Part 4: Polarity & intermolecular forces

1. Electronegativity

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. It is about the bonding pair, not a lone pair.

• Across a period it increases: nuclear charge rises, atomic radius falls and shielding stays similar, so the bonding pair is attracted more strongly.
• Down a group it decreases: the atom is larger and there is more shielding, so the bonding pair is further from the nucleus and attracted less.

Fluorine is the most electronegative element. Oxygen, nitrogen and chlorine also have high electronegativities.

2. Polar bonds

When two bonded atoms have different electronegativities, the bonding pair is shared unequally. The more electronegative atom gets a partial negative charge (δ−) and the other a partial positive charge (δ+): H^δ+–Cl^δ−. Bonds between identical atoms (Cl–Cl) are non-polar.

Bonding is a continuum: the bigger the electronegativity difference, the more polar the bond, until it becomes more ionic in character. There is no single universal cut-off value that separates "covalent" from "ionic".

3. Polar and non-polar molecules

A molecule can contain polar bonds and still be non-polar. A molecule is non-polar when its bond dipoles cancel as vectors. Equal dipoles cancel in arrangements such as linear CO₂, trigonal planar BF₃ and tetrahedral CCl₄. In bent H₂O and pyramidal NH₃ they do not cancel. It is the dipoles that cancel, not the bonds.

Do not say "molecules with lone pairs are polar": XeF₄ shows this is false.

4. London forces

Electrons are always moving. At any instant, the electron cloud in a molecule may be uneven, creating an instantaneous dipole. This induces a dipole in a neighbouring molecule, and the opposite partial charges attract. These are London (dispersion) forces. They act between all molecules, including non-polar ones and noble-gas atoms.

London forces are stronger when:

• the molecule has more electrons, making its electron cloud more easily distorted;
• the molecule has a larger contact surface area. Unbranched molecules can make more extensive contact than compact branched isomers of the same formula, so they have stronger London forces and higher boiling points.

Down Group 17, from F₂ to I₂, the molecules have more electrons, so London forces and boiling points increase. Boiling points also rise along a homologous series as the chain gets longer.

5. Permanent dipole–dipole forces

Polar molecules have permanent dipoles. The δ+ end of one molecule attracts the δ− end of a neighbour. These forces act in addition to London forces, never instead of them.

6. Hydrogen bonding

A hydrogen bond forms when hydrogen is covalently bonded to N, O or F, and is attracted to a lone pair on an N, O or F atom in another molecule.

• The N–H, O–H or F–H bond is very polar, so the H is strongly δ+.
• Draw the hydrogen bond as a dotted line from the δ+ H to the lone pair on the N, O or F of the other molecule.
• Show the δ+ and δ− charges and the lone pair. The X–H···Y arrangement is close to linear (about 180°).

Donors and acceptors

Ethanol (CH₃CH₂OH) has an O–H group and lone pairs on O, so it can hydrogen bond with itself. Ethanal (CH₃CHO) has lone pairs on O but no H bonded to O, so it cannot hydrogen bond with itself; it can still accept hydrogen bonds from water.

7. Comparing boiling points

Don't apply a fixed ranking such as "hydrogen bonds > dipoles > London" when the molecules differ in size. Compare the total intermolecular forces.

8. Ice and anomalous hydrides

H₂O, NH₃ and HF have much higher boiling points than you would expect from the trend in the heavier hydrides of their groups (H₂S, PH₃, HCl). This is because they form hydrogen bonds, which need extra energy to overcome.

In ice, each water molecule forms up to four hydrogen bonds: two through its own H atoms and two to its lone pairs. This makes an open tetrahedral network, with molecules held further apart on average than in liquid water. So ice is less dense than water. On melting, some hydrogen bonds break and the network partly collapses, so the molecules pack more closely. (This tetrahedral arrangement between molecules is separate from the 104.5° angle inside each water molecule.)

9. The charged-rod test

A thin stream of a polar liquid (such as water) bends towards a charged rod: the permanent dipoles turn so that their oppositely charged end faces the rod, and are attracted. A non-polar liquid such as cyclohexane is barely affected. This is a simple observation, not a precise measure of how polar a substance is.

10. Quick checks

Finesse practice: indicative answers to check your reasoning, not official mark allocations.

11. Sources