J. Phys. Chem. B 2009, 113, 759–766
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Solvation of Nitrophenol Isomers: Consequences for Solute Electronic Structure and Alkane/Water Partitioning William H. Steel,† James B. Foresman,† Daniel K. Burden,‡ Yuen Y. Lau,§ and Robert A. Walker*,| Department of Physical Sciences, York College of PennsylVania, York, PennsylVania 17405, and Department of Chemistry, UniVersity of Wisconsin, Madison, Wisconsin 53706, Department of Chemistry and Biochemistry, UniVersity of California, Los Angeles, California 90095, and Department of Chemistry and Biochemistry and Chemical Physics Program, UniVersity of Maryland, College Park, Maryland 20742. ReceiVed: June 12, 2008; ReVised Manuscript ReceiVed: August 27, 2008
Solute partitioning across a variety of alkane/aqueous interfaces was examined as a function of solute and alkane solvent structure. Solutes include p-nitrophenol (PNP), 3,5-dimethyl-p-nitrophenol (3,5-DMPNP), and 2,6-dimethyl-p-nitrophenol (2,6-DMPNP), the latter two being isomers distinguished solely by the location of methyl substituents on the aromatic ring. The alkane solvents included cylohexane, methylcyclohexane, octane, and iso-octane (2,2,4-trimethylpentane). PNP partitioned preferentially into the water by factors as high as 160:1. The dimethyl isomers partitioned more equally between water and the different alkanes. 2,6DMPNP showed a 3-fold greater affinity for the alkane phase than 3,5-DMPNP. Ab initio calculations were used to characterize the molecular and electronic structure of the three solutes and to quantify individual contributions to each solute’s solvation energy in model aqueous and alkane phases. Differences between 2,6-DMPNP and 3,5-DMPNP partitioning are interpreted based on the ability of the methyl groups in 2,6DMPNP to weaken hydrogen bonding between the phenol group and adjacent water molecules. This diminished solvation interaction reduces the barrier to solute migration into the nonpolar organic phase despite the fact that 2,6-DMPNP has a larger (calculated) permanent, ground-state dipole than 3,5-DMPNP. 1. Introduction Alkane/aqueous, liquid-liquid systems serve as important models for many phenomena, including membrane permeability, solvent extraction, and colloid stability.1-6 Given the disparate nature of the two bulk phases, solutes typically show very different alkane/aqueous partitioning affinities, depending on solute molecular structure. Although partitioning reflects an equilibrium distribution that minimizes solute free energy between the alkane and aqueous phases, distinguishing the molecular forces that move a solute from one phase into another presents challenges. Solute partitioning depends upon numerous effects, including individual solute-solvent interactions, the creation and destruction of solute cavities in the two phases, and the restructuring of solvent around a solute once it has migrated from one phase to the other.7-10 These effects are more complicated than studies of solute solubilities (in a single phase)11-13 because equilibrium solute concentrations in partitioning experiments are determined by differential solvation effects between the two solvents.8 Sometimes these differences are self-evident, such as differences in hydrogen bonding opportunities between aqueous and aprotic organic solvents, but questions concerning changes in overall system entropy and free energy can be subtle and on the order of kT at ambient temperatures. Nevertheless, a clear understanding of how solvation properties promote or inhibit solute partitioning marks an essential prerequisite for developing and improving models * To whom correspondence should be addressed. E-mail: rawalker@ umd.edu. † York College of Pennsylvania. ‡ University of Wisconsin. § University of California. | University of Maryland.
describing solute phase transfer in physiological, environmental, and industrial applications.14-21 The work described below represents our initial efforts to identify the roles played by solute structure and solute solvation in controlling partitioning behavior between aqueous solutions and different alkanes. One general application of alkane/aqueous systems is their ability to mimic hydrophilic/hydrophobic boundaries found in nature.22,23 As a consequence, solute partitioning across an alkane/aqueous interface has direct bearing on issues related to public health and environmental safety. For example, bioaccumulation describes how pollutants become more concentrated in organisms higher up in a food chain. Pollutants that are more soluble in hydrophobic environments are strong candidates to accumulate inside of food sources and become significant health threats compared to solutes that remain solvated in water or preferentially adsorb to particulates and inorganic substrates.24,25 A second application where partitioning plays an important role is in determining membrane affinity. Potential oral drugs must undergo extensive testing to determine if they will actively partition into the blood stream from the intestines and from the blood into the hydrophobic region of a cell membrane.26 Such testing can be done rigorously using natural membranes (such as cells from intestinal walls) or through less labor intensive, empirical methods referred to as “high-throughput screenings” (HTS).2 The log P scale describing solute partitioning between water and 1-octanol is an example of one such HTS.6,27 In recent years numerous experiments have clarified how solvent and solute structure influence solvation at interfaces,28-31 but these results do not speak to the balance of forces that lead to the equilibrium distribution of solutes between two bulk phases. Differences in solvent-solute interactions as well as configurational changes in the two solvents can lead to
10.1021/jp805184w CCC: $40.75 2009 American Chemical Society Published on Web 12/30/2008
760 J. Phys. Chem. B, Vol. 113, No. 3, 2009 significant differences in solute partitioning in different liquid/ liquid systems, even when intuition based on “like dissolves like” would predict similar behavior.32 For example, the polar solute p-nitrophenol shows an alkane/water partitioning coefficient of 0.009 when the alkane is cyclohexane, but this equilibrium constant increases more than 2-fold (to 0.025) when the alkane is n-hexadecane.29 This result may seem surprising given that the two alkane solvents have similar static dielectric constants (2.04 and 2.06, respectively) and similar densities (0.778 g/cm3 and 0.773 g/cm3).33 However, n-hexadecane is a flexible, linear alkane, capable of rearranging itself to create a solute cavity, limiting the entropic penalty of accommodating the solute, and optimizing noncovalent solvent-solute interactions. In contrast, cyclohexane only has one structural degree of freedom (“chair” to “boat” isomerization) and is less able to accommodate a migrating solute. Solvents studied in the experiments described below are all relatively small alkanes, including cyclohexane, methyl-cyclohexane, octane,and isooctane (2,2,4-trimethylpentane). Generally, such systems are treated with a classical, mole-fraction approach to solute partitioning.7,8,10,34-36 Experiments described below examine the aqueous/alkane partitioning behavior of three related solutes, p-nitrophenol (PNP), 3,5-dimethyl-p-nitrophenol (3,5-DMPNP), and 2,6dimethyl-p-nitrophenol (2,6-DMPNP). The latter two solutes are isomers differing only in the relative positioning of the two methyl groups on the aromatic ring. Nitrophenols are widely encountered in nature and have also been associated with phytotoxicityandreducedreproductiveratesinsimpleorganisms.37,38 Given the potential harmful effects of these species, models used to estimate steady-state concentrations and to identify resevoirs where nitrophenols can accumulate require accurate information about analyte partitioning. Complementing the experimental studies are a series of high-level, ab initio calculations intended to identify those contributions to overall solvation energies (and differences in solvation energies) that are responsible for solute partitioning between aqueous and alkane phases. The goal of this comparison is to move beyond simple functional group considerations and begin to identify those molecular properties that control solute partitioning between weakly associating liquid phases. PNP shows the smallest (oil/water) partitioning constants with Kalkane/water ≈ 0.01 for all of the alkane-aqueous systems studies. Much more surprising are the Kalkane/water constants for 2,6-DMPNP that are ∼3-fold greater than those for 3,5-DMPNP despite the fact that the 2,6 isomer has a larger calculated dipole moment. Simple dipolar considerations would predict that the solute having a larger dipole should show a stronger preference for the polar, aqueous phase. Ab initio calculations suggest that the counterintuitive observations depend in large part on the ability of the methyl groups in 2,6-DMPNP to weaken hydrogen bonding between the phenol group and adjacent water molecules, thereby destabilizing solvation in the aqueous phase and reducing the barrier to solute migration from the aqueous phase into the alkane. The combination of experimental results coupled with ab initio calculations provides benchmarks and insight for evolving models of solute partitioning that need to predict differences in partitioning behavior with chemical accuracy (∼1 kcal/mol). 24,39-42 2. Experimental Considerations 2.1. Materials. The structures of solutes studied in the partitioning experiments are shown in Figure 1a. These solutes were chosen primarily because they share similar functional group compositions. Both PNP and 2,6-DMPNP were purchased
Steel et al. from Aldrich and were used as received. The purities of PNP and 2,6-DMPNP were reported as 99+% and 98%, respectively. 3,5-DMPNP was synthesized and purified according to previously published procedures.22,43 Briefly, 3,5-dimethylphenol (3,5-DMP) was nitrated with nitric acid. Then, 610.8 mg of 3,5-DMP was dissolved into excess diethyl ether under N2. While the dissolved 3,5-DMP solution was being stirred in a flask over an ice bath, a total of 0.3125 mL of 3.2 M nitric acid was slowly added to the solution. After 90 min the reaction had gone to completion and excess ether was removed with a rotary evaporator. The product mixture was then dissolved in chloroform and extracted three times with water. The product mixture was separated using several iterations of flash chromatography with a mobile phase of methylene chloride. Several impurities were separated from the 3,5-DMPNP, including the ortho isomer of 3,5-DMPNP, 3,5-dimethyl-orthonitrophenol; 3,5-dimethyl-4,6-dinitrophenol, where nitro groups has been added to both ortho and para positions on the ring; and leftover reactant, 3,5-DMP. Additional products remained on the column and were unable to be identified. Yields of 20% were slightly less than the reported 25%.22,43 The melting point of synthesized 3,5-DMPNP was 105-108 °C, in agreement with previously reported values of 106-108 °C.22,43 NMR spectroscopy and mass spectroscopy also confirmed product purity. The alkanes used in the partitioning experiments were cyclohexane, methylcyclohexane, octane, and iso-octane. They were purchased from J.T. Baker, Sigma-Aldrich, Sigma-Aldrich, and Fisher Scientific, respectively. The solvents were used as received, and their purities were reported as 99.9, 99, 99.8, and 98% respectively. Analysis with mass spectrometry, NMR spectrometry, and (liquid/liquid) interfacial tension measurements failed to identify any specific impurities. These alkanes all have static dielectric constants of 2.0 ( 0.05. The UV transparency of all solvents down to 250 nm allowed solute concentrations to be determined accurately using UV-vis absorbance spectrometry. 2.2. Experimental Characterzations. Solute partitioning in these experiments is described through the ratio of solute concentration in the alkane relative to that in the aqueous phase:
Keq )
[solute in organic phase] [solute in aqueous phase]
(1)
Concentrations of solutes are determined by a Beer’s Law analysis. Such a technique requires calibration curves covering a wide range of solute concentrations. Serial dilutions were carried out six times to create seven solutions with target concentrations of 1.00, 0.500, 0.250, 0.125, 0.063, 0.031, and 0.016 mM. A linear least-squares regression line yielded a slope equal to the molar absorption coefficient of the solvent in the solute. The accuracy of the calibration curves is reflected in the y-intercept of the linear regression. Any deviations from ideality were minimal, resulting in uncertainties in concentration of less than 5%. The linear dependence of absorbance on concentration over 2 orders of magnitude enabled calculations to assume activity coefficients of unity. To determine the partitioning coefficient, Keq, at 295 K, alkane/aqueous systems were assembled in the following manner: first, 50 mL solutions of a known concentration (0.250 mM or 0.500 mM) were created in both water and the organic solvent being tested. Each solution was transferred to its own separatory funnel. Next, 50 mL of the complementary solvent was added to each separatory funnel. Then, 50 mL of pure water
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Figure 1. Molecular structures of the solutes used in these studies (top) and a comparison between HOMO (middle) and LUMO (bottom) electron densities of the different solutes. Electronic structure was calculated using methods described in the text.
was added to the organic solution, and 50 mL of pure organic solvent was added to the aqueous solution. The systems were then agitated and allowed to equilibrate for 24-72 h. Following equilibration, aliquots were taken from the aqueous and organic layer of each system. The absorbance of each aliquot was measured and converted to a corresponding bulk solution concentration. The solute concentration in each sample was calculated using the extinction coefficient from the corresponding calibration curve. Care was taken to ensure that all of the solute could be accounted for in the two phases based on concentrations determined from the absorbance measurements. One potential sink for solutes was acid-base equilibrium leading to the formation of phenoxide anions in the aqueous phase. To prevent the formation of phenoxide anions from the parent phenols in aqueous solution, all aqueous phases were acidified. The acidified solutions were made from deionized water (Millipore Milli-Q Water, >18 MΩ · cm) and hydrochloric acid and had a pH of ∼3.5. In solutions made from deionized water that hadn’t been acidified (pH ∼5.3-5.6), the presence of phenoxide was readily apparent both visibly (as a pale yellow tint to the aqueous phase) and in the UV absorbance spectra (as a pronounced absorbance centered at ∼420 nm). With pKa values between 7.1 and 7.4, neutral phenol solutes were present in excess of 1000:1 in acidified aqueous solutions, ensuring accurate concentration measurements. 2.3. Computational Methods. Computations were performed using Gaussian03.44 Structures for the three solutes were optimized in the gas-phase using the MP2 method and the
6-311++G(2d,p) basis set. Solvent effects were investigated two different ways. First, solvation free energies were estimated in both cyclohexane and water using the standard PCM UAHF model45 applied to the MP2 gas-phase geometries. Here the electrostatic contribution is evaluated at the MP2/ 6-311++G(2d,p) level of theory. Second, MP2 optimizations were perfomed on water-solute complexes (hydroxy hydrogen to water oxygen) in order to estimate hydrogen bond binding enthalpies. Thermal contributions were obtained from the 0.988 scaled B3LYP/6-311++G(2d,p) frequencies, and the counterpoise method was used to correct for basis set superposition error. We also examined the effect of rotating the nitro group with respect to the plane of the aromatic ring in both 4-nitrophenol and 3,5-dimethyl-4-nitrophenol. Constrained optimizations were performed using the B3LYP/6-311+G(2d,p) model at 10 degree increments starting from a CCNO dihedral angle of zero and going to ninety degrees. These geometries were then used in two ways. First, MP2/6-311++G(2d,p) single point calculations were done at each point to compute the gas-phase barrier to rotation. Next, TDDFT calculations in conjunction with the B3LYP/6-311+G(2d,p) model were used to calculate the energies of the lowest three excited states of these two solutes as a function of nitro twist. This model was also used to obtain the dipole moments of the excited states at the MP2 equilibrium structures by using numerical differentiation of an applied electric field.
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Figure 2. Torsional barriers (a) and ground-state dipole moments (b) for PNP and 3,5-DMPNP. Data for 2,6-DMPNP tracks results for PNP quite closely. Note the asymmetry in the torsional potential for 3,5DMPNP as θ approaches 90°.
3. Results and Discussion 3.1. Gas-phase Equilibrium Structures. Calculated structures of the three solutes are shown in Figure 1a. PNP and 2,6DMPNP both have the nitro group coplanar with the aromatic ring, whereas the adjacent methyl groups of 3,5-DMPNP force the nitro group into a conformation that is twisted ∼55° relative to the plane of the ring. This orientation prevents the nitro group of 3,5-DMPNP from participating in the resonance structure of the ring and has a very pronounced effect on overall solute electronic structure. The HOMOs and LUMOs for all three solutes are shown in Figure 1. The electronic structure of PNP and 2,6-DMPNP are quite similar with the LUMO of each solute, reflecting some of the charge transfer character that contributes to the S1 excited state. In contrast, the LUMO of 3,5-DMPNP keeps the electron density of the nitro group decoupled from the π-structure of the ring. On the basis of the series of gas-phase calculations, the methyl groups in the 3- and 5- positions also reduce the corresponding barrier to -NO2 internal rotation. In 3,5-DMPNP the barrier to rotation is asymmetric and is generally less than 1 kcal/mol unless the -NO2 is nearly coplanar, which is a geometry that lies ∼5 kcal/mol higher in energy than the minimum energy configuration. In contrast, the barrier to -NO2 rotation in PNP (and 2, 6-DMPNP, data not shown) is symmetric and rises to nearly 6 kcal/mol when θ ) 90°, despite the absence of steric constraints (Figure 2a). Twisting the -NO2 group also leads a reduced ground-state dipole (3.78 D, calculated) for 3,5DMPNP, compared to those of PNP (4.82 D, calculated) and 2,6-DMPNP (5.43 D, calculated), where the -NO2 group is coplanar with the aromatic ring (Figure 2b). We note that the ground-state dipoles of PNP and 2,6-DMPNP have been measured in solution to be 4.8 and 5.0 D, respectively.46-48 Although these experimental dipole values do not exactly match the gas-phase, calculated values, we are nevertheless confident in the relative ordering of dipoles (µ3,5-DMPNP < µPNP < µ2,6-DMPNP) as well as the approximate differences in magnitudes. Measurements of the solutes’ electronic structures reported below bolster this confidence. 3.2. Solvatochromic Activity. Electronic Structure. One goal of these studies is to identify how different solute-solvent
Figure 3. (a) Experimental solvatochromic data and (b) calculated electronic excitation energies for PNP and 3,5-DMPNP as a function of -NO2 torsion angle. For the experimental data in panel (a), representative solvents (and their corresponding Onsager values) include octane (0.39), isooctane (0.39), cyclohexane (0.41), methylcyclohexane (0.41), diethyl ether (0.69), chloroform (0.72), 1-octanol (0.85), 1-butanol (0.92), methanol (0.95), DMSO (0.97), and water (0.98). Arrows in panel (b) point to the equilibrium -NO2 orientation (relative to the plane of the ring) for both PNP (blue) and 3,5-DMPNP (red).
interactions correlate with observed partitioning behavior in alkane/aqueous systems. A property that provides insight into solute/solvent interactions is a solute’s solvatochromic response in different solvents.49-51 Solvatochromism refers to a solute’s solvent-dependent absorbance and/or emission spectra, and this property can be exploited to quantify local polarity around the solute30,34,52,53 or to calculate differences in permanent dipoles between a solute’s ground and excited electronic states.49-51,54 Figure 3 shows bulk solution absorbance maxima in a variety of solvents as a function of each solvent’s Onsager polarity function, f(), where f() is related to a solvent’s static dielectric constant, ,55
f() )
2( - 1) 2 + 1
(2)
According to the data shown in Figure 3, the three solutes show two distinctive types of solvatochromic behavior: the electronic transitions of both PNP and 2,6-DMPNP depend sensitvely on solvent polarity, with λmax shifting approximately 30 nm to longer wavelength as solvent polarity increases from that of alkanes ( ≈ 2.0) to that of water ( ≈ 80). In contrast,
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TABLE 1: Calculated Gas-phase Dipole Moments and Experimental Partitioning Coefficients Partitioning Results (Along with Calculated Gas-phase, Ground and Excited State Dipoles) for PNP, 2,6-DMPNP, and 3,5-DMPNP in Four Alkane/Water Systemsa solute
µgs(MP2) (D)
µgs(DFT) (D)
µex(TD-DFT) (D)
cyclohexane/water
methylchex/water
PNP 2,6-DMPNP 3,5-DMPNP
4.82 5.43 3.78
5.46 6.12 4.13
11.03 11.85 7.98
0.009 ((0.001) 1.47 ((0.01) 0.38 ((0.03)
0.010 ((0.001) 1.63 ((0.02) 0.47 ((0.02)
octane/water 0.0080 ((0.0008) 1.31 ((0.05) 0.37 ((03)
iso-octane/water 0.0060 (( 0.0002) 0.855 ((0.002) 0.34 ((0.01)
a
Results describe the ratio of solute concentration in the organic phase to that in the aqueous phase. Results are reported as the averages of at least four independent measurements.
the response of 3,5-DMPNP in this same collection of solvents is attenuated with λmax shifting only 10 nm between the alkane (344 nm) and aqueous (355 nm) limits. Also it is worth noting that the excitation energies of PNP and 2,6-DMPNP overlap closely in a given solvent, the excitation energy of 3,5-DMPNP in the same solvent is shifted more than 30 nm to longer wavelength. These differences in solvatochromic behavior can be related to each solute’s ground and excited-state electronic structure. For its electronic excitation to shift to lower energy (or longer wavelength), a solute must have a larger dipole in its excitedstate relative to its electronic ground state. TD-DFT calculations performed on the three solutes (Table 1, columns 3 and 4) provide a comparison of the ground and excited-state dipole moments that supports the experimental observations. More specifically, the magnitude of the calculated dipole change provides insight into the observed differences between 3,5DMPNP and the other two solutes, PNP and 2,6-DMPNP. The differences in permanent dipole moments between the ground and first excited states, ∆µ, are ∼+5.5 D for both PNP and 2,6-DMPNP but only +3.9 D for 3,5-DMPNP. For both PNP and 2,6-DMPNP, these values compare reasonably well to experimental values reported in the literature.46-48 Given solute equilibrium structures, we also calculated the excitation energies of PNP and 3,5-DMPNP in a polarizable continuum and reproduced qualitatively the experimental results (Figure 3b). An interesting observation is that if the -NO2 group of PNP were twisted out of plane, then one would expect a large jump in the aborbance maximum due to a different excited state (that has less CT character) becoming more dominant. Combining the results from ab initio calculations with experimental data shows that the twisted nitro group alters the electronic structure in 3,5-DMPNP, which is a change that results in the reduction of the dipole moment of 3,5-DMPNP and the corresponding change in dipole moment upon excitation. This result is supported by the second trend observed in Figure 3a, namely that the solvatochromic window of PNP and 2,6DMPNP (∼30 nm) is enhanced relative to that of 3,5-DMPNP (10 nm). Such an enhancement is a reasonable consequence of a larger change in dipole moment being more sensitive to solvation in polar solvents such as water and DMSO. These interactions would lead to differentially stronger solvent-solute interactions in a solute’s excited-state and a correspondingly larger decrease in electronic transition energy. 3.3. Aqueous/Alkane Partitioning of Nitrophenols. Although the solvatochromic results might lead one to anticipate that the less polar 3,5-DMPNP should show a stronger affinity than either PNP or 2,6-DMPNP for solvation in nonpolar, organic phases, partitioning coefficients indicate otherwise. Partitioning results clearly show that other structural and electronic properties of solutes must be considered in order to accurately predict solvation preferences. The room temperature (22 ( 1 °C) alkane/aqueous partitioning coefficients of PNP,
2,6-DMPNP, and 3,5-DMPNP were determined for four different alkane/water systems. Cyclohexane, methylcyclohexane, octane, and iso-octane were used as the organic phases in the alkane/aqueous systems. Data for all 12 experiments are reported in Table 1. Two general trends stand out: (1) 2,6-DMPNP is the most hydrophobic of the three solutes, with partitioning ratios that are ∼3 fold larger than those of 3,5-DMPNP for a given alkane/aqueous system (and ∼150-times larger than the partitioning ratio of PNP); and (2) methylcyclohexane is the alkane that is best able to accommodate the nitrophenol-based solutes, as evidenced by the largest partitioning coefficients; isooctane is the least accommodating solvent. For a given solute, the partitioning coefficient can vary by up to a factor of 2 between these two aqueous/alkane systems. Between these two limits, cyclohexane/aqueous systems have consistently larger partitioning ratios than octane/aqueous systems. Solute Dependence of Partition Coefficients. As expected, PNP shows the smallest partition coefficients in all alkane/ aqueous systems. The polar -OH and -NO2 groups and the large ground-state dipole (4.82 D) make this solute very hydrophilic. The partition ratios of 0.6 to 1.0 × 10-2 (depending on the specific alkane/aqueous system) compare very favorably to literature values reported for similar organic/aqueous systems.56-58 Partition coefficients are often predicted by a solute’s functional group composition.14-16,19,59,60 However, differences between the partition coefficients for 2,6-DMPNP and 3,5DMPNP need to be interpreted not based simply on functional group composition, but rather on the basis of how solute molecular structure influences solvent-solute interactions. From data in Table 1, 2,6-DMPNP is 3-4 times more soluble in the organic phase relative to water compared to 3,5-DMPNP. This result is surprising given that 2,6-DMPNP has a significantly larger ground-state dipole (5.43 D) at its equilibrium geometry compared to the equilibrium ground-state dipole of 3,5-DMPNP (3.78 D). On the basis of the simple adage “like dissolves like”, one would expect 2,6-DMPNP to have smaller, not larger, alkane/water partitioning coefficients compared to 3,5-DMPNP. The methyl groups of 2,6,DMPNP do not interfere with the structure of this isomer’s nitro group, nor do they alter the solute’s electronic structure significantly from that of PNP. However, the methyl groups do influence the ability of the -OH group to interact with surrounding solvent. Specifically, the relatively bulky methyl groups on either side of the -OH can prevent the solute from forming strong hydrogen bonds with adjacent water molecules. Table 2 summarizes the results of ab initio calculations that determined ∆Hform of the hydrogen bond between an explicit water molecule and the nitrophenol hydroxyl group. These experiments were carried out for nitrophenol-water systems in the gas phase, but the results are consistent with predictions based on the determination of isomer solvation energies in different polarizable continuum systems (vide infra). The important results from these calculations are
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TABLE 2: Solute-Water Binding Energies [kcal/mol]a theory
pnp
3,5-dimethyl-pnp 2,6-dimethyl-pnp
MP2/6-311++G(2d,p) 5.87
5.35
4.33
a
This quantity reports the negative of the enthalpy change for the formation of the gas-phase solute-water complex from the solute and water (estimated at 298K). Geometry is optimized at the MP2 level. BSSE corrections are applied using counterpoise. Thermal contributions are from the 0.988 scaled B3LYP 6-311++G(2d,p) frequencies.
that the water molecule is more weakly bound to the hydroxyl group of 2,6-DMNP (∆Hform ) -4.33 kcal/mol) than it is to either PNP (-5.87 kcal/mol) or 3,5-DMPNP (-5.35 kcal/mol). This weaker hydrogen bonding reduces (i.e., makes more positive) the solvation energy of 2,6-DMPNP in aqueous solution and lowers the barrier to migration from the aqueous to the organic phase. Such effects have been observed previously in systems where intramolecular hydrogen bonding make certain isomers appear more lipophilic than than isomers incapable of such interactions.57-59,61-63 However, we propose that for 2,6DMPNP the effects of the methyl groups on alkane/water partitioning are not intramolecular but rather intermolecular. The methyl groups in 2,6-DMPNP weaken solvation interactions with water and in doing so promote solvation in the nonpolar alkane phase. These effects were explored in two ways. First, from ab initio calculations, we deconvolved each solute’s solvation energies in both polar (water) and nonpolar (cyclohexane) media into electrostatic, cavitation, dispersion, and repulsion interactions.45 The results appear in Table 3. Electrostatic contributions arise from the favorable interactions of a solute’s charge distribution and the continuum dielectric of the solvent. Cavitation terms are due to the energy cost of organizing a shell of solvent molecules around the solute. Dispersion and repulsion components arise from the intermolecular forces between any two molecules. Empirical parameters from the solvent are used to define all but the electrostatic pieces of the complete Hamilitonian for the solvated system. The differences in calculated solvation free energies (∆(∆Gsolv)) track the observed differences in solute partitioning between alkane and aqueous phases. PNP has the largest difference (∆(∆Gsolv) ) +5.7 kcal/mol), favoring higher concentrations in the aqueous phase. 2,6-DMPNP has the smallest difference (2.6 kcal/mol, again favoring partitioning in the aqueous phase), meaning that, of the three solutes, 2,6DMPNP should be the most soluble in alkanes and should show the largest alkane/aqueous partitioning coefficients. We note that the absolute magnitudes of these solvation energy differences will not be quantitatively accurate given that PCM calculations necessarily neglect all issues related to solvent structure, but the trends predicted by these calculations agree qualitatively with the results summarized in Table 1. Closer inspection of the calculated data show that these differences in ∆Gsolv result primarily from the relatively small (in magnitude) solvation
Figure 4. Schematic representations of a solute’s change in solvation energy as it migrates from an aqueous phase into the organic phase. The energy axis is arbitrary, but the relative separations between solvated solutes quantitatively reflect the differences reported in Table 3. PNP has the largest difference in solvation energies between water and alkanes, and 2,6-DMPNP has the smallest. The small energetic minimum between the two equilibrated states (aqueous and organic) indicates that all three solutes are surface active and create small surface excesses at the aqueous/alkane interface.
energy of 2,6-DMPNP in water rather than differences in solute solubilities in the alkane continuum (that differ by less than 0.5 kcal/mol). Schematically, these efffects are illustrated in Figure 4. SolWent Dependence of Partition Coefficients. Despite the data showing clear patterns of general behavior, interpreting solvent-dependent differences in partitioning coefficients of the same solute is not straightforward. Of the alkane solvents tested in these experiments, methylcyclohexane was the most accommodating solvent (with the largest Kalkane/water equilibrium constants), and isooctane was the least accommodating solvent. The partitioning coefficients for each solute were approximately 1.5-fold larger in the methylcyclohexane/aqueous systems than in the isooctane/aqueous systems. Between these two limits, the cyclohexane/aqueous partitioning coefficients were always larger than those for the aqueous/octane systems. Beyond these empirical observations, quantitative interpretation of the results is difficult. What follows is a brief discussion of solvation properties that may influence the solvent dependence observed in the alkane/aqueous partitioning experiments. From the standpoint of a polarizable continuum model, all four organic solvents should solvate a given solute equally well. In reality, however, numerous studies of solute solubility and partitioning have shown that the identity of the alkane solvent can affect the solvent’s ability to solvate a solute.34,56,61,64-66 Results reported in Table 1 are consistent with previous reports noting that partitioning coefficients of substituted phenols are significantly larger for cyclohexane/aqueous systems compared to linear alkane/aqueous systems, although the origins of these
TABLE 3: Contributions to the Free Energy of Solvation [kcal/mol] MP2/6-311++G(2d,p) Optimized Gas-phase Geometries Solvation Model: HF 6-31G(d) PCM UAHF pnp
3,5-dimethyl-pnp
2,6-dimethyl-pnp
term
water
cyclohexane
water
cyclohexane
water
cyclohexane
electrostatic cavitation dispersion repulsion ∆Gsolv
-11.22 17.80 -18.66 3.40 -8.68
-2.21 13.51 -17.17 2.85 -3.02
-10.83 21.82 -22.48 4.04 -7.45
-2.01 16.50 -20.67 3.38 -2.80
-8.49 22.13 -22.33 3.50 -5.19
-1.74 16.70 -20.54 2.93 -2.65
Solvation of Nitrophenol Isomers
J. Phys. Chem. B, Vol. 113, No. 3, 2009 765
TABLE 4: Geometric Properties of the Alkane Solventsa Molar ∆(∆G)FH ∆(∆G)FH ∆(∆G)FH volume (PNP) (2,6-DMPNP) (3,5-DMPNP) (cm3/mole) methylcyclohexane cyclohexane octane isooctane
127.7 107.8 162.6 165
+0.07 +0.20 +0.09 +0.28
-4.18 -4.01 -4.19 -3.93
-3.42 -3.19 -3.42 -3.36
a Molar volumes represent reported experimental values. ∆(∆G)FH data were calculated using the equilibrium partitioning coefficients reported in Table 1 and solute molar volumes of 94.0 cm3/mole for PNP, 139 cm3/mole for 2,6-DMPNP, and 138 cm3/mole for 3,5-DMPNP. The PNP molar volume is an experimentally determined value. Molar volumes for 2,6 and 3,5-DMPNP were extrapolated from calculated quantities. ∆(∆G)FH quantities reported in kcal/mole.
results have not been correlated with explicit solvent structure.56,65 An analysis of the contribution made to ∆(∆Gsolv) by a Flory-Huggins volume dependence on solvation predicts that the largest organic solvent should favor larger (alkane/aqueous) partitioning coefficients through the following expression: 32
∆(∆G)FH ) -RTln
(
[x]alk Vx Vx - RT [x]aq Vaq Valk
)
(3)
where ∆(∆G)FH is a Flory-Huggins-based description of the change in solvation energy when a solute migrates from one phase to another, [x]i is the concentration of the solute in the alkane and aqueous phases, and Vi are the molar volumes of the solute, water, and organic species. Given that Vaq and Vx are constant for a given solute (and greater than unity), and given that Vaq will be much smaller than Valk, the largest alkane solvents will make the smallest (positive) contribution to ∆(∆G)FH. Such differences are apparent when comparing the aforementioned partitioning of PNP between water and cyclohexane compared to the water/hexadecane system.29 Testing the FH description of nitrophenol partitioning in the alkane/aqueous systems required evaluating the geometric properties of the four organic solvents and three solutes used in these studies. Results are shown in Table 4. Solvent molar volumes were determined from reported solvent densities (at 295 K). Several trends are apparent in the reported results. First, the FH description predicts that iso-octane should be the least accommodating solvent (with the most positive ∆(∆G)FH) and that methylcyclohexane should be more accommodating than cyclohexane. However, the FH description also predicts that methylcyclohexane and isooctane should show similar partition coefficients, regardless of solute identity. Instead, experimental data show octane and cyclohexane to behave similarly. 4. Conclusions To account for the solvent-dependent differences in observed partitioning, we propose that one should consider instead the individual contributions to a solute’s solvation energy in the different organic solvents. These terms include attractive electrostatic and dispersive components (leading to negative ∆Gsolv) as well as terms describing the energy required for cavity formation and a “repulsive” term that keeps solvent out of the newly created cavity.45 These latter two termsscavity formation and repulsive interactionsslead to positive contributions to ∆Gsolv. Inspection of these quantities for the different isomers embedded in polarizable continuum models (Table 3) shows
that a systematic analysis of each term individually will be difficult if the solvent molecules must be treated explicitly. Results from calculation will need to be able to identify small, quantitative differences between contributions that are much smaller in magnituge than the contributions themselves. For example, observed differences in partitioning imply that 3,5-DMPNP is up to ∼2-times more soluble in methylcyclohexane than in isooctane. At 22 °C, a 2-fold difference in partitioning reflects a difference in solvation energy (∆Gsolv) of ∼0.39 kcal/mol that is more negative for methylcyclohexane than for isooctane. The terms in Table 2 are all much larger than this difference in relative solvation energies. Furthermore, the individual contributions theselves can not be isolated and independently tested. One might speculate that the energy required to create a cavity in methyl-cyclohexane should be smallest (i.e., least positive) given this solvent’s remarkably low melting temperature (-126 °C) and that the dispersive interactions should be greatest (i.e., largest magnitude, negative sign) in cyclohexane given the solvent’s anticipated registry (and strong, attractictive dispersive interactions) with the aromatic ring of the nitrophenol solutes. Such speculation, though, is based on expectation and not on empirical evidence. We do note that studies of a nonpolar solute (anthracene) solubility in alkane solvents reported anthracene to be more soluble in octane compared to isooctane and more soluble in methylcyclohexane compared to cyclohexane. However, in the anthracene studies, octane was the most accommodating solvent, whereas in the partitioning studies described above show clearly that nitrophenols are more soluble in both methylcyclohexane and cyclohexane compared to octane (and isooctane). Acknowledgment. This work was supported by the National Science Foundation through the Chemistry Division’s Experimental Physical Chemistry program (CHE0608122). Support for D.K.B. was provided by a grant to the University of Maryland from the Howard Hughes Medical Institute Undergraduate Science Education Program. J.B.F. thanks Gaussian Inc., for its financial support. W.H.S. thanks the Faculty Development Committee of York College for its support. R.A.W. gratefully acknowledges fellowship support from Durham University’s Institute for Advanced Study (UK) during the initial stages of manuscript preparation. References and Notes (1) Freiser, H. Chem. ReV. 1988, 88, 611. (2) Lennernas, H. J. Pharm. Pharmacol. 1997, 49, 627. (3) Safran, S. A. Statistical Thermodynamics of Surfaces, Interfaces and Membranes; Addison Wesley: Reading, MA, 1994. (4) Watarai, H. Trends Anal. Chem. 1993, 12, 313. (5) Wimley, W. C.; White, S. H. Biochemistry 1992, 31, 12813. (6) Woodrow, B. N.; Dorsey, J. G. EnViron. Sci. Technol. 1997, 31, 2812. (7) Ben-Amotz, D.; Raineri, F. O.; Stell, G. J. Phys. Chem. B 2005, 109, 6866. (8) Chan, H. S.; Dill, K. A. Annu. ReV. Biophys. Biomol. Struct. 1997, 26, 425. (9) Dill, K. A. Science 1990, 250, 297. (10) Widom, B.; Ben-Amotz, D. Proc. Natl. Acad. Sci., USA 2006, 103, 18887. (11) Acree, W. E. J. Chem. Soc., Faraday Trans. 1991, 87, 461. (12) Cabani, S.; Conti, G.; Mollica, V.; Bernazzani, L. J. Chem. Soc., Faraday Trans. 1991, 87, 2433. (13) Marche, C.; Ferronato, C.; Jose, J. J. Chem. Eng. Data 2004, 49, 937. (14) Constantinou, L.; Gani, R. AIChE J. 1994, 40, 1697. (15) Endo, S.; Schmidt, T. C. Fluid Phase Equilib. 2006, 246, 143. (16) Espinosa, S.; Diaz, S.; Fornari, T. Fluid Phase Equilib. 2005, 231, 197. (17) Fuentes, E.; Baez, M. E.; Reyes, D. Anal. Chim. Acta 2006, 578, 122.
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