Article pubs.acs.org/JPCB
Effects of Carboxylate Anions on the Molecular Organization of H2O as Probed by 1-Propanol Takemi Kondo, Yuji Miyazaki, and Akira Inaba Research Center for Structural Thermodynamics, Graduate School of Sciences, Osaka University, Toyonaka, Osaka 560-0043 Japan
Yoshikata Koga* Department of Chemistry, The University of British Columbia, Vancouver, BC Canada V6T 1Z1 S Supporting Information *
ABSTRACT: We characterized the effects of carboxylate anions, formate (OFm−), acetate (OAc−), and propionate (OPr−), on the molecular organization of liquid H2O by the 1-propanol (1P) probing methodology. The latter thermodynamic methodology provides two indices: one pertaining to the hydration number, nH, and the other being related to the net increase/decrease of the entropy−volume cross fluctuation of the system. The results indicated that OFm− is a hydration center with nH = 1.2 ± 0.5 and leaves the bulk H2O away from the hydration shell unperturbed. We suggest that this single H2O hydrates preferentially one of the O’s in the COO− group, showing the hydration center character. The values of nH for OAc− and OPr− were found to be 3.7 ± 0.8 and 9 ± 2, respectively, out of which one H2O molecule is used for hydrating the COO− side and the remaining 2.7 and 8 H2O molecules hydrate the respective alkyl group. Hence, OPr− is more hydrophobic than OAc− in terms of the hydration number. However, both alkyl moieties seem to equally retard the hydrogen bond probability of bulk H2O away from hydration shells around nonpolar sites, as much as the probing 1P does.
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INTRODUCTION We have devised a thermodynamic methodology that characterizes what non-electrolyte1,2 and electrolyte solutes3 do to the molecular organization of H2O. While a detailed description of the methodology has been given elsewhere,1−3 an account is given below. We have been advancing a differential approach in solution thermodynamics and applied it to studies of aqueous solutions.4 We experimentally determine the excess partial molar quantities of solute B in B−H2O (abbreviated as W) system. For enthalpy, for example, it is written as ⎛ ∂HE ⎞ ⎟⎟ HBE ≡ ⎜⎜ ⎝ ∂nB ⎠ p , T , n
gain a deeper insight into a complex system−aqueous solution. Accordingly, we defined and evaluated a third derivative of GE, E , as the enthalpic B−B interaction, HBB ⎛ ∂H E ⎞ ⎛ ∂H E ⎞ E HBB ≡ N ⎜⎜ B ⎟⎟ = (1 − x B)⎜⎜ B ⎟⎟ ⎝ ∂nB ⎠ ⎝ ∂x B ⎠
where xB is the mole fraction of B. In evaluating we do not use any fitting function but take the derivative of the far right of E thus obtained is completely model-free. eq 2 graphically. HBB According to the definition, eq 2, this quantity provides information about the effect of B on the enthalpic situation of B. Hence, it signifies the B−B interaction in terms of enthalpy.4 E We stress that HBE and HBB pertain only to solute B in a complex mixture. Or, we perturb the entire system by an infinitesimal addition of B, the target component, in order to isolate B’s effects on the appropriate level of thermodynamic quantity in question, HE and HBE, respectively, as eqs 1 and 2 indicate. We used these second and third derivatives and their entropy and volume analogues on binary aqueous solutions of nonelectrolytes and learned that the mixing schemes, the
(1)
W
where nB and nW are the molar amounts of B and W, respectively. HE is the total excess enthalpy of the system with the total molar amount N = nB + nW. In the (p, T, ni) variable system, HE is obtained by taking the derivative of GE once with respect to T, as HE = GE − T(∂GE/∂T) and hence HBE is a second derivative of GE. While HE (first derivative) is the excess enthalpy of the entire system, HBE (second derivative) provides information about the enthalpic contribution toward the entire system by B or the actual enthalpic situation of B in a complex mixture. Thus, the second derivative quantity gives more detailed information in general. This approach we pursued to © 2012 American Chemical Society
(2) E HBB ,
Received: December 8, 2011 Revised: February 25, 2012 Published: February 27, 2012 3571
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associated with phase transitions. The crucial difference is E that here H1P1P is a third derivative quantity, while the latter Cp is a second derivative of G associated with phase transitions. We suggest therefore that the present crossover is more subtle than the phase transition, and we call it the mixing scheme transition. We suggest further that point X is the onset and Y the end point of the mixing scheme transition.4 The composition region beyond point Y is where mixing scheme II is operative.4 We also learned that hydrophiles, in mixing scheme I, form hydrogen bonds directly to the (temporarily existing) hydrogen bond network of liquid water rather than stripping a certain number of H2O molecules from it to form hydration shells. By so doing, they do not break the hydrogen bond connectivity but they are now impurities in the hydrogen bond network. This would break the H donor/acceptor symmetry of pure liquid water and hence retard the degree of fluctuation enjoyed in pure H2O. When the solute mole fraction reaches about 0.1, the solute composition becomes too high to keep this mode of mixing scheme, and the solute starts to aggregatethe onset of mixing scheme II.4,8,9 We point out in passing that, for a E hydrophilic solute B, the xB-dependence pattern of HBB in binary B−H2O is qualitatively different from that of hydroE phobic B in that there is no peak type anomaly. Rather, HBB decreases rapidly from xB = 0 to point X, where the slope changes rather suddenly to mixing scheme II.4,9 In this work, we use hydrophobes/hydrophiles to mean the moieties that do to liquid water as described above. Another set of the second derivatives of G is the response functions, heat capacity (Cp), compressibility (κT), and thermal expansivity (αp). They are related to the mean square fluctuations of entropy and volume, and the entropy−volume cross fluctuation, respectively. The last quantity is particularly important in discussing modification to liquid water by the reasons discussed below. We defined earlier the normalized entropy−volume cross fluctuation, SVΔ, as4
molecular level mode of mixing in aqueous solutions, depend crucially on the composition. We learned that, in the H2O-rich region, the integrity of liquid H2O is maintained. Namely, while minor modification by the presence of B within a limited concentration range is operative, the bulk water remains as a highly interacting system via hydrogen bonding in addition to van der Waals interactions but also highly fluctuating system; the hydrogen bonds are forming/breaking rapidly, and yet at any given instance, the hydrogen bond network is bondpercolated.4,4 In the solute-rich region, the solute molecules cluster together in a similar manner as in the pure solute liquid. H2O molecules, in this region, interact with such solute clusters almost independently as a single molecule. In the intermediate composition region, two kinds of clusters mix together, one rich in H2O and the other in solute molecules. These two kinds of clusters are reminiscent of the H2O-rich and solute-rich regions.4 We name these qualitatively different mixing schemes as mixing schemes I, II, and III from the H2O-rich side. While the integrity of liquid water is retained in mixing scheme I, the molecular organization of bulk water is modified gradually as the solute mole fraction xB increases, the detail of which is crucially dependent on the nature of solute. Our studies on a series of mono-ols with a varying size of alkyl group revealed that the effect of a hydrophobic moiety on H2O within mixing scheme I is that a hydrophobe forms a hydration shell, the hydrogen bond probability of which is higher than pure bulk H2O similar to the classical “iceberg” formation.4 Recent theoretical works support the existence of more ordered H2O molecules in the vicinity of such hydrophobic solute as benzene and cyclohexane.6,7 More importantly, however, we realized that the hydrogen bond probability of bulk H2O away from “icebergs” is reduced progressively as xB increases. At the threshold value of xB depending on the strength of hydrophobicity, the probability of hydrogen bonding reaches the bond percolation threshold and the connectivity of the hydrogen bond network is lost, and mixing scheme II sets in.4 For a typical hydrophobe 1-propanol (1P), i.e., B = 1P, E H1P1P shows a peak type anomaly as shown in Figure 1. The top of the peak, point X in the figure, is regarded as the crossover point, just as familiar heat capacity (second) anomalies
SV
Δ ≡ RT α p/Vm
(3)
where R is the gas constant and Vm the molar volume of the system. The details about the definition and the physical significance of the normalized fluctuations are discussed extensively elsewhere.4 We then define the partial molar normalized entropy−volume cross fluctuation of solute B, a third derivative of G, as SVΔ ≡ N (∂ SV Δ/∂n ) = (1 − x )(∂ SV Δ/∂x ) B B B B
(4)
This quantity ΔB signifies the effect of solute B on the normalized entropy−volume cross fluctuation of the system. Within mixing scheme I where the integrity of liquid water is maintained, SVΔB could provide important information as to how the solute B modifies the molecular organization of water. Unlike ordinary liquids, water is unique; one such manifestation is that αp is negative below 4 °C. This came from the fact that a positive local/temporal volume variation during the course of fluctuation could be associated with a negative entropy variation due to instantaneous/local formation of hydrogen bonds. This effect is apparently dominant below 4 °C in pure water. If addition of solute at a fixed temperature reduces the hydrogen bond probability of the bulk H2O gradually as discussed above, this negative contribution to SVΔ decreases and hence the net value of SVΔ increases gradually. As a result, its mole fraction derivative, SVΔB, calculated by eq 4 should be SV
E Figure 1. H1P1P and SVΔ1P in the 1P−H2O binary system at 25 °C. See text.
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E E Figure 2. (a) Effects of “hydration centers” on the H1P1P pattern per unit increase in xS0. (b) Effects of “hydrophobes” on the H1P1P pattern per unit 0 0 E increase in xS . See text about point P and xS in the figure. (c) Effects of “hydrophiles” on the H1P1P pattern per unit increase in xS0. (d) Effect of an E pattern per unit increase in xS0. “amphiphile” on the H1P1P
positive. Figure 1 also shows SVΔ1P for B = 1-propanol (1P) against the mole fraction of 1P, x1P. As is evident in the figure, SV Δ1P is indeed positive but also it shows an anomalous peak, E .4 very much similar to the x1P-dependence pattern of H1P1P SV Indeed, on scaling the ordinate for Δ1P by a single factor with the zero point fixed, both patterns match exactly! We point out in passing that the mean square volume fluctuation calculated using the κT data indicated an initial decrease within mixing scheme I.4 This latter finding suggested that the hydration shells around 1P with a higher hydrogen bond probability (“icebergs”) do not participate in fluctuation, or the lifetime of icebergs is longer than that of fluctuation. If icebergs form as a part of fluctuation within the same time scale, the mean square volume fluctuation should increase on addition of 1P.4 Thus, SV Δ1P reflects the situation of the bulk H2O away from “icebergs”. Recent femtosecond pump−probe spectroscopic studies10−12 suggested H2O molecules in hydration shells have slower dynamics than those in bulk H2O, irrespective of whether the core is an ion10,12 or a neutral nonelectrolyte.11,12
While direct comparison between findings about dynamics and about thermodynamic situation may be dangerous, both findings are surely related. The findings about Figure 1 that the x1P-dependence pattern E of both SVΔ1P and H1P1P match exactly on scaling the ordinate by a single factor indicates that both quantities share the same fundamental cause, and we suggest that the 1P−1P interaction is operative via bulk H2O away from hydration shells where the S−V cross fluctuation is dominant. If so, in a ternary system 1P−S−H2O, where S is a test sample, the x1P-dependence E pattern of H1P1P must reflect the effect S made on the bulk H2O, whatever happened to the vicinity of S. Thus, by trying out what could be regarded as standard S’s, the effects of which on H2O are known, we could catalogue the induced changes to E the H1P1P pattern. We could then learn the effect of unknown sample S on bulk H2O relative to these chosen standards. This is the basis for what we call the 1P-probing methodology. E To reiterate, we evaluate H1P1P in a ternary system, 1P−S− H2O, where S is the sample whose effect on H2O is to be 3573
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E learned. We first experimentally determine H1P by titrating a small enough amount of 1P into the S−W system, the initial mole fraction of which is xS0 = nS/(nS + nW) before titration. We E E then evaluate H1P1P graphically from the data of H1P against x1P E with a chosen xS0. The induced change on the H1P1P pattern as a function of xS0 is sought, and compared with the same for standard S, thus identifying the characteristics of the effect of S on H2O. Figure 2 shows sketches of such induced changes by a variety of standards. Figure 2a is for S = NaF and NaCl.3 On addition E pattern is squashed to the west. This of these salts, the H1P1P indicates that on addition of these salts the available H2O E pertains on molecules for 1P to interact is reduced, since H1P1P the thermodynamic situation of 1P as stressed above. We thus interpret this westward shift to be due to hydration. From the westward shift of point X as a function of xS0, the hydration number, nH, can be evaluated for the salt. The results are nH = 7.5 ± 0.6 and 19 ± 2 for NaCl and NaF, respectively. Theoretical studies13,14 and an X-ray/neutron diffraction study15 concluded that Na+ is hydrated by 5.2 molecules of H2O and leaves the bulk water away from hydration shells unperturbed. Our application of the 1P-probing methodology on NaBr and NaI indicated nH for Na+ is 5.2−5.6.3 This observation led us to set the reference nH for Na+ to be 5.2. If we assume the complete dissociation within the range of xS0 studied, xS0 < 0.035 for NaCl and xS0 < 0.0043 for NaF,3 it follows that nH = 2.3 ± 0.6 and 14 ± 2 for Cl− and F−, respectively. This complete dissociation assumption may be acceptable, though indirectly, from the fact that the westward shift of point X is linear to xS0 in the range studied. E Furthermore, the values of H1P1P at the start of titration, i.e., at x1P = 0, and those at point X remain the same on addition of S indicate that the bulk H2O away from the hydration shell is left unperturbed by the presence of Na+, Cl−, and F−, consistent with the conclusion by the theoretical studies mentioned above,13,14 and the suggestions by the femtosecond pump− probe spectroscopic studies.10,12 We call this class of ions the hydration center. Subsequently, we found Ca2+ (nH = 6 ± 2) and NH4+ (nH = 1 ± 1) belong to this class.3 Figure 2b is for S = hydrophobes.1 For an almost equally E pattern shifts hydrophobic solute, i.e., 2-propanol, the H1P1P parallel to the west. This is only natural considering for the sake of argument the case in which 1P itself was added as S to xS0 shown in Figure 2b. Then, the subsequent titration starts at this point P, and follows the remaining course from point P up to point X, which appears as if the titration starts at x1P = 0 and follows the course shifted exactly parallel to the west. A stronger (or weaker) hydrophobe than the probing 1P like tertbutanol (or methanol), the westward shift is farther (or lesser), reflecting that nH is larger (or smaller). This westward shift of point X indicates also unavailability of bulk H2O for 1P to interact just as the case of the hydration center. In addition, the E value of H1P1P at point X shifts northward (or southward), reflecting the fact that the ability of tert-butanol (or methanol) to reduce the hydrogen bond probability of bulk H2O away from “icebergs” is more (or lesser), as we learned from earlier studies.4 Figure 2c is for S = hydrophiles, such as urea1,16 and glycerol.1,8 Hydrophiles form hydrogen bonds directly to the hydrogen bond network of H2O and keep the hydrogen bond connectivity of the network, which could be understood as the reason for the observation that point X does not show a westward shift. The southward shift apparent in the figure could
be interpreted as the retardation of the net entropy−volume cross fluctuation due to the presence of hydrophiles within the network. Indeed, this observation was instrumental in gaining the insight into mixing scheme I for the hydrophiles discussed above. Amphiphiles’ effects, Figure 2d, are a combination of those by hydrophobic and hydrophilic moieties. Their westward and southward components show the contributions of hydrophobic and hydrophilic moieties, respectively. Nonelectrolytes that are known as “osmolytes” among bioscientists that accumulate in various organisms under water stress17 were found by the present methodology to be amphiphiles with very weak hydrophobic and with equally weak hydrophilic contributions.18 Typical constituent ions of a group of low melting point electrolytes known as “room temperature ionic liquids”, on the other hand, show an amphiphilic response to these 1Pprobing studies with extremely strong hydrophobic and equally strong hydrophilic characteristics. These findings fit rather nicely into the special characteristics of ionic liquids, low melting points for ionic compounds, in particular.19−21 In a preliminary study applying the present methodology to acetate anion, CH3COO−, the results showed surprisingly a similar response as a hydrophobe!22 It seemed natural then that ions are either hydration centers or hydrophiles. As discussed above about ionic liquid ions, it is now apparent that ions could have a hydrophobic propensity. In order to investigate this point further, we here apply the same methodology to Na salts of a series of carboxylate anions: formate (OFm−), acetate (OAc−), and propionate (OPr−).
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EXPERIMENTAL SECTION
Sodium formate (Aldrich, 99.998%), sodium acetate (Aldrich, 99.995%), and sodium propionate (Aldrich, 99%) were vacuum-dried prior to use. The initial aqueous solutions were prepared using Milli-Q H2O immediately before use. 1Propanol (Aldrich, 99.7%) was used as supplied. Due care was exercised not to contaminate moisture from atmosphere by performing 1-propanol (1P) titration in a dry nitrogen atmosphere. The calorimeter used was a semi-isothermal titration calorimeter (LKB 8700) with the following modifications. The power to the resistor (nominal 50 Ω) in the glass cell was supplied by a DC constant voltage/current source (type 2555, YEW, Japan). The thermistor (nominal 2 kΩ) inserted in the cell was calibrated against an ITS-90 calibrated Pt thermometer (nominal 25 Ω, Minco Products, Inc., USA). The excess partial E molar enthalpy of 1P, H1P , was measured making sure that the temperature variation of the solution in the cell during the course of titration was within 298.15 ± 0.1 K. In this manner, E the value of H1P at the infinite dilution in H2O was within the standard value −10.16 ± 0.02 kJ mol−1.23
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RESULTS AND DISCUSSION E The raw data of the excess partial molar enthalpy of 1P, H1P , in the ternary system, 1P−S−H2O, at the initial mole fraction of the test sample S, xS0, are given in Table S1 of the Supporting Information and plotted in Figure 3. For the runs with the highest value of xS0 in each series, it was suspected to be very close to a liquid−liquid phase separation due to somewhat slow dissolution of 1P after titration, although we did not observe turbidity in the resulting solution in the cell. It is well-known that NaOAc crystals show an extensive supersaturation in 3574
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aqueous solution. While the chosen values of xS0 are well within the saturation values in binary aqueous solutions, xS0(sat) = 0.204, 0.160, and 0.157 at 20 °C for NaOFm, NaOAc, and NaOPr, respectively, we would like to be cautious about possible phase separations in the presence of 1P. Thus, we do E not use the H1P data for the highest xS0 series for further analysis. We next evaluate the 1P−1P enthalpic interaction, E , defined by eq 2 above with B = 1P. As mentioned above, H1P1P E graphically without resorting to any fitting we evaluate H1P1P function in using the second equality of eq 2.8 The results are shown in Figure 4. For the binary 1P−H2O system (xS0 = 0), we repeated the graphical differentiation three times on the same E raw H1P data. They are separately plotted in each of parts a, b, and c of Figure 4. The comparison provides an estimate in the E uncertainty in H1P1P to be ±10 kJ mol−1. From this exercise, we also estimate the error in the x1P locus of point X to be about ±0.001, as is evident in Figure 5 at xS0 = 0. Figure 4a shows that NaOFm has the hallmark of the hydration center. See Figure 2a. Since Na+ is known to be a hydration center, OFm− anion should also be a hydration center and the bulk H2O away from the hydration shells remains unperturbed. On the other hand, NaOAc and NaOPr, Figure 4b and c, behave as “hydrophobes” in the same manner as case C in Figure 2b. Namely, both form a hydration shell with an enhanced hydrogen bond probability. E Since the heights of the peak in the figures, the values of H1P1P at point X, do not change as xS0 increases within ±10 kJ mol−1, the hydrogen bond probability of bulk H2O away from the hydration shell decreases in the same degree as 1P does to H2O. Since Na+ is already known to be the hydration center that leaves the bulk H2O away from the hydration shells unperturbed, we suggest both OAc− and OPr− behave as hydrophobes toward the present methodology and also reduce the hydrogen bond probability of bulk H2O away from hydration shells in the same degree as 1P does to H2O. Figure 5 shows the x1P loci of point X against xS0. For all three salts, the xS0 dependence appears linear, indicating that apparently there is no ion pairing or aggregation, as was evident in aqueous solutions of room temperature ionic liquids based on 1-butyl-3-methylimidazolium cation.24,25 For these ionic liquids, the equivalent plots of Figure 5 showed a clear break in slope at the same mole fraction where critical aggregations of imidazolium ions were reported.26−28 If the ionpairing of the sort suggested for MgSO4 involving H2O molecules and counterions29 were to be operative for the present carboxylate ions, though controversial,30−32 there would be a similar break or at least a convex curvature upward. The slopes of Figure 5 will indicate how much H2O molecules be removed from the bulk by hydration per a unit increase in xS0, i.e., the hydration number, nH. They are 6.4 ± 0.5, 8.9 ± 0.8, and 14 ± 2 for NaOFm, NaOAc, and NaOPr, respectively. The difference among carboxylates is that NaOFm leaves the hydrogen bond status of bulk H2O away from hydration shells unperturbed, while NaOAc and NaOPr reduce the hydrogen bond probability of bulk H2O in almost the same manner as the probing 1P, as discussed above. Since the hydration number for Na+ is well established within the present methodology as 5.2,3 and the presence of Na+ does not alter the bulk H2O away from the hydration shell, the hydration numbers for carboxylate anions are 1.2 ± 0.5, 3.7 ± 0.8, and 9 ± 2 for OFm−, OAc−, and OPr−, respectively. However, there is an important qualitative difference in that OFm− does not alter the hydrogen bond status of the bulk H2O away from hydration shells, while both OAc− and OPr− reduce
E Figure 3. (a) The excess partial molar enthalpy of 1-propanol, H1P , in 1P−S−H2O at 25 °C. S is sodium formate. S = NaOFm. (b) The E , in 1P−S−H2O at 25 excess partial molar enthalpy of 1-propanol, H1P °C. S is sodium acetate. S = NaOAc. (c) The excess partial molar E , in 1P−S−H2O at 25 °C. S is sodium enthalpy of 1-propanol, H1P propionate. S = NaOPr.
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Figure 5. The x1P-loci of point X against xS0. The uncertainty in x1P at point X is estimated as ±0.001.
it in the same manner as the probing 1P does to H2O. A recent IR spectroscopic study on the hydrogen bonding to the COO− part of carboxylate anion concluded that its two proton accepting sites are cooperatively nonequivalent and, once one of the two O atoms forms a hydrogen bond with a proton donor, the second one becomes dormant.33 It was also pointed out by photoelectron spectroscopy and ab initio computational study that the first H2O molecule that hydrates the COO− part does so in a bidentate fashion.34 If so, then the present findings indicate that the H−C side of OFm− is not hydrated but the methyl and ethyl sides of OAc− and OPr− are hydrated by 2.7 ± 0.5 and 8 ± 2 H2O molecules and that the latter hydration is of a hydrophobic nature, while the hydration of OFm− has a “hydration center” characteristic for COO− only. The same conclusion is given by another recent study by dielectric relaxation spectroscopy for OFm− and OAc− that H2O hydrates to the “hydrophilic moiety” of OFm−, but the latter contains hydration around the “hydrophobic” moiety,35 though the hydration numbers are somewhat different from ours. It is indeed dangerous to compare hydration numbers among literature data, since the definitions and the methodologies and inherent assumptions in each methodology are quite different. For experimental investigations, nH for Na+ is different; our nH(Na+) is 5.2, while that by Rahman et al.35 varies from about 4.5 to to 3 as the mole fraction increases from 0 to 0.05. Consequently, their hydration number for OFm− varies from 0 to 1 and that for OAc− from 1 to 2 in the same concentration range.35 Another ab initio molecular dynamics study (being a theoretical work, no countercation is required and hence no need to worry about its hydration behavior),36 OFm−'s hydration number by their study suggests about 2.5. An X-ray and attenuated total reflectance IR study on concentrated KOFm−H2O concluded their “hydration number” for the anion varies from 4.5 to about 1.5 as the mole fraction increases from 0 to 0.4, while that for K+ varies from about 6 to less than 1 in the same mole fraction range.37 These “hydration numbers” including our nH’s are all based on different definitions and different methodologies with different assumptions, and determined in different concentration ranges. Thus, intercomparison among them is meaningless. It is
E Figure 4. (a) The 1P−1P enthalpic interaction, H1P1P , in 1P−S−H2O at 25 °C. S is sodium formate. S = NaOFm. (b) The 1P−1P enthalpic E , in 1P−S−H2O at 25 °C. S is sodium acetate. S = interaction, H1P1P E , in 1P−S−H2O at NaOAc. (c) The 1P−1P enthalpic interaction, H1P1P 25 °C. S is sodium propionate. S = NaOPr.
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interesting to note, however, that they are all within the same order of magnitude. E The fact that the H1P1P pattern in Figure 4b and c for OFm− − and OPr shifts to the west in a parallel manner keeping the peak height almost the same indicates that the hydrogen bond probability of bulk H2O is reduced as much as the probing 1P does to H2O, as stressed repeatedly above. The latter fact is surprising, and its origin must be studied further. We point out that the hydration number of methanol by the same methodology is 3.4,1,2 as compared with 3.7 for OAc− including E nH for the COO− side, but the H1P1P value of point X decreases (southward shift), indicating that the degree of reduction of the hydrogen bond probability by methanol is weaker than that of the probing 1P. We have not yet applied the same methodology for ethanol, which might be a starting point toward further understanding. Such a project is under way. Thus, we support the conclusion that OFm− acts as what we call a “hydration center”, while OAc− and OPr− as what we call a “hydrophobe” which forms a hydration shell and reduces the hydrogen bond probability of bulk H2O away from hydration shells, in addition to a “hydration center” of the COO− side . The hydrophobicity is stronger for OPr− in terms of nH, but the degree of reduction of the hydrogen bond probability of bulk H2O is the same for both as that of the probing 1P. In closing, we comment on recent criticisms or controversies38−41 about “iceberg” formation by hydrophobic solutes. While these and other studies10−12 support the existence of a hydration shell that has a slower dynamics than that of bulk H2O, they conclude that the hydrogen bond probability of H2O in the hydration shell is not particularly enhanced. However, the samples that were used as an example of a hydrophobe for all these studies are either tetramethyl urea or trimethyl-N-oxide or both. On the other hand, our recent study using the present methodology indicated that both these compounds do not behave as a typical hydrophobe but as an amphiphile with a very weak hydrophobicity and equally weak hydrophilicity.2 Namely, the methyl groups attached to the N atom do not behave as a hydrophobic moiety toward the present 1P-probing methodology.
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ASSOCIATED CONTENT
S Supporting Information *
Table S1 containing the excess partial molar enthalpy of 1E propanol, H1P . This material is available free of charge via the Internet at http://pubs.acs.org.
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AUTHOR INFORMATION
Notes
The authors declare no competing financial interest.
ACKNOWLEDGMENTS We thank the Ministry of Education, Science and Culture of Japan for financial support. Contribution No. 23 from Research Center for Structural Thermodynamics, Graduate School of Science, Osaka University.
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REFERENCES
(1) Koga, Y. Solution Thermodynamics and Its Application to Aqueous Solutions: A Differential Approach; Elsevier: Amsterdam, The Netherlands, 2007; Chapter VII. (2) Koga, Y.; Westh, P.; Nishikawa, K.; Subramanian, S. J. Phys. Chem. B 2011, 115, 2995. (3) Reference 1 but chapter VIII. 3577
dx.doi.org/10.1021/jp2118407 | J. Phys. Chem. B 2012, 116, 3571−3577
