Liquid intermolecular attractive

Density functional theory (DPT) calculations [18] suggest the tangential pulling force Ft as a solid/liquid adhesion originated in long-range solid/liquid intermolecular attraction... 

A substance exists as a liquid rather than a gas because attractive forces between molecules (intermolecular attractive forces) are greater in the liquid than in the gas phase. Attractive forces between neutral species (atoms or molecules, but not ions) are referred to as van der Waals forces and may be of three types ... 

Induced-dipole/induced-dipole attractions are ver-y weak forces individually, but a typical organic substance can par ticipate in so many of them that they are collectively the most impor tant of all the contributor s to intermolecular- attraction in the liquid state. They are the only forces of attraction possible between nonpolar- molecules such as alkanes. 

The magnitude of this effect depends on the strength of the attractive forces and hence on the nature of the gas. Intermolecular attractive forces are stronger in C02 than they are in 02, which explains why the deviation from ideality of Vmis greater with carbon dioxide and why carbon dioxide is more readily condensed to a liquid than is oxygen. 

Methyl ethyl ether is a gas at room temperature (boiling point = 8 °C), but 1-propanol, shown in Figure 11-13. is a liquid (boiling point = 97 °C). The compounds have the same molecular formula, C3 Hg O, and each has a chain of four inner atoms, C—O—C—C and O—C—C—C. Consequently, the electron clouds of these two molecules are about the same size, and their dispersion forces are comparable. Each molecule has an s p -hybridized oxygen atom with two polar single bonds, so their dipolar forces should be similar. The very different boiling points of 1-propanol and methyl ethyl ether make it clear that dispersion and dipolar forces do not reveal the entire story of intermolecular attractions. 

The graph in Figure 11-37 shows that adding heat to boiling water does not cause the temperature of the water to increase. Instead, the added energy is used to overcome intermolecular attractions as molecules leave the liquid phase and enter the gas phase. Other two-phase systems, such as an ice-water mixture, show similar behavior. 

There are similarities between the intermolecular attractions used to describe on a molecular level (1) viscosity, (2) surface tension, (3) the rate of evaporation and resulting vapor pressure of a liquid. For compounds in the liquid phase that have strong intermolecular forces of attraction operating between its molecules ... 

In the body of a liquid, intermolecular forces pull the molecules in all directions. At the surface of the liquid, the molecules pull down into the body of the liquid and from the sides. There are no molecules above the surface to pull in that direction. The effect of this unequal attraction is that the liquid tries to minimize its surface area. The minimum surface area for a given quantity of matter is a sphere. In a large pool of liquid, where sphere formation is not possible, the surface behaves as if it had a thin stretched elastic membrane or skin over it. The surface tension is the resistance of a liquid to an increase in its surface area. It requires force to break the attractive forces at the surface. The greater the intermolecular force, the greater the surface tension. Polar liquids, especially those that utilize hydrogen bonding, have a much higher surface tension than nonpolar liquids. 

Air at room temperature and pressure consists of 99.9% void and 0.1% molecules of nitrogen and oxygen. In such a dilute gas, each individual molecule is free to travel at great speed without interference, except during brief moments when it undertakes a collision with another molecule or with the container walls. The intermolecular attractive and repulsive forces are assumed in the hard sphere model to be zero when two molecules are not in contact, but they rise to infinite repulsion upon contact. This model is applicable when the gas density is low, encountered at low pressure and high temperature. This model predicts that, even at very low temperature and high pressure, the ideal gas does not condense into a liquid and eventually a solid. 

The cohesive energy density of a low molecular weight liquid is given by the heat of vaporization of that liquid, expressed per unit volume. Verify the value given for the CED of acetone in Table 3.1 from the facts that AHv = 30.2 kJ mole 1 and p = 0.792 g cm 3. Would it be better to use the change in internal energy on vaporization At/, rather than AH, as a measure of intermolecular attraction Does it make much difference quantitatively whether At/, or AHv is used ... 

Now we can see how a chemical s structure causes it to have a particular vapor pressure. This is possible because, as a first approximation, the free energy of vaporization, A G, mostly differs from compound to compound due to differences in those substances enthalpies of vaporization, Avap//,-. These enthalpies reflect the sum of intermolecular attractions that act to hold those liquid molecules together. Thus, we can expect that substances that exhibit high vapor pressures have structures that do not enable the molecules to have strong intermolecular attractions. Conversely, molecules with low vapor pressures must have structures that cause the molecules to be substantially attracted to one another. 

Moreover, this relation between chemical structure and vapor pressure also holds because enthalpies and entropies of vaporization are directly related, in general. Recall that the entropy of vaporization reflects the difference of a molecule s freedom in the gas phase versus the liquid phase (A pS = Si% - SiL). At ambient pressures, we may assume that differences in Avap5) between different compounds are primarily due to differences in molecular freedom in the liquid phase. (The freedom of the molecules in the gas phase is not that different between compounds). Hence, not surprisingly, molecules that exhibit stronger intermolecular attractions... 

For bipolar organic liquids, especially for hydrogen-bonding liquids such as alcohols and amines, the tendency to orient in the liquid phase, due to these highly directional intermolecular attractions, is greatly increased by this intermolecular interaction. We can see the effect of this in the significantly larger entropies of vaporization of bipolar chemicals, like aniline, phenol, benzyl alcohol, or ethanol (Table 4.2). 

Van der Waals sought to address two basic defects of the KMT noninteracting point mass picture (i) neglect of the finite molecular volume that distinguishes molecules from mathematical points and (ii) neglect of the intermolecular attraction that leads to condensation (liquid formation) at sufficiently low temperature. Whereas the ideal gas equation (2.2) exhibits no vestige of condensation phenomena, the Van der Waals equation (2.13) is intended to provide a unified description of gas-liquid ( fluid ) behavior, exhibiting the essential commonality that must be shared by these disparate forms of matter at the molecular level. 

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