Salting-In and Salting-Out of Hydrophobic Solutes in Aqueous Salt Solutions

Abstract

We present results on the thermodynamic and structural aspects of the hydration of hydrophobic solutes in three tetramethylammonium [N(CH 3 ) 4 + ] salt solutions at various concentrations obtained from molecular dynamics simulations. Monovalent counterions of different sizessF -, Cl -, and a relatively large model ion BIsare chosen in order to cover a range of kosmotropic to chaotropic behaviors. Chemical potentials of hard-sphere solutes obtained using test particle insertions display both salting-in and salting-out effects depending on the type of salt. Water and salt-ion densities in the vicinity of hard-sphere solutes are calculated. Small and strongly hydrated Fions (kosmotropes) are excluded from the vicinity of hydrophobic solutes, leading to an increase in local water densities near hydrophobic solutes (i.e., preferential hydration). This increases the excess chemical potential of hydrophobic solutes in solution which leads to salting-out. Opposite behavior is observed for large, less favorably hydrated BIions (chaotropes) which associate strongly with hydrophobic solutes. Compressive forces due to neighboring water molecules, cations, and anions on the surface of the hard sphere solute are calculated. We find that water molecules make the most significant contribution toward the total compressive force. This explains the observed linear correlation between the extent of preferential hydration or dehydration of the solute surface and salting-out or salting-in effects. The trends in the thermodynamics of hydration of hydrophobic solutes upon addition of salts are explained in terms of the structural hydration of individual salt ions.

1H NMR and Molecular Dynamics Evidence for an Unexpected Interaction on the Origin of Salting-In/Salting-Out Phenomena

“…It is evident from Figures 9 and 10 that association of ClO 4 -with the IL cation persists over the entire range of concentrations studied. RDFs of the salt anion around H(6), H(7), H (8), and H(9) in the IL cation are in stark contrast to those obtained for Cl -and SO 4 2-discussed below. The RDFs show that ClO 4 -interacts strongly with the alkyl tails of the IL cation, as shown by the ordering in the first and second solvation shells of H(6), H(7), H (8), and H(9).…”

“…RDFs of the salt anion around H(6), H(7), H (8), and H(9) in the IL cation are in stark contrast to those obtained for Cl -and SO 4 2-discussed below. The RDFs show that ClO 4 -interacts strongly with the alkyl tails of the IL cation, as shown by the ordering in the first and second solvation shells of H(6), H(7), H (8), and H(9). The intensity of the first solvation peak is comparable to that observed for the imidazolium ring hydrogen atoms.…”

“…The most interesting and unexpected results were obtained for the RDFs of salt anions and the IL cation protons H(6), H(7), H (8), and H(9). Figure 8 shows the interaction between Cl and the alkyl chain protons.…”

“…11 The dependency of the aqueous solubility of IL in pure water compared to that in the presence of salt with the chemical shift deviations for protons H(7), H (8), and H(9) is shown in Figure 5, for the sodium-based salts. Note that no correlation was found for the other protons present at the imidazolium cation.…”

“…[1][2][3][4] Recent experimental data and simulation results have cast doubts on this paradigm of the change in the bulk water structure as the main phenomenon behind the effect of salts on the molecules' aqueous solubility. [4][5][6][7][8][9][10] Both experimental studies [4][5][6][7] and molecular simulation calculations [8][9][10] are converging upon the idea that ions of high charge densityswith water-ion interactions stronger than water-water interactionssare excluded from the vicinity of the solute due to their preferential hydration, decreasing the solute's solubility in water. On the other hand, low charge density ionsswith water-ion interactions weaker than water-water interactionssdirectly interact with the solute thus stabilizing it in water, and therefore increasing the solute's solubility.…”