Introduction to Intermolecular Forces分子间作用力入门

Why does water boil at 100 °C?

Water (H₂O) and hydrogen sulfide (H₂S) have the same bent shape, and the H₂S molecule is larger and heavier. Yet water boils at 100 °C, while hydrogen sulfide boils at −59.6 °C and is a gas at room temperature. As lesson 2.4.2 pointed out, boiling water doesn’t break its O–H bonds; it just separates the molecules from each other. So a boiling point depends on the attractions between molecules.

Intermolecular forces

Molecules attract one another through intermolecular forces分子间作用力. These are much weaker than chemical bonds: covalent bonds usually take hundreds of kJ/mol to break, intermolecular forces only a few to a few tens of kJ/mol.

• When a substance melts or boils, it is the intermolecular forces that are overcome; the molecules themselves are unchanged. Only a chemical reaction breaks chemical bonds.
• For a substance made of molecules, the stronger the intermolecular forces, the higher its melting and boiling points.

American and Chinese textbooks divide these forces differently. US Honors and AP courses use three types: London dispersion forces色散力, dipole–dipole forces偶极–偶极作用, and hydrogen bonding氢键. Chinese textbooks call the general attraction between all molecules van der Waals forces范德华力, which roughly covers the first two, and treat hydrogen bonding separately. (American books sometimes use “van der Waals forces” in this broad sense too.)

London dispersion forces

There are dispersion forces between all molecules, nonpolar ones included. The electrons in a molecule are always moving, and at any instant they may bunch up on one side, creating an instantaneous dipole. That dipole pushes the electrons of a neighboring molecule to one side, inducing a dipole there, and the two attract.

• The larger the molecule and the more electrons it has, the stronger the dispersion forces. A large electron cloud is easier to distort, or more polarizable. That is why the halogens go from gas (F₂) to solid (I₂) down the group (lesson 2.3.6), and why noble-gas boiling points rise from helium to radon.
• The more surface two molecules share, the stronger the dispersion forces. Long, stretched-out molecules pack against each other more closely than nearly spherical ones.

Dipole–dipole forces

Polar molecules (lesson 2.4.7) have a partially positive end and a partially negative end. The δ+ end of one molecule attracts the δ− end of another: this is a dipole–dipole force. Among molecules of similar size, polar molecules attract each other more strongly than nonpolar ones and have higher boiling points.

Hydrogen bonding

When hydrogen is bonded to nitrogen, oxygen, or fluorine, which are small and very electronegative, the shared electrons are pulled strongly away from hydrogen, leaving it almost a bare proton with a large partial positive charge. That hydrogen is strongly attracted to a lone pair on an N, O, or F atom in another molecule. This attraction is a hydrogen bond, often written X–H···Y (X and Y are N, O, or F; the dots are the hydrogen bond).

• A hydrogen bond is not a chemical bond; it is an especially strong intermolecular force. It is stronger than dispersion and dipole–dipole forces but much weaker than a covalent bond.
• Hydrogen bonding gives NH₃, H₂O, and HF unusually high boiling points compared with the other hydrides in their groups.
• In ice, each water molecule is hydrogen-bonded to 4 neighbors in an open network with lots of empty space, so ice is less dense than liquid water and floats.

Use boiling-point data to compare the three forces: click the hydrides in the first plot, then compare three substances of similar molar mass, and alkanes of different sizes and shapes:

How to compare boiling points

To compare the boiling points of molecular substances, think in this order:

1. Is there hydrogen bonding? Molecules with H bonded to N, O, or F boil much higher.
2. How big are the molecules? For similar structures, the larger the molecule (more electrons, larger molar mass), the stronger the dispersion forces and the higher the boiling point.
3. If the sizes are similar, compare polarity and shape. Polar molecules boil higher than nonpolar ones, and long molecules boil higher than nearly spherical ones.

This is only a rough order, because dispersion forces can be strong in big molecules. Octane (C₈H₁₈) has only dispersion forces, yet it boils at 125.6 °C, higher than water.

Example 1: H₂O versus H₂S

Why does water (100 °C) boil so much higher than hydrogen sulfide (−59.6 °C)?

Example 2: Boiling points of the halogens

Rank F₂, Cl₂, Br₂, and I₂ from lowest to highest boiling point, and explain.

Common mistakes

Quick check

Summary

• Intermolecular forces are much weaker than chemical bonds. Melting and boiling overcome intermolecular forces; only chemical changes break chemical bonds.
• American textbooks divide intermolecular forces into London dispersion forces, dipole–dipole forces, and hydrogen bonding; Chinese textbooks group the first two as van der Waals forces. Dispersion forces act between all molecules and grow with size; polar molecules add dipole–dipole forces; H bonded to N, O, or F allows hydrogen bonding.
• To compare boiling points, check for hydrogen bonding, then size, then polarity and shape. Hydrogen bonding makes the boiling points of NH₃, H₂O, and HF unusually high and makes ice less dense than water.