8.4 Intermolecular Forces in the Liquid Phase

Intermolecular Forces: Overview

Intermolecular forces (IMFs) are attractive forces that act between separate molecules, as opposed to intramolecular forces (covalent and ionic bonds), which hold atoms together within a single molecule. IMFs are dramatically weaker than intramolecular bonds — typically in the range of about 1 to 40 kilojoules per mole, compared to roughly 150 to 950 kilojoules per mole for covalent bonds — but they are what determine nearly every bulk physical property of a substance: melting point, boiling point, viscosity, surface tension, vapor pressure, and solubility.

As a general rule, stronger intermolecular forces mean higher melting and boiling points, because more thermal energy is required to pull molecules apart from each other (melting) or to separate them entirely into the gas phase (boiling). The three intermolecular forces tested on the MCAT, from strongest to weakest per interaction, are hydrogen bonding, dipole-dipole interactions, and London dispersion forces (a category within the broader term Van der Waals forces).

London Dispersion Forces

London dispersion forces, the weakest individual intermolecular force, arise in every atom and molecule, polar or nonpolar, from the constant, random motion of electrons. At any given instant, electron distribution is not perfectly even, creating a temporary, instantaneous dipole; this instantaneous dipole induces a complementary, temporary dipole in a neighboring molecule, and the two attract each other briefly before the fluctuation reverses. Because they don't require a permanent dipole, dispersion forces are the only intermolecular force available to purely nonpolar species — noble gases, diatomic elements like nitrogen and oxygen, and nonpolar hydrocarbons.

Dispersion force strength increases with polarizability — how easily a molecule's electron cloud can be distorted — which in turn increases with molecular size, number of electrons, and available surface area for contact between neighboring molecules. Two classic MCAT trends follow directly from this:

• Halogen boiling points increase down the group: F2 (boils at −188°C, a gas at room temperature), Cl2 (−34°C, gas), Br2 (59°C, liquid at room temperature), I2 (184°C, solid at room temperature). None of these nonpolar diatomic molecules can hydrogen bond or form permanent dipoles, so this steady increase is due entirely to larger, more polarizable electron clouds as atomic number increases down the group.
• Branching lowers boiling point among constitutional isomers: straight-chain alkanes have higher boiling points than their branched isomers of the same molecular formula, because an elongated, unbranched chain allows more surface-area contact with neighboring molecules, maximizing the cumulative dispersion force, while a compact, branched structure reduces that contact area.

Despite being the weakest force per interaction, dispersion forces can dominate in sufficiently large molecules — this is why large nonpolar molecules like long-chain fatty-acid tails or large hydrocarbons can still have substantial boiling points and significant intermolecular attraction, simply from the sheer number of atoms contributing dispersion contacts.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between molecules that possess a permanent dipole moment (polar molecules, as introduced in section 8.1). The partially positive end of one polar molecule is electrostatically attracted to the partially negative end of a neighboring molecule, and these interactions are, on average, stronger than dispersion forces between molecules of comparable size and molecular weight, though still weaker than hydrogen bonding.

A useful comparison: acetone (a polar molecule with a strong carbonyl dipole, boiling point 56°C) boils at a substantially higher temperature than a nonpolar hydrocarbon of similar molecular weight, which relies on dispersion forces alone. Dipole-dipole interactions require the molecules to be reasonably close and favorably oriented (positive end near negative end), so they are generally shorter-range and more geometry-dependent than dispersion forces.

Hydrogen Bonding

Hydrogen bonding is a particularly strong, specific type of dipole-dipole interaction that occurs when a hydrogen atom is covalently bonded directly to a small, highly electronegative atom carrying lone pairs — specifically nitrogen, oxygen, or fluorine (a common mnemonic is 'N, O, F'). Because these atoms pull electron density strongly away from hydrogen, and because hydrogen has no inner-shell electrons to shield its nucleus, the resulting hydrogen atom carries an unusually large partial positive charge and can interact strongly with a lone pair on a nitrogen, oxygen, or fluorine atom of a neighboring molecule.

Hydrogen bonds are roughly an order of magnitude weaker than a typical covalent bond, but several times stronger than a typical dipole-dipole interaction between molecules lacking N-H, O-H, or F-H bonds.

A classic demonstration compares the boiling points of period 2 hydrides to their heavier period 3 analogs: water (H2O, boiling point 100°C) and ammonia (NH3, boiling point −33°C) both boil far higher than their heavier counterparts hydrogen sulfide (H2S, boiling point −60°C) and phosphine (PH3, boiling point −88°C), even though H2S and PH3 have greater molecular weight, which would ordinarily predict stronger dispersion forces and a higher boiling point. The period 2 hydrides can hydrogen bond because nitrogen and oxygen are small and highly electronegative; sulfur and phosphorus are larger and less electronegative, so H2S and PH3 rely only on weaker dipole-dipole and dispersion forces.

Biological significance of hydrogen bonding is extensive:

• DNA base pairing: adenine pairs with thymine through two hydrogen bonds, while guanine pairs with cytosine through three hydrogen bonds — the extra hydrogen bond makes GC-rich DNA regions more thermally stable (higher melting temperature) than AT-rich regions.
• Protein secondary structure: alpha helices and beta sheets are held together by hydrogen bonds between the backbone amide N-H (donor) and carbonyl C=O (acceptor) groups along the polypeptide chain.
• Water's unusual properties: water's extensive three-dimensional hydrogen-bonding network gives it an unusually high boiling point, high heat of vaporization, high surface tension, and the well-known property that ice is less dense than liquid water (the rigid, open hydrogen-bonded lattice in ice holds molecules farther apart, on average, than in the liquid), which allows ice to float and insulate bodies of water beneath it.

Intermolecular Forces, Solubility, and Biological Structure

Intermolecular forces govern solubility through the general principle 'like dissolves like': a solute dissolves readily in a solvent when the solute-solvent interactions that form are comparable in strength to the solute-solute and solvent-solvent interactions that must be broken. Polar and hydrogen-bonding solutes (such as small alcohols, sugars, and ionic compounds) dissolve well in polar, hydrogen-bonding solvents like water, because new solute-water hydrogen bonds and dipole interactions offset the energy needed to disrupt water's existing hydrogen-bond network. Nonpolar solutes dissolve well in nonpolar solvents through dispersion forces alone, but dissolve poorly in water, because inserting a nonpolar molecule into water disrupts water's hydrogen-bonding network without providing any comparably strong replacement interaction.

This principle directly explains the boiling-point gap between constitutional isomers with different functional groups: 1-butanol (boiling point approximately 118°C) can hydrogen bond through its hydroxyl group, while its constitutional isomer diethyl ether (boiling point approximately 35°C, same molecular formula C4H10O) has an oxygen with no attached hydrogen, so it can only accept, never donate, a hydrogen bond, leaving it with much weaker net intermolecular forces overall (comparable to pentane, boiling point 36°C, which has similar molecular weight but only dispersion forces available).

Amphipathic molecules — those with both a polar or charged region and a nonpolar region in the same molecule — are central to biological structure because of how intermolecular forces sort their two halves. Phospholipids have a polar phosphate head group that hydrogen bonds and forms dipole interactions with surrounding water, and a nonpolar fatty-acid tail that cannot participate in those interactions. In water, phospholipids spontaneously assemble into bilayers, with polar heads facing outward toward water and nonpolar tails clustering together in the interior, driven by the hydrophobic effect: water preferentially maximizes its own hydrogen bonding, so nonpolar tails are excluded from the aqueous environment and pushed together, an entropically favorable rearrangement of water molecules rather than a direct attraction between the tails themselves. The same hydrophobic effect drives protein folding, where nonpolar amino acid side chains cluster in a folded protein's interior, away from surrounding water, while polar and charged side chains remain on the solvent-exposed surface, frequently forming hydrogen bonds with water or with each other.