Intermolecular Forces分子间作用力

From molecules to the properties of substances

Unit 2 looked at the chemical bonds inside a single molecule or crystal. Unit 3 takes a different view: when many particles come together, how do they attract one another, and how do those attractions set a substance’s boiling point, vapor pressure, and solubility? Lesson 2.4.8 introduced dispersion forces, dipole–dipole forces, and hydrogen bonding. AP topic 3.1 builds on that: identify which interactions act between any two particles, and explain their strengths from the structure of the particles.

All Coulombic attractions

Every intermolecular force is a Coulombic attraction between charged regions: ions carry whole-number charges, polar molecules carry partial charges, and even nonpolar molecules develop instantaneous dipoles from moment to moment. Larger charges and closer approach mean stronger attraction. All of these are much weaker than chemical bonds.

For particles of similar size, the order is roughly: ion–dipole > hydrogen bonding > dipole–dipole > dipole–induced dipole and dispersion. But dispersion forces grow quickly as particles get larger, and between large molecules they can be the strongest interaction.

AP note: don’t use “London dispersion forces” and “van der Waals forces” as synonyms. The term van der Waals forces means different things in different books, and sometimes covers all intermolecular forces.

Dispersion forces and polarizability

The ease with which an electron cloud can be distorted, and so develop instantaneous and induced dipoles, is called polarizability可极化性. The more electrons and the larger the electron cloud, the greater the polarizability and the stronger the dispersion forces:

For the same reason, $\ce{CCl4}$ is a liquid at room temperature while $\ce{CI4}$, with far more electrons, is a solid.

Dispersion forces also depend on the contact area接触面积 between molecules. The three isomers of pentane are all $\ce{C5H12}$ with the same number of electrons: long-chain n-pentane (36.1 °C) packs most closely and boils highest; ball-shaped neopentane (9.5 °C) has the least contact area and boils lowest; isopentane (27.8 °C) is in between. (Boiling points: CRC.)

Extra attractions from polarity

• Dipole–dipole forces: between polar molecules, the δ+ end of one molecule attracts the δ− end of another. These act in addition to dispersion forces, so for molecules of similar size, polar substances boil higher than nonpolar ones: propane −42.1 °C, dimethyl ether −24.8 °C. Isomers work the same way: cis-1,2-dichloroethene is polar and the trans isomer is nonpolar (U2.7), so the cis isomer has the higher boiling point and lower vapor pressure. $\ce{CH2F2}$ is polar and $\ce{CF4}$ is nonpolar; even though $\ce{CF4}$ has more electrons, $\ce{CH2F2}$ boils higher.
• Dipole–induced dipole forces: a polar molecule distorts the electron cloud of a nearby nonpolar molecule and induces a dipole in it. The larger the dipole of the polar molecule and the more polarizable the nonpolar molecule, the stronger the attraction. These act between water and $\ce{O2}$ or between water and benzene, which is why a little $\ce{O2}$ dissolves in water.
• Ion–dipole forces: an ion attracts the oppositely charged end of a polar molecule: $\ce{Na+}$ attracts the O end of water (δ−), and $\ce{Cl-}$ attracts the H end (δ+). Ions carry whole-number charges, so ion–dipole forces are generally stronger than dipole–dipole forces. By Coulomb’s law, the larger the ion’s charge and the smaller its radius, the more strongly it attracts water: $\ce{Mg^2+}$ (72 pm) > $\ce{Ca^2+}$ (100 pm) > $\ce{Ba^2+}$ (135 pm) (Shannon radii). Working backward, which end of the water molecules faces an ion tells you whether the ion is positive or negative. These forces are what allows ionic compounds to dissolve in water (U3.7).

Hydrogen bonding: who gives, who takes

A hydrogen bond needs two things: an H-bond donor氢键给体 (H bonded to N, O, or F) and an H-bond acceptor氢键受体 (a lone pair on N, O, or F).

• Pure dimethyl ether has no hydrogen bonding: all its H atoms are bonded to C, so it has no donor. But its O has lone pairs that can accept an H from water, so dimethyl ether and water do hydrogen-bond to each other.
• Ethanol (78.3 °C) and dimethyl ether (−24.8 °C) share a molecular formula yet boil more than 100 degrees apart: ethanol has an O–H and hydrogen-bonds to itself (the isomers of U2.5).
• A hydrogen bond is an attraction between molecules, not the covalent O–H bond inside a molecule. In a particle diagram of water, draw the hydrogen bond from an H atom on one molecule to an O atom on a neighboring molecule, usually as a dashed line.
• Hydrogen bonds can also form between different parts of the same large molecule. The two strands of DNA pair up through hydrogen bonds between bases (2 between A and T, 3 between G and C), and proteins fold into specific shapes through hydrogen bonds, dispersion forces, and other noncovalent interactions between parts of the molecule.

Choose two particles to see which interactions act between them and which ends face each other in the particle diagram:

Intermolecular forces and properties

When a liquid boils, its molecules must separate completely, overcoming the intermolecular forces. So:

• stronger forces mean a higher boiling point;
• stronger forces mean a lower vapor pressure at a given temperature (fewer molecules escape into the gas phase).

Melting points also tend to rise with stronger forces, but melting only rearranges the molecules, and how well they pack matters too, so the pattern is less clean than for boiling points.

The heat needed to turn 1 mol of a liquid into a gas is the enthalpy of vaporization汽化热 ΔH_vap. The stronger the intermolecular forces, the larger it is. It is always much larger than the enthalpy of fusion: melting only weakens the attractions while the molecules stay in contact, whereas vaporization separates the molecules completely. For water, the enthalpy of fusion is 6.01 kJ/mol and the enthalpy of vaporization is 40.7 kJ/mol (CRC), which is why the boiling plateau of a heating curve is so much longer than the melting plateau (1.2.2).

To compare boiling points, first list every interaction present in each substance, then ask: Is there hydrogen bonding? Is the molecule polar? How many electrons, and what shape? Don’t skip the last question: molecules with many electrons can have dispersion forces that outweigh everything else. $\ce{I2}$ is nonpolar, yet its boiling point (184.3 °C) is far above that of polar HCl (−85.1 °C); octane (125.6 °C) has only dispersion forces yet boils above water.

Test these patterns against boiling-point data:

Example 1: Interactions between two particles

List every interaction between each pair of particles, and say which is strongest: (1) $\ce{CH3OH}$ and $\ce{H2O}$; (2) $\ce{I2}$ and $\ce{H2O}$; (3) $\ce{Na+}$ and $\ce{H2O}$.

Example 2: Why does nonpolar I₂ boil higher than polar HCl?

$\ce{I2}$ boils at 184.3 °C, and HCl at −85.1 °C. HCl molecules have dipole–dipole forces between them, while $\ce{I2}$ has only dispersion forces. Explain this result.

Common mistakes

Quick check

AP-style free response

Summary

• Intermolecular forces are all Coulombic attractions: dispersion forces (all particles), dipole–induced dipole and dipole–dipole forces (polar molecules), hydrogen bonding (H bonded to N, O, or F plus a lone pair on N, O, or F), and ion–dipole forces (ions with polar molecules).
• Dispersion forces grow with polarizability: more electrons, a larger electron cloud, and more contact area mean stronger dispersion forces.
• Stronger forces mean higher boiling points and lower vapor pressures. To compare, list every force first, then consider hydrogen bonding, polarity, and size; dispersion forces can dominate in large molecules.