation for this counterintuitive observation, but a few points do warrant emphasis. Water functions as both a Lewis acid and base, so the consideration of the cation−anion interaction independent of the interaction energy of the ionic species with water provides an overly simplistic starting point for understanding ion solvation in water. Simple point-charge models also appear insufficient for describing the ion pairing in aqueous solution. The residence time of water molecule in the first solvation shell of a Mg2+ and Ca2+ cation exceeds the residence time of numerous anions. This clearly does not correlate with the simple -minded energetics expected for charge−charge versus charge−dipole interactions. Lastly, the anions we observe in contact with Mg2+ and Ca2+ also correspond to the anions that tend to be segregated to the surface of aqueous solutions and dissolved proteins.1,2,60−64

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ASSOCIATED CONTENT

S Supporting Information *

FTIR spectra for all solutions studied with 2DIR spectroscopy and population and orientational relaxation dynamics measured with polarization resolved pump−probe spectroscopy. This material is available free of charge via the Internet at http:// pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Tel: 650-926-2382. E-mail: kgaff[email protected]. Present Address †

Minbiao Ji: Department of Chemistry and Chemical Biology, Harvard University, Cambridge, Massachusetts, 02138, United States. Notes

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The authors declare no competing financial interest.

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CLOSING REMARKS The interaction of charged species in aqueous solution has important implications for chemical, biological, and environmental processes. Despite the obvious importance of ion hydration and assembly, many phenomena currently defy clear explanation. We have presented a 2DIR spectroscopy study of

ACKNOWLEDGMENTS We acknowledge support from the AMOS program within the Chemical Sciences, Geosciences, and Biosciences Division of the Office of Basic Energy Sciences, Office of Science, U.S. Department of Energy. E

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