UNIT 2 CHEMISTRY · PEARSON AAQ · B1.7
Intermolecular Forces
FREE CHEMISTRY REVISION · TOPIC 16
Explain physical properties through attractions between molecules.
Identify temporary and permanent dipoles, hydrogen-bond donors and acceptors, and compare the total attractions that control molecular behaviour.
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · B1.7
What you will learn
Explain London dispersion, permanent dipole–dipole attractions and hydrogen bonding; identify all the forces present in a molecular substance; use them to explain differences in boiling point, volatility, viscosity and related physical properties.
1. Between molecules versus within molecules
Intermolecular forces act between molecules. Covalent bonds hold atoms together within each molecule. When a simple molecular liquid boils, its molecules separate; the covalent bonds within them usually remain intact. More energy is needed to separate molecules when their attractions are stronger, giving a higher boiling point at the same external pressure.
Melting disrupts the ordered solid arrangement, but molecules remain close together in a liquid and still attract one another. Melting points depend strongly on crystal packing as well as attraction strength, so boiling-point trends are usually easier to compare directly. Giant covalent and ionic solids need different bonding explanations; do not apply a molecular-force model indiscriminately.
2. London dispersion: temporary dipole–induced dipole
Electrons are always moving. At an instant an atom or molecule may have an uneven electron distribution, producing a temporary (instantaneous) dipole. Its electric field distorts a neighbouring electron cloud, producing an induced dipole. The correlated distributions give a net attraction, called London dispersion.
London forces occur between all atoms and molecules, including non-polar particles, polar molecules and substances that hydrogen bond. They do not require a permanent dipole. The fluctuations continue: do not draw the temporary partial charges as permanently fixed on a non-polar molecule.
A larger, more easily distorted electron cloud is more polarisable, generally giving stronger London attractions. Within comparable families, more electrons and larger size often accompany higher boiling points. Relative molecular mass is a useful correlation, but mass itself is not the electrostatic cause. Molecular shape also matters: elongated molecules can have greater surface contact than compact, highly branched isomers.
3. Permanent dipole–permanent dipole interactions
Polar molecules already have a permanent separation of charge because their bond dipoles do not cancel. In favourable orientations the δ+ end of one attracts the δ− end of another. These are permanent dipole–dipole attractions. Thermal motion continually changes orientations; a liquid is not an immobile array of neatly aligned molecules.
These attractions act in addition to London forces. For substances of similar size and shape, a permanent dipole can increase the total attraction and boiling point. A large non-polar molecule can nevertheless have stronger total attractions than a much smaller polar molecule. Polarity alone cannot rank every boiling point.
4. Hydrogen bonding: donor, hydrogen and acceptor
For the common cases in this specification, hydrogen bonding involves H covalently bonded to N, O or F attracted to an available lone pair on N, O or F in a neighbouring molecule. The strongly polar N–H, O–H or F–H bond leaves H δ+. The small, electron-rich acceptor attracts it. A hydrogen bond is directional and is conventionally drawn as a dotted or dashed line from H to the acceptor atom.
Water, ammonia, hydrogen fluoride and alcohols can hydrogen bond between their own molecules. In ethanol, C₂H₅OH, the O–H group supplies a donor and the oxygen lone pairs supply acceptors. A molecule merely containing H and O somewhere is insufficient: the hydrogen must be attached to the appropriate donor atom.
Methoxymethane, CH₃OCH₃, has oxygen lone pairs but no O–H bond. It can accept hydrogen bonds from water, but cannot form the corresponding conventional hydrogen-bond network with other methoxymethane molecules. HCl does not meet the usual N/O/F donor condition used here. These distinctions are more useful than the rule “has hydrogen, so hydrogen bonds”.
| Substance | London | Permanent dipole–dipole | Hydrogen bonding between its own molecules |
|---|---|---|---|
| CH₄ | Yes | No: symmetric and non-polar | No |
| HCl | Yes | Yes | No in the usual AAQ N/O/F model |
| CH₃OCH₃ | Yes | Yes | No O–H/N–H/F–H donor |
| C₂H₅OH | Yes | Yes | Yes: O–H donor and O lone pairs |
| H₂O | Yes | Yes | Yes: extensive network |
5. Explain physical properties with a controlled comparison
Boiling point: stronger total intermolecular attractions require more energy to separate molecules into the gas. At a fixed external pressure, boiling occurs when the vapour pressure equals that pressure. Volatility: a more volatile liquid escapes more readily into the gas; weaker attractions generally give greater volatility and higher vapour pressure at a fixed temperature.
Viscosity: stronger attractions and molecular entanglement can resist flow. Hydrogen-bond networks can increase viscosity, but molecule size, shape and temperature also matter. Surface tension: cohesive attractions contribute to the energetic cost of making a liquid surface. Water is developed in B1.8.
Solubility: dissolving replaces solute–solute and solvent–solvent attractions with solute–solvent attractions. Hydrogen bonding with water can favour mixing, while a large non-polar region can reduce water solubility. “Like dissolves like” is a guide, not a complete explanation; the balance of interactions and entropy matters.
For similar-sized species, hydrogen bonding is often especially influential, followed by additional permanent-dipole attractions. Do not use a universal ranking of whole substances based only on the named force: numerous London interactions in a large molecule can outweigh the attractions in a smaller hydrogen-bonded one. Compare structure, size, polarizability, shape and available donors/acceptors.
Worked examples
Example 1 · Ethane versus butane
Both are essentially non-polar and have London forces. Butane has more electrons and a larger, more polarisable cloud. Its total attractions are stronger, so more energy is required for vaporisation: butane has the higher boiling point. Do not claim that butane gains hydrogen bonding.
Example 2 · Pentane isomers
n-Pentane and highly branched 2,2-dimethylpropane have the same formula and electron count. The elongated n-pentane shape allows greater intermolecular contact than the compact branched shape. London attractions are more effective and n-pentane has the higher boiling point. Shape explains a difference that mass cannot.
Example 3 · Ethanol versus methoxymethane
Both have formula C₂H₆O, so their similar mass and electron count make this a useful comparison. Ethanol has an O–H donor and O lone pairs, allowing intermolecular hydrogen bonds. Methoxymethane lacks that donor. Ethanol has stronger total attractions and a higher boiling point.
Example 4 · Identify forces in HCl
Cl is more electronegative than H, so the molecule has a permanent dipole. HCl therefore has permanent dipole–dipole attractions and London forces. It does not form the usual strong N/O/F hydrogen-bond network. State all present forces rather than only the strongest named one.
Example 5 · Explain a phase change
When liquid ethanol boils, energy overcomes intermolecular attractions, including hydrogen bonding. Ethanol molecules enter the gas without needing to break all their C–C, C–H or O–H bonds. Breaking those covalent bonds would change chemical identity, not simply vaporise the liquid.
Common misconceptions and exam tips
“Non-polar particles have no attractions.” They have London forces.
“Polar molecules do not have London forces.” London forces are present in every molecular substance.
“Any H-containing molecule hydrogen bonds.” Identify the appropriate donor bond and acceptor lone pair.
“Boiling breaks the covalent bonds.” In molecular vaporisation the molecules largely remain intact.
“Hydrogen bonding always beats London forces.” Compare total attractions and molecular size; do not apply an absolute ranking.
For comparison questions name the force, explain its structural origin, compare total attraction strength, then link to energy and the observed property. For a drawn hydrogen bond label δ+, δ− and the donor/acceptor atoms, and place the dotted line H···acceptor, not between two hydrogens.
Your revision summary
London forces arise from fluctuating and induced dipoles and act in all molecular substances. Polar molecules additionally experience permanent dipole attractions. Appropriate N–H, O–H or F–H donors and lone-pair acceptors allow hydrogen bonding. Physical properties depend on total attractions, molecular size, shape and, particularly for melting, packing.
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