Intermolecular Forces: Boiling Point, Vapor Pressure and Solubility

Attractions between molecules, not bonds within them

Covalent bonds hold the atoms of a molecule together. Intermolecular forces (IMFs) are the much weaker attractions between separate molecules, and they are what must be overcome when a liquid boils or a solid melts. When water boils, the gas that forms is still made of H₂O molecules; no O–H bonds break. Only the attractions between neighboring molecules are disrupted.

Because these forces govern how strongly molecules cling to one another, they control a set of physical properties together. Stronger IMFs generally mean a higher boiling point, a higher enthalpy of vaporization, a lower vapor pressure at a given temperature, and often higher viscosity and surface tension. If you can rank IMFs, you can usually rank all of these at once.

London dispersion forces in every substance

Dispersion forces arise because electron distributions fluctuate. At any instant the electrons in a molecule may be slightly uneven, creating a temporary dipole that induces a matching dipole in a neighbor. The two attract briefly, and the pattern repeats across the whole liquid. Every atom and molecule experiences dispersion forces, polar or not.

The strength of dispersion forces depends on polarizability, meaning how easily the electron cloud can be distorted. Larger molecules with more electrons are more polarizable, so dispersion forces generally increase with molar mass within a family of similar compounds. Shape also matters: long, extended molecules have more surface in contact with their neighbors than compact, branched ones of the same formula, so they attract each other more strongly.

Do not underestimate dispersion. In large molecules it is often the largest contribution to the total attraction, even when the molecule is also polar.

Dipole–dipole attractions and hydrogen bonding

Polar molecules have permanent dipoles, so in a liquid they tend to orient with the partial positive end of one molecule near the partial negative end of another. This dipole–dipole attraction adds to the dispersion forces that are always present. Deciding whether a molecule is polar requires both its bond polarities and its shape.

Hydrogen bonding is an especially strong dipole–dipole attraction. It requires a hydrogen atom covalently bonded to nitrogen, oxygen or fluorine (the donor) and a lone pair on a nearby N, O or F atom (the acceptor). The very polar N–H, O–H or F–H bond leaves the small hydrogen atom with a large partial positive charge that can approach the acceptor's lone pair closely.

A molecule can be an acceptor without being a donor. Acetone, (CH₃)₂CO, has a lone pair on oxygen but all of its hydrogens are bonded to carbon, so acetone molecules cannot hydrogen bond with each other. They can, however, accept hydrogen bonds from water, which is one reason acetone mixes with water in all proportions.

Ion–dipole forces and dissolving ionic solids

When an ionic solid dissolves in water, each ion becomes surrounded by water molecules oriented toward it: the partially negative oxygen ends face cations, and the partially positive hydrogen ends face anions. This ion–dipole attraction releases enough energy to offset much of the energy needed to separate the ions from the crystal. The attraction is stronger for ions of higher charge and smaller size.

Ion–dipole forces explain why potassium bromide dissolves readily in water but hardly at all in hexane, which has no dipole to attract ions. The general rule "like dissolves like" summarizes this: substances dissolve best in solvents whose intermolecular forces are similar to their own. Iodine, a nonpolar molecule, dissolves well in hexane, where dispersion forces between solute and solvent are comparable to those within each pure substance.

Worked example: three liquids with almost the same molar mass

Rank butane (CH₃CH₂CH₂CH₃), acetone ((CH₃)₂CO) and 1-propanol (CH₃CH₂CH₂OH) by boiling point. First compare molar masses, using IUPAC standard atomic weights rounded to two decimal places (H 1.01, C 12.01, O 16.00 g/mol). Butane, C₄H₁₀: 4(12.01) + 10(1.01) = 58.14 g/mol. Acetone, C₃H₆O: 3(12.01) + 6(1.01) + 16.00 = 58.09 g/mol. 1-Propanol, C₃H₈O: 3(12.01) + 8(1.01) + 16.00 = 60.11 g/mol. The masses differ by less than 4%, so dispersion forces should be roughly similar and the differences will come from polarity.

Next classify the forces. Butane contains only C–C and C–H bonds and is essentially nonpolar, so it has dispersion forces only. Acetone has a polar C=O bond in a trigonal planar arrangement that does not cancel, so it adds dipole–dipole forces, but it has no O–H bond and cannot donate hydrogen bonds. 1-Propanol has an O–H group, so it adds hydrogen bonding. The prediction is butane < acetone < 1-propanol.

The NIST Chemistry WebBook lists normal boiling points of about 273 K (0 °C) for butane, 329.3 K (56.2 °C) for acetone and 370.3 K (97.2 °C) for 1-propanol, converting with T(°C) = T(K) − 273.15. The order matches the prediction. The same reasoning ranks vapor pressure at room temperature in the reverse order: butane is a gas at 25 °C, acetone evaporates quickly, and 1-propanol evaporates more slowly.

Worked example: when molecular size outweighs a dipole

Now compare diethyl ether ((C₂H₅)₂O), 1-butanol (CH₃CH₂CH₂CH₂OH) and hexane (C₆H₁₄). Diethyl ether and 1-butanol are isomers, both C₄H₁₀O: 4(12.01) + 10(1.01) + 16.00 = 74.14 g/mol. Hexane is heavier: 6(12.01) + 14(1.01) = 86.20 g/mol. Diethyl ether is polar but has no O–H bond, 1-butanol can hydrogen bond, and hexane is nonpolar.

Between the two isomers, hydrogen bonding decides: 1-butanol should boil far above diethyl ether. Hexane is the instructive case. A polarity-only argument would put it last, but hexane is larger, has more electrons and is an extended chain, so its dispersion forces are larger than the ether's dispersion plus its weak dipole–dipole attraction.

The NIST values confirm this: diethyl ether 307.7 K (34.6 °C), hexane 341.9 K (68.8 °C) and 1-butanol 390.6 K (about 117 °C). The lesson is that polarity comparisons are safest between molecules of similar size. When the sizes differ substantially, consider dispersion first and then ask whether a dipole or hydrogen bonding is strong enough to overturn it. Solubility follows the forces too: hexane is essentially insoluble in water, while 1-butanol, which can hydrogen bond with water, is partly soluble.

Errors to avoid when comparing intermolecular forces

Before ranking, list every force each substance has rather than just the strongest one, and note the molar masses. This prevents the most common mistake, which is comparing only labels such as "polar" and "nonpolar" while ignoring size.

The errors below come up repeatedly in explanations of boiling point and solubility.

• Saying that boiling breaks covalent bonds. Boiling separates molecules; the molecules stay intact.
• Claiming hydrogen bonding for any molecule that contains hydrogen. The hydrogen must be bonded to N, O or F.
• Treating nonpolar substances as having no intermolecular forces. All substances have dispersion forces.
• Ranking molecules of very different size by polarity alone, as the hexane example shows.
• Reversing the vapor pressure trend. Stronger IMFs mean lower vapor pressure and a higher boiling point.
• Treating "like dissolves like" as a law rather than a rule of thumb. Solubility also depends on entropy and temperature.