Part 3: Polarity, intermolecular forces and water

Use electronegativity and three-dimensional shape to predict dipoles, then explain boiling and solubility with the correct forces.

Bond polarity and molecular polarity are different

Electronegativity measures an atom’s ability to attract the bonding electron pair in a covalent bond. On the Pauling scale it generally rises across a period and falls down a group; fluorine is the most electronegative element. A difference pulls the pair towards one atom, giving δ− there and δ+ at the other end. These are partial, not full ionic, charges.

A molecule is polar if its bond dipoles do not cancel in three dimensions. Linear CO₂ and tetrahedral CCl₄ have cancelling dipoles; bent H₂O and CH₃Cl do not. Having polar bonds alone does not prove a polar molecule. Bonding spans degrees of ionic character, so an electronegativity difference is a guide rather than an absolute universal cutoff.

Add bond dipoles using the geometry

CO₂ has two polar C=O bonds. They point in opposite directions along a straight line and cancel, so the molecule is non-polar. Water also has polar bonds, but they meet at an angle; their effects do not cancel, giving an overall dipole. A drawing makes the difference clearer than the word “symmetrical” on its own.

CCl₄ has four identical polar bonds arranged tetrahedrally; their vector sum is zero. Replacing one Cl with H changes the dipole contributions, so CHCl₃ is polar even though the carbon still has tetrahedral geometry. Geometry and bond identities both matter.

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London forces, permanent dipoles and hydrogen bonds

London forces arise when fluctuations in electron density create instantaneous dipoles that induce dipoles nearby. They occur between all atoms and molecules. Larger, more polarisable electron clouds and greater intermolecular contact generally strengthen them. Polar molecules additionally have permanent dipole–dipole attractions. OCR groups these two types under van der Waals forces.

Hydrogen bonding involves H covalently bonded to N, O or F attracting a lone pair on N, O or F in a neighbouring molecule (or sometimes another part of the same molecule). In water, label Oδ−–Hδ+ and show a dotted attraction from Hδ+ to an oxygen lone pair on another molecule. H bonding is not an additional covalent O–H bond.

Compare similar-sized molecules to isolate the effect of hydrogen bonding. Do not claim a small hydrogen-bonding molecule must boil above every much larger molecule: overall London interactions also matter.

Choose the explanation that distinguishes the pair

Compare ethanol and dimethyl ether: both are C₂H₆O with the same electron count. Ethanol has an O–H donor and oxygen lone pairs, so its molecules can hydrogen bond to each other. The ether has oxygen lone pairs but no O–H/N–H/F–H donor, so it cannot make the same hydrogen-bond network in a pure sample. It still has London forces and permanent dipole attractions.

Now compare the ether with water as solvent. Its oxygen can accept a hydrogen bond from a water molecule, so “cannot hydrogen bond to itself” does not mean “cannot hydrogen bond with water”. Identify the donor and acceptor for the actual pair of molecules.

For pentane versus 2,2-dimethylpropane, both have the same formula and electron count, so a more-electrons argument cannot distinguish them. Pentane offers more intermolecular contact and generally stronger overall London attractions, giving the higher boiling point. For Cl₂ versus I₂, molecular size and polarisation instead provide the useful distinction.

Water’s unusual properties and molecular solubility

Water has a relatively high boiling point for such a small molecule because substantial energy is needed to overcome intermolecular hydrogen bonding. In ice the hydrogen-bonded open arrangement holds molecules farther apart on average than in liquid water, so ice is less dense. Each water molecule can participate as two H-bond donors and two acceptors in the ideal ice network.

A molecular solute dissolves favourably when new solvent–solute interactions compensate for interactions disrupted in both substances. Small alcohols can hydrogen bond to water and mix well; long hydrocarbon sections reduce water compatibility. Non-polar substances often dissolve better in non-polar solvents. Ionic solubility depends on lattice and hydration energetics, so avoid an “all polar things dissolve” rule.

Draw the hydrogen bond to the electron pair

Draw both covalent water molecules first. Put δ− on each O and δ+ on its H atoms. Show two oxygen lone pairs. The dotted intermolecular line must run from a hydrogen bonded to oxygen in one molecule to an oxygen lone pair in the neighbouring molecule. It must be visually different from the solid covalent O–H bonds.

Each water molecule can donate through two H atoms and accept at two lone pairs in the ideal ice network. The open arrangement holds molecules farther apart on average than in liquid water, reducing density. Do not explain this using stretched covalent bonds, air in the gaps or weaker London forces.

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Quick checks

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

Sources