Difference between revisions of "Solvation and hydrophobic forces"
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== Solvation force - oscillatory ==
== Solvation force - oscillatory ==
Revision as of 18:17, 29 September 2008
Molecular ordering at surfaces
The liquid density at:
- (a) The vapor/liquid interface.
- (b) The solid/liquid interface.
- (c) The solid/liquid/solid interface.
This effect seems to be on a molecular level but does it have any truth on a macroscopic level? Or microfluidics? What length scales are important?
Look who worked on this topic! Strange coincidence or the one and the same? Below is a paper abstract.
Ordering in a Fluid Inert Gas Confined by Flat Surfaces
Stephen E. Donnelly,1* Robert C. Birtcher,2 Charles W. Allen,2 Ian Morrison,1 Kazuo Furuya,3 Minghui Song,3 Kazutaka Mitsuishi,3 Ulrich Dahmen4
High-resolution transmission electron microscopy images of room-temperature fluid xenon in small faceted cavities in aluminum reveal the presence of three well-defined layers within the fluid at each facet. Such interfacial layering of simple liquids has been theoretically predicted, but observational evidence has been ambiguous. Molecular dynamics simulations indicate that the density variation induced by the layering will cause xenon, confined to an approximately cubic cavity of volume approx 8 cubic nanometers, to condense into the body-centered cubic phase, differing from the face-centered cubic phase of both bulk solid xenon and solid xenon confined in somewhat larger (>=20 cubic nanometer) tetradecahedral cavities in face-centered cubic metals. Layering at the liquid-solid interface plays an important role in determining physical properties as diverse as the rheological behavior of two-dimensionally confined liquids and the dynamics of crystal growth.
1 Joule Physics Laboratory, Institute for Materials Research, University of Salford, Manchester M5 4WT, UK. 2 Materials Science Division, Argonne National Laboratory, Argonne IL 60439, USA. 3 National Institute for Materials Science, 3-13 Sakura, Tsukuba 305, Japan. 4 National Center for Electron Microscopy, LBNL, Berkeley, CA 94720, USA.
This improved packing has been shown to actually cause colloidal gas to aggregate. Through studies of phase transitions that occur through surface modifications of colloidal systems, Roke et al. demonstrated how the increased density from intermolecular packing increased the van der Waals attractions between the particles which could then cause the colloidal gas to aggregate.
Read more in: S. Roke, O. Berg, J. Buitenhuis, A. van Blaaderen, and M. Bonn, "Surface molecular view of colloidal gelation," Proc. Nat. Acad. Sci. USA 103, 13310-13314(2006). http://www.imsc.res.in/~sitabhra/teaching/cmp03/class3.html
Solvation force - oscillatory
Note that the number of spheres in contact with the surface also varies with the maxima and minima.
The corresponding oscillatory solvation forces are shown in the lower graph.
Meaning of solvation: Any stabilizing interaction of a solute (or solute moiety) and the solvent or a similar interaction of solvent with groups of an insoluble material (i.e. the ionic groups of an ion-exchange resin). Such interactions generally involve electrostatic forces and van der Waals forces, as well as chemically more specific effects such as hydrogen bond formation. <http://goldbook.iupac.org/S05747.html>
Intermolecular interaction is definitely critical to understanding the solvation force. Polar solvents are those with a molecular structure that contains dipoles. Such compounds are often found to have a high dielectric constant. The polar molecules of these solvents can solvate ions because they can orient the appropriate partially charged portion of the molecule towards the ion in response to electrostatic attraction. This stabilizes the system. Water is the most common and well-studied polar solvent, but others exist, such as acetonitrile, dimethyl sulfoxide, methanol, propylene carbonate, ammonia, ethanol, and acetone. These solvents can be used to dissolve inorganic compounds such as salts.
Solvation involves different types of intermolecular interactions: hydrogen bonding, ion-dipole and dipole-dipole attractions or van der Waals forces. The hydrogen bonding, ion-dipole, and dipole-dipole interactions occur only in polar solvents. Ion-ion interactions occur only in ionic solvents. The solvation process will only be thermodynamically favored if the overall Gibbs energy of the solution is decreased compared to the Gibbs energy of the separated solvent and solid (or gas or liquid). This means that the change in enthalpy minus the change in entropy (multiplied by the absolute temperature) is a negative value, or that the Gibbs free energy of the system decreases.
Conductivity of a solution depends on the solvation of their ions.
Measured oscillatory forces
Forces between mica sheets separated by a liquid. The doted line is the theoretical calculation.
Not so good!!
Molecular “ordering” at a surface is a significant factor.
The hydrophobic efffect
Hydrophobicity refers to the physical property of a molecule that is repelled from a mass of water. Hydrophobic molecules tend to be non-polar and thus prefer other neutral molecules and nonpolar solvents. Hydrophobic molecules in water often cluster together forming micelles. This minimizes the the surface contact between hydrophobic molecules and water. Water on hydrophobic surfaces will exhibit a high contact angle. Examples of hydrophobic molecules include the alkanes, oils, fats, and greasy substances in general. Hydrophobic materials are used for oil removal from water, the management of oil spills, and chemical separation processes to remove non-polar from polar compounds.
According to thermodynamics, matter seeks to be in a low-energy state, and bonding reduces chemical energy. Water is electrically polarized, and is able to form hydrogen bonds internally, which gives it many of its unique physical properties. But, since hydrophobes are not electrically polarized, and because they are unable to form hydrogen bonds, water repels hydrophobes, in favour of bonding with itself. It is this effect that causes the hydrophobic interaction — which in itself is incorrectly named as the energetic force comes from the hydrophilic molecules. Thus the two immiscible phases (hydrophilic vs. hydrophobic) will change so that their corresponding interfacial area will be minimal. This effect can be visualized in the phenomenon called phase separation.
Water near a “hydrophobic” surface re-arranges to maintain hydrogen bonding. This decreases the entropy and hence raises the free energy. Hydrophobic molecules can dissolve in water to a small extent, which is known as hydrophobic hydration. Small non-polar molecules have the highest degree of solvation in water. Larger hydrophobic molecules (particularly if they have charged areas) due to entropic effects. Water molecules arrange themselves around the hydrophobic molecules without breaking hydrogen bonds (meaning there's great flexibility in spatial arrangement) or experiencing great losses of energy. The interaction between the dissolved hydrophobic and water molecules can be described by Van der Waals forces.
There is a strong dependence on temperature and pressure on a hydrophobic molecule's ability to become solvated. At very low temperatures, there is a shift in the characteristic interaction between water and non-polar groups, and once again at a certain point the penalty of solvating a non-polar group decreases. In other words, the entropic penalty for solvating non-polar molecules disappears at low temperatures where the solvent-solvent interaction energy becomes comparable to the solvent-solute interaction energy. At high pressures, the solute/solvent system starts behaving increasingly more like a hard sphere system. Solvation happens at the hard sphere limit. Hydrophobic hydration actually causes entropy to go down, because there is increased ordering of water molecules and correlated Van der Waals interactions. At the same time, enthalpy decreases. As we learn in a first semester chemistry course, a process where dS is negative and dH is negative is spontaneous at low temperatures. As said before, this is true. Furthermore, we expect that as temperature increase, solvation decreases. In fact, as temperature elevates, solvation reaches a minimum. However, after this minimum (which depends on all sorts of parameters, such as solute radii, pressure, etc), solvation actually increases. At high temperatures, the entropy of the system increases, and there comes a point where the system has enough energy that there is an extremely small penalty in solvating non-polar groups. More simply, greatly increased temperature causes there is so much entropy in the system that water molecules stop caring what they are in contact with, allowing for greater solvation.
Many water molecular models have been developed to help discover the structure of water. Different models have been created for different applications, but the truth is that no water model captures all of the properties of water at varied thermodynamic conditions. There are, however, some very robust models for certain applications. Some famous models are TIP3P and SPC/E. One of the biggest problems with physical modeling of water comes from determining which values and parameters one should seek to reproduce. Some models focus on bulk properties such as density of vapor pressure. Others attempt to reproduce the electrostatics well, such as accurately predicting dipole moments. Yet other models aim to reproduce the vibrational spectrum and transport properties of a water system. These models look at things like infrared and x-ray absorption data, diffusion and viscosity. On top of everything, there is great controversy about how water molecules should be represented. Because it is so difficult to model water, and because there is so little agreement amongst experts, it is often thought to be best to develop models that accurately predict different properties of water for a very specific system (such as the interaction of a biological molecule with water).
The hydrophobic effect is extremely important in the following molecular interactions:
- (a) Solubility
- (b) Micellization
- (c) Association
- (d) Protein folding
Less so in
- (e) Adhesion
- (f) Wetting
- (g) Flocculation
- (h) Flotation.
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