Intermolecular forces

Intermolecular forces (IMFs) act between molecules. They are much weaker than the covalent bonds within molecules.

• London dispersion forces: between all molecules and atoms. A momentary uneven electron distribution (an instantaneous dipole) induces a dipole in a neighbour, and the two attract.
• Dipole–dipole forces: between polar molecules, those with a net dipole moment. The δ+ end of one molecule attracts the δ− end of another.
• Hydrogen bonding: an especially strong dipole–dipole attraction between an H atom bonded to N, O or F on one molecule and a lone pair on N, O or F of another, as in O–H···O between water molecules.

Ionic compounds are held together by ion–ion attractions, far stronger than any of these: NaCl boils at 1686 K.

Dispersion forces depend on polarizability, how easily an electron cloud is distorted.

• More electrons and larger size give stronger forces: He 4.2 K < Ne 27.1 < Ar 87.3 < Kr 119.9 < Xe 165.0 K, and $\ce{F2}$ 85.0 < $\ce{Cl2}$ 239.1 < $\ce{Br2}$ 332.0 < $\ce{I2}$ 457.6 K.
• Shape and contact area: long, unbranched molecules touch along more of their length than compact, branched ones. For the three $\ce{C5H12}$ isomers: pentane 309.2 K > 2-methylbutane 301.0 K > 2,2-dimethylpropane (neopentane) 282.7 K.
• For large molecules dispersion can outweigh polarity: nonpolar $\ce{CCl4}$ (349.9 K) boils above polar $\ce{CHCl3}$ (334.3 K).

For molecules of similar size, the usual order of strength is dispersion only < dipole–dipole < hydrogen bonding. Even a hydrogen bond (about 20 kJ/mol in water) is small next to a covalent O–H bond (463 kJ/mol).