Hydrocarbons dispersion forces

A thin film of hydrocarbon spread on a horizontal surface of quartz will experience a negative dispersion interaction. Treating these as 1 = quartz, 2 = n-decane, 3 = vacuum, determine the Hamaker constant A123 for the interaction. Balance the negative dispersion force (nonretarded) against the gravitational force to find the equilibrium film thickness. 

In order to include other interactions such as dipolar or hydrogen bonding, many semiempirical approaches have been tried [196, 197, 200], including adding terms to Eq. X-45 [198, 201] or modifying the definition of [202, 199]. Perhaps the most well-known of these approaches comes from Fowkes [203, 204] suggestion that the interactions across a water-hydrocarbon interface are dominated by dispersion forces such that Eq. X-45 could be modified as... 

As already mentioned molecules cohere because of the presence of one or more of four types of forces, namely dispersion, dipole, induction and hydrogen bonding forces. In the case of aliphatic hydrocarbons the dispersion forces predominate. Many polymers and solvents, however, are said to be polar because they contain dipoles and these can enhance the total intermolecular attraction. It is generally considered that for solubility in such cases both the solubility parameter and the degree of polarity should match. This latter quality is usually expressed in terms of partial polarity which expresses the fraction of total forces due to the dipole bonds. Some figures for partial polarities of solvents are given in Table 5.5 but there is a serious lack of quantitative data on polymer partial polarities. At the present time a comparison of polarities has to be made on a commonsense rather than a quantitative approach. 

Dispersion forces are ubiquitous and are present in all molecular interactions. They can occur in isolation, but are always present even when other types of interaction dominate. Typically, the interactions between hydrocarbons are exclusively dispersive and, because of them, hexane, at S.T.P., is a liquid boiling at 68.7°C and is not a gas. Dispersive interactions are sometimes referred to as hydrophobic or lyophobic particularly in the fields of biotechnology and biochemistry. These terms appear to have arisen because dispersive substances, e.g., the aliphatic hydrocarbons, do not dissolve readily in water. Biochemical terms for molecular interactions in relation to the physical chemical terms will be discussed later. 

Dispersive forces are more difficult to describe. Although electric in nature, they result from charge fluctuations rather than permanent electrical charges on the molecule. Examples of purely dispersive interactions are the molecular forces that exist between saturated aliphatic hydrocarbon molecules. Saturated aliphatic hydrocarbons are not ionic, have no permanent dipoles and are not polarizable. Yet molecular forces between hydrocarbons are strong and consequently, n-heptane is not a gas, but a liquid that boils at 100°C. This is a result of the collective effect of all the dispersive interactions that hold the molecules together as a liquid. 

Cation-7t and tu-tu Attractive interactions involving tc systems. The interaction energy depends on both the nature of the tc system and the nature of the caUon. When the ligand is a metal caUon, electrostaUc forces dominate the interacUon. When the ligand is a non-polar molecule (hydrocarbons, etc.) the dispersive interacUons dominate. A combinaUon of electrostaUc and dispersive forces governs the interacUon when the ligand is polar. 

AOT is an anionic surfactant complexed to the counterion, usually sodium. The water molecules in the intramicellar water pool are either free or bound to the interface. The bound water can interact with various parts of the surfactant. These interactions include hydrogen-bonding interactions with oxygen molecules on the sulfonate and succinate groups, ion-dipole interactions with the anionic surfactant headgroup and counterion, dipole-dipole interactions with the succinate group, and dispersive forces with the hydrocarbon tails. 

Dispersion forces are weak intermolecular forces. They are stronger, however, when the hydrocarbon part of a molecule is very large. Thus, a large molecule has stronger dispersion interactions than a smaller molecule. 

An examination of table 5.13 shows that the effects of substituting the —H group with any group will increase the boiling point, except for —F. The hydrocarbon groups add electrons, and thus increase the dispersion forces. Since the group contributions A Tb are ranked in the order... 

Heats of adsorption measurements do not lead to very specific interpretation since the isosteric heat of adsorption (AH) arises from both nonspecific interactions, which occur in all cases of adsorption, and from specific interactions with the hydroxy groups nevertheless, valuable conclusions about the binding forces can be deduced. Saturated hydrocarbons, e.g., n-pentane, have a value of — AH of 8.0 kcal/mole, while saturated ethers have values of around 16 kcal/mole.14 Probably dispersion forces only are involved in the former case and additional specific interaction with the silanol-OH occurs in the second case. On graphite, where there is no specific interaction, the heats of adsorption of hydrocarbons and ethers are very similar.17 The heat of adsorption of furan (11 kcal/mole) is 5 kcal/mole less than that of tetrahydrofuran this again indicates the effect that delocalization of electrons by the double bonds has on the binding forces.14... 

The values of — AH for benzene are in the range 10-12 kcal/mole,15,19-20 being intermediate between values attributed to pure dispersion forces for saturated hydrocarbons and those in which more specific forces are involved. Furthermore, Ron and coworkers calculated the entropies of adsorption for benzene and concluded that the mobile gas model of adsorption was applicable, and Whalen18 found no simple relationship between the hydroxy site content and benzene adsorption. These results confirm the conclusions reached from the infrared data that benzene adsorption is essentially due to dispersion forces which should be greater than with saturated compounds, and that no hydrogen bonding is involved. 

Azulene. The absorption spectrum of azulene, a nonbenzenoid aromatic hydrocarbon with odd-membered rings, can be considered as two distinct spectra, the visible absorption due to the 1Lb band (0-0 band near 700 nm) and the ultraviolet absorption of the 1L0 band.29 This latter band is very similar to the long wavelength bands of benzene and naphthalene CLb) and shows the same 130 cm-1 blue shift when adsorbed on silica gel from cyclohexane.7 As in the case of benzene and naphthalene, this blue shift is due to the fact that the red shift, relative to the vapor spectra, is smaller (305 cm"1) for the adsorbed molecule than in cyclohexane solution (435 cm"1). Thus it would appear that the red shifts of the 1La band are solely due to dispersive forces interacting with the aromatic molecule, in agreement with Weigang s prediction,29 and dipole-dipole interaction is negligible. 

The refractive index, d, is a measure of induced polarizability. Dispersion forces are especially high for aromatic hydrocarbons, which have highly polarizable k electrons. This is reflected in the high refractive indices of aromatic compounds, often 0.1 to 0.2 units higher than comparable nonaromatic compounds (table 3.5). Solvents with high polarizabilities are often good solvents for soft anions (i.e., those with high polarizabilities) such as SCN, F, and fF... 

London Dispersion Forces. These are due to the attraction of dipoles which arise from the arrangement of the elementary charges. Dispersion forces act between all molecular types and especially in the separation of nonpolar substances (e.g., saturated hydrocarbons). 

There is, however, a second component that results from dispersion energies. Dispersion forces are weak in water because of the low polarizability of oxygen (Table 11.2) and because of the low atom density (the dispersion energies are additive). This is an additional factor favoring the self-association of hydrocarbons, as they have a higher atom density and the polarizability of —CH2— is greater than that of O. 

The forces involved in the interaction al a good release interface must be as weak as possible. They cannot be the strong primary bonds associated with ionic, covalent, and metallic bonding neither arc they the stronger of the electrostatic and polarization forces that contribute to secondary van der Waals interactions. Rather, they are the weakest of these types of forces, the so-called London or dispersion forces that arise from interactions of temporary dipoles caused by fluctuations in electron density. They are common to all matter. The surfaces that are solid at room temperature and have the lowest dispersion-force interactions are those comprised of aliphatic hydrocarbons and fluorocarbons. 

A key advantage of semiempirical methods is that they give heats of formation directly. Small cyclic hydrocarbons are typically computed to be too stable, and sterically crowded structures are predicted to be too unstable. This is because semiempirical methods do not describe weak interactions well, e.g. those arising from London dispersion forces thus, they would not be suitable to describe, for instance, molecular structures that rely heavily on hydrogen bonding interactions. 

This equation assumes that the tip and sample are of the same material and no deformation occurs. If the tip and the surface are made of different materials, the cohesion energy wc=2y should be replaced by the work of adhesion. For a hydrocarbon polymer, where mainly dispersion forces are responsible for the tip-surface interaction, the work of adhesion can be estimated as wa =2(YtdYs)1/2>... 

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