Account for the following observations in terms of the type and strength of intermolecular forces. (a) The melting point of solid xenon is \(-112^{\circ} \mathrm{C}\) and that of solid argon is \(-189^{\circ} \mathrm{C}\). (b) The vapor pressure of diethyl ether \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OC}_{2} \mathrm{H}_{5}\right)\) is greater than that of water. (c) The boiling point of pentane, \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{3} \mathrm{CH}_{3}\), is \(36.1^{\circ} \mathrm{C}\), whereas that of 2,2 -dimethylpropane (also known as neopentane \(), \mathrm{C}\left(\mathrm{CH}_{3}\right)_{4}\), is \(9.5^{\circ} \mathrm{C}\).

Understanding why different substances have distinct melting points helps us comprehend their unique physical properties. Melting points are a measure of the temperature at which a solid turns into a liquid. This physical change happens when molecules have enough thermal energy to overcome the intermolecular forces that hold them in a fixed position within the solid state.

For noble gases like xenon and argon, the melting points are significantly affected by the presence of dispersion forces, also known as London forces. Since these gases exist as individual atoms, the strength of their dispersion forces—and hence their melting points—correlate with the size of their electron clouds. Xenon, with a larger electron cloud than argon, exhibits stronger dispersion forces, which translates to a higher melting point. This example makes it clear that even seemingly weak intermolecular forces are critical in determining the melting point of substances.

Vapor pressure is a fascinating concept, referring to the pressure exerted by a vapor in equilibrium with its liquid or solid form at a given temperature. Essentially, it's a measure of a substance's tendency to evaporate.

When comparing substances like diethyl ether and water, we notice that diethyl ether has a higher vapor pressure. This is because diethyl ether can't form hydrogen bonds, which are very strong types of intermolecular forces. Hydrogen bonds in water significantly lower its vapor pressure, as more energy is required to break these bonds and allow molecules to escape into the gaseous state. In contrast, diethyl ether’s primary intermolecular forces are weaker dipole-dipole interactions, making it easier for molecules to evaporate and thus contributing to its higher vapor pressure.

The boiling point is the temperature at which a liquid's vapor pressure equals the external pressure surrounding the liquid, and thus it begins to transform into gas. Substances with high intermolecular forces typically have higher boiling points because more energy is needed to separate the molecules.

Looking at substances like pentane and neopentane, we observe that pentane has a higher boiling point. This is attributed to its linear structure which allows for more extensive dispersion forces due to a larger surface area for interactions. Neopentane, with its more compact shape, has less surface area and hence weaker dispersion forces. Therefore, less energy is required to reach neopentane’s boiling point, making it lower than that of pentane. This example demonstrates the direct relationship between molecular structure, intermolecular forces, and boiling points.

Dispersion forces, also known as London forces, are the weakest intermolecular forces but are universal; they occur in all molecules, polar and nonpolar alike. These forces originate from temporary dipoles that develop when the electron clouds of atoms or molecules fluctuate momentarily.

Since the strength of dispersion forces increases with the size of the electron cloud, larger molecules or atoms have stronger dispersion forces. This accounts for the difference in melting and boiling points of substances, as seen in noble gases and hydrocarbon chains. Larger, more elongated hydrocarbon chains have more surface area to interact, leading to stronger dispersion forces and higher melting and boiling points compared to their branched or smaller counterparts.

Hydrogen bonding, a strong type of intermolecular force, occurs when a hydrogen atom is bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine and is attracted to another electronegative atom. This special interaction has a critical role in determining physical properties like boiling points and solubility.

Water molecules, for instance, form extensive hydrogen bonds. This results in water's high boiling point compared to other substances of similar size and mass. The ability to form hydrogen bonds also explains why substances with these types of intermolecular forces often have lower vapor pressures, as more energy is needed to break the bonds during evaporation.

Dipole-dipole interactions are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. These interactions are stronger than dispersion forces but weaker than hydrogen bonding.

They affect how molecules align with one another and influence properties such as melting points, boiling points, and vapor pressures. A polar molecule like diethyl ether will have a higher vapor pressure because it relies mainly on dipole-dipole interactions, which are weaker compared to the hydrogen bonds in water, which lower its vapor pressure. Understanding the relative strengths of these forces helps explain the physical behaviors of different substances.