Chloromethane \(\left(\mathrm{CH}_{3} \mathrm{Cl}\right)\), methane, and acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) form molecular solids. (a) What types of forces hold these molecules in a molecular solid? (b) Place the solids in order of increasing melting point.

Molecular solids are types of chemical compounds where the individual molecules are held together by intermolecular forces rather than ionic or covalent bonds. These forces are relatively weak compared to the bonds within molecules, which makes molecular solids typically soft and with low melting and boiling points.

In the given exercise, chloromethane, methane, and acetic acid all form molecular solids. However, the types of intermolecular forces at play differ between them, which affects their physical properties such as melting points. For students aiming to grasp these concepts, it is important to understand these differing forces so they can predict and explain the behaviors of various molecular solids.

Dipole-dipole interactions occur when two polar molecules interact with each other. These molecules have areas of partial positive and negative charge due to an uneven distribution of electrons, which allows them to attract each other much like magnets. An example provided in the exercise is chloromethane (CH₃Cl), which is a polar molecule.

To understand this further, imagine each chloromethane molecule having a slight positive charge near the hydrogen atoms and a slight negative charge near the chlorine atom, leading to an attraction between oppositely charged ends of nearby molecules. This type of interaction is stronger than London dispersion forces found in nonpolar molecules.

London dispersion forces are the weakest type of intermolecular force. They are found in all molecular solids, including nonpolar molecules like methane (CH₄). They arise from temporary dipoles that occur when the electrons in a molecule are unevenly distributed at a particular instant in time, which in turn induces a temporary dipole in neighboring molecules.

In simpler terms, even though methane molecules do not have a permanent dipole, the random movement of electrons can create a fleeting positive and negative side that attracts other molecules momentarily. This fleeting nature means that the London dispersion forces are much weaker than permanent dipole-dipole interactions.

Hydrogen bonding is a particularly strong type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative elements, such as nitrogen, oxygen, or fluorine. In acetic acid (CH₃COOH), the hydrogen atoms of the -OH group can form hydrogen bonds with the oxygen atoms of other acetic acid molecules.

This strong attraction is due to the significant difference in electronegativity between the hydrogen and the oxygen, which pulls the electrons away from the hydrogen, creating a highly polar bond. The positively charged hydrogen is then strongly attracted to the negative regions of nearby molecules. This is why, among the molecular solids mentioned in the textbook exercise, acetic acid has the highest melting point—it takes more energy to break these strong hydrogen bonds.

The melting point of a substance is the temperature at which it changes from a solid to a liquid. This physical property is greatly influenced by the strength of the intermolecular forces within a molecular solid. Stronger intermolecular forces require more thermal energy to overcome, resulting in a higher melting point.

In the context of molecular solids, it is crucial to recognize that melting point is indicative of the robustness of the intermolecular forces binding the solid together. From the exercise, it is clear that acetic acid, with substantial hydrogen bonding, has the highest melting point, followed by chloromethane with its dipole-dipole interactions, and finally methane, which is bound by weak London dispersion forces.