Part 4: Polarity and intermolecular forces

Distinguish an unequal bond from an overall molecular dipole, then identify the attractions between complete molecules.

Electronegativity describes attraction within a bond

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. It is not the energy required to remove an electron. If two bonded atoms have different electronegativities, electron density is pulled towards the more electronegative atom, which becomes δ−; the other becomes δ+.

In H–Cl, chlorine is more electronegative and the bond is Hδ+–Clδ−. Delta means partial charge: hydrogen has not become a free H⁺ ion simply because the bond is polar. Electronegativity generally rises across a period and falls down a group because nuclear attraction competes with increasing distance and shielding.

Ionic and covalent bonding are limiting models

Bonding forms a continuum from nearly equal sharing to strongly unequal electron distribution and ionic character. Small electronegativity differences favour covalent character; larger differences generally increase polarity and ionic character. There is no universal numerical cutoff that perfectly classifies every compound in every physical state.

For example, sodium chloride is well described by a giant ionic lattice, whereas silicon tetrachloride consists of covalent molecules. Aluminium chloride shows substantial covalent character and can form Al₂Cl₆ molecules in appropriate conditions. Its behaviour cannot be inferred by assuming every metal–non-metal compound has identical purely ionic bonding.

In 8CH0/01 June 2023 Q5(c–d), the evidence needed both ends of the continuum and a link from each selected substance’s bonding to the supplied properties. A formula containing a metal is a clue, not a replacement for the conductivity and phase-change data.

Bond dipoles add as vectors

A molecule is polar when its bond dipoles and electron distribution produce a net permanent dipole. Work out its shape before deciding. CO₂ has polar C=O bonds, but the molecule is linear and the equal opposite bond dipoles cancel. CCl₄ is tetrahedral with four equal C–Cl bond dipoles that cancel in three dimensions.

H₂O is bent, so its O–H dipoles do not cancel; water is polar. NH₃ is pyramidal and polar. CH₃Cl is tetrahedral but its substituents are not all the same, so its dipoles do not cancel. The presence of a lone pair is not itself a universal test: square-planar XeF₄ can be non-polar despite central lone pairs.

A charged rod can deflect a thin water stream because the field attracts the polar liquid through orientation and polarisation. Use a narrow, reproducible stream and the same rod distance when comparing samples. Treat deflection as supporting evidence, not a unique chemical identification; induced polarisation and experimental conditions can also matter.

London forces arise from fluctuating electron distributions

At any instant, an atom or molecule’s electron distribution can be uneven, creating an instantaneous dipole. This changes the electron distribution in a neighbouring particle and induces another dipole. Attraction between the correlated partial charges produces London forces, also called instantaneous dipole–induced dipole attractions.

All atoms and molecules have electrons and therefore London attractions, including non-polar molecules and noble-gas atoms. Larger, more easily polarised electron clouds generally produce stronger attractions; surface contact between molecules also matters. A molecule does not need a permanent dipole to have a boiling temperature above absolute zero.

For a similar family such as alkanes, more electrons and a larger polarisable cloud tend to increase boiling temperature. This does not mean molar mass itself is a force, or that an instantaneous dipole becomes a permanent ion.

Polar molecules have an additional attraction

Permanent dipole–dipole attractions occur between the δ+ region of one polar molecule and the δ− region of another. Polar molecules also have London forces; identifying one interaction does not remove the others.

Compare similarly sized molecules before concluding that polarity raises boiling temperature. A very large non-polar molecule can have stronger total intermolecular attraction than a small polar molecule because its London contribution is large. The simplistic ranking ‘all permanent-dipole forces exceed all London forces’ is therefore unsafe.

Hydrogen bonding needs a donor and an acceptor

At this level, a hydrogen bond forms between a δ+ hydrogen covalently bonded to N, O or F and a lone pair on N, O or F in another molecule or another part of a molecule. The N–H, O–H or F–H bond is strongly polar. The hydrogen bond is drawn as a dashed attraction from that hydrogen towards the accepting lone pair; it is not the O–H covalent bond itself.

Water can donate through its two O–H hydrogens and accept through oxygen’s lone pairs. Liquid ammonia has N–H donors and a nitrogen lone pair; liquid HF has an F–H donor and fluorine lone pairs. For analogous compounds, methanol and methylamine can hydrogen bond to themselves. Methoxymethane has oxygen lone pairs but no O–H bond: it can accept hydrogen bonds from water but does not provide the usual donor for self hydrogen bonding.

A clear drawing includes δ+ on the donor hydrogen, δ− on the N/O/F atoms, the acceptor lone pair and a dashed intermolecular link. The donor bond and hydrogen-bond direction are often close to linear, but do not justify that geometry by claiming hydrogen has two ordinary covalent electron pairs.

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Quick checks

Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.

Sources