Use intermolecular attractions to explain chain-length and branching trends, water’s unusual behaviour and why a polar solvent does not dissolve every polar substance.
Phase changes separate particles without changing molecular identity
For a molecular liquid to boil, molecules must overcome intermolecular attractions sufficiently to enter the gas. The covalent bonds inside a molecule generally remain intact. Stronger total attractions require more energy and usually produce a higher boiling temperature and lower volatility at a given temperature.
Volatility means how readily a liquid forms vapour. Compare substances at the same temperature and pressure, and distinguish boiling from decomposition. Saying ‘more bonds’ is not enough: name the relevant intermolecular attraction and explain why it is stronger.
Chain length and branching change London attractions
Along an alkane homologous series, increasing chain length gives more electrons and a larger, more polarisable electron cloud. London attractions become stronger and boiling temperature generally rises. On a graph, put the stated independent variable, such as carbon number, on the horizontal axis and boiling temperature with units on the vertical axis; use the observed data rather than assuming a perfect straight line.
Compare structural isomers with the same molecular formula: a less-branched chain can make more extensive contact with neighbouring molecules. A compact branched molecule has less effective contact and generally weaker London attractions, so it boils at a lower temperature. Branching does not reduce the number of electrons for isomers.
Worked comparison: pentane and 2,2-dimethylpropane both contain 42 electrons per molecule, because C₅H₁₂ gives 5(6) + 12(1) = 42. The lower boiling temperature of the compact isomer is therefore explained by shape/contact rather than fewer electrons.
An alcohol usually boils above an alkane with a similar number of electrons because alcohol molecules form hydrogen bonds as well as London attractions. Ethanol and propane each have 26 electrons, but ethanol has an O–H donor and oxygen lone pairs. Stronger attractions between ethanol molecules reduce its volatility.
HF is anomalous within the hydrogen halides
From HCl to HBr to HI, the increasing size and polarisability of the electron cloud strengthens London attractions and raises boiling temperature overall. Their permanent dipoles do not give the correct trend on their own; HCl is more polar than HI but has the lower boiling temperature.
HF forms hydrogen bonds because hydrogen is bonded directly to fluorine. The additional intermolecular association gives a much higher boiling temperature than the simple HCl-to-HI trend would predict. Do not explain this by the strength of the H–F covalent bond: boiling HF separates molecules, not hydrogen and fluorine atoms.
Hydrogen bonding gives water unusual physical properties
Water has relatively high melting and boiling temperatures for such a small molecule because hydrogen bonding adds substantial attraction between molecules. Energy is needed to disrupt these intermolecular associations. Liquid water still has extensive, changing hydrogen bonding; melting does not remove every hydrogen bond.
In ordinary ice, water molecules form an open hydrogen-bonded network with approximately tetrahedral organisation around each oxygen. Each water can donate two hydrogen bonds and accept two. The open arrangement holds molecules farther apart on average than in liquid water. Some of that open structure collapses on melting, so liquid water is denser and ice floats.
Hydrogen bonds do not make the individual H₂O molecules swell. The density change concerns their arrangement and average spacing. Likewise, describing ice as a giant covalent network would be wrong: the O–H covalent bonds lie within distinct water molecules, which are connected by intermolecular hydrogen bonding.
Water hydrates ions, but not every ionic solid dissolves
Dissolution of an ionic solid requires disruption of lattice attractions and rearrangement of water–water interactions. New ion–dipole attractions form when water surrounds the ions: oxygen’s δ− end points towards a cation, while the δ+ hydrogen ends point towards an anion. This process is hydration.
Whether much dissolves depends on the balance of these interactions and the overall tendency to mix. Some ionic solids, such as sodium chloride, dissolve readily; others, such as barium sulfate, are sparingly soluble. Ionic bonding does not guarantee high water solubility. A more detailed energy and entropy treatment appears in Year 13.
When asked why a salt can dissolve, describe attraction between the ion and the appropriate end of water. Do not call the process covalent bonding of all the ions to water or imply hydration changes Na⁺ back into neutral sodium.
New solute–solvent attractions must compensate for disrupted ones
Small alcohols mix well with water because their O–H groups can form hydrogen bonds with water, replacing some water–water and alcohol–alcohol attractions. As the hydrocarbon chain grows, its non-polar region becomes a larger proportion of the molecule, so solubility generally falls.
Halogenoalkanes can have polar C–X bonds but are usually poorly soluble in water. They cannot form sufficiently favourable hydrogen bonds with water to compensate for disrupting the water network and separating solute molecules. Polarity alone is therefore not a guarantee of water solubility.
A non-polar solute such as iodine often dissolves better in a non-polar organic solvent because London attractions between solute and solvent can replace those in the separate substances. ‘Like dissolves like’ summarises this comparison; it does not replace an explanation of which interactions are broken and formed.
Propanone illustrates the donor/acceptor distinction: its carbonyl oxygen can accept hydrogen bonds from water even though pure propanone has no O–H donor for self hydrogen bonding. In a real solvent choice, consider chemical compatibility, separation method, flammability and toxicity as well as solubility. A solvent that reacts with the sample is unsuitable for a simple physical extraction.
Quick checks
Original Finesse questions. Reveal the indicative worked solutions after attempting each question; these are not official Edexcel mark allocations.
Q1. Why does hexane generally boil above butane?Show answer
Hexane has more electrons and a larger, more polarisable cloud, giving stronger London attractions between molecules. More energy is needed to overcome them. The C–C bonds are not broken on boiling.
Q2. Two C₆H₁₄ isomers have different boiling temperatures. Why cannot this be explained by different electron numbers?Show answer
Both have 6(6) + 14 = 50 electrons. Their different branching changes molecular shape and contact between neighbours, changing the total London attraction.
Q3. Explain both HF’s anomalously high boiling temperature and the rise from HCl to HI.Show answer
HF forms hydrogen bonds. HCl, HBr and HI do not meet the normal hydrogen-bonding condition, and their increasing electron-cloud polarisability down the series strengthens London attractions. These are two different parts of the explanation.
Q4. A student says ice floats because freezing breaks all hydrogen bonds. Correct this.Show answer
Ice retains an ordered, open hydrogen-bonded network that spaces molecules farther apart. Liquid water has a less open arrangement and is denser. Some hydrogen bonds are disrupted on melting, but liquid water still hydrogen bonds extensively.
Q5. Predict the orientation of water around Mg²⁺ and explain why that alone does not prove every magnesium salt is soluble.Show answer
Water’s δ− oxygen points towards Mg²⁺, forming ion–dipole attractions. Solubility also depends on the lattice and all other interactions and mixing effects. Strong hydration alone does not establish the complete balance for every salt.
Sources
Sources and examiner guidance (reviewed 9 October 2026)
- Pearson Edexcel 9CH0 specification, Issue 3 — Topic 2, printed pp. 9–11, checked against 8CH0 printed pp. 7–9. Reviewed 9 October 2026.
- Chemrevise — Edexcel Bonding — All 14 pages reviewed as a secondary coverage check. Explanations, worked applications and exercises are original.
Finesse Tuition is not endorsed by AQA or Chemrevise. All explanations and examples here are our own.