Kaplan MCAT General Chemistry Review

Dispersion forces are the weakest of all the intermolecular interactions because they are the result of induced dipoles that change and shift moment by moment. They do not extend over long distances and are, therefore, significant only when molecules are close together. The strength of the London force also depends on the degree to which and the ease by which the molecules can be polarized (that is to say, how easily the electrons can be shifted around). Large molecules in which the electrons are far from the nucleus are relatively easy to polarize and, thus, possess greater dispersion forces. Think about two plates—a big dinner plate and a small teacup saucer—with nearly equal numbers of marbles on each. The marbles will roll around more, and may be distributed more unequally, on the bigger plate than on the smaller saucer. Nevertheless, don’t underestimate the importance of the dispersion forces. If it weren’t for them, the noble gases would not liquefy at any temperature because no other intermolecular forces exist between the noble gas atoms. The low temperatures at which the noble gases liquefy is to some extent indicative of the magnitude of the dispersion forces between the atoms.

DIPOLE–DIPOLE INTERACTIONS

Polar molecules tend to orient themselves in such a way that the opposite ends of the respective molecular dipoles are closest to each other: The positive region of one molecule is close to the negative region of another molecule. This arrangement is energetically favorable because an attractive electrostatic force is formed between the two molecules.

Dipole–dipole interactions are present in the solid and liquid phases but become negligible in the gas phase because of the significantly increased distance between gas particles. Polar species tend to have higher melting and boiling points than those of nonpolar species of comparable molecular weight. You should realize that London forces and dipole–dipole interactions are different not in kind but in degree. Both are electrostatic forces between opposite partial charges; the difference is only in the strength and in the permanence of the molecular dipole.

HYDROGEN BONDS

Hydrogen bonds are a favorite of the MCAT! So understand them well, and you will be able to answer every question pertaining to their nature, behavior, and effects. A hydrogen bond is actually only a specific, unusually strong form of dipole–dipole interaction, which may be intra- or intermolecular. Please understand that hydrogen bonds are not actually bonds—there is no sharing or transfer of electrons between two atoms. When hydrogen is bound to a highly electronegative atom, such as fluorine, oxygen, or nitrogen, the hydrogen atom carries little of the electron density of the covalent bond. The hydrogen atom acts essentially as a naked proton (it is a little scandalous!). The positively charged hydrogen atom interacts with the partial negative charge of oxygen, nitrogen, or fluorine on nearby molecules. Substances that display hydrogen bonding tend to have unusually high boiling points compared with compounds of similar molecular formula that do not hydrogen bond. The difference derives from the energy required to break the hydrogen bonds. A very good example of this difference, which you absolutely need to understand for the MCAT, is the difference between a carboxylic acid, which can form hydrogen bonds between molecules of itself, and the acyl halide, a derivative of the carboxylic acid, which has substituted a halogen for the hydroxyl group and which cannot form hydrogen bonds with molecules of itself. The boiling point of propanoic acid (MW = 74 g/mol) is 141°C, but the boiling point of propanoyl chloride (MW = 92.5 g/mol) is 80°C. Hydrogen bonding is particularly important in the behavior of water, alcohols, amines, and carboxylic acids. It is no overstatement to say that if it weren’t for water’s ability to form hydrogen bonds and exist in the liquid state at room temperature, we would not exist (at least not in the form we recognize as “human”).

Conclusion