What forces act between molecules, and how do they explain physical properties?
Van der Waals (London) forces, permanent dipole-dipole forces and hydrogen bonding, how they arise from electronegativity and polarity, and how they explain boiling points, solubility and the anomalous behaviour of water.
A CCEA A-Level Chemistry answer on intermolecular forces, covering van der Waals (London) forces, permanent dipole-dipole forces and hydrogen bonding, how polarity and electronegativity create them, and how they explain trends in boiling point, solubility and the anomalous properties of water and ice.
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What this dot point is asking
CCEA wants you to identify and explain the three intermolecular forces, link them to electronegativity and molecular polarity, rank their relative strengths, and use them to explain physical properties such as boiling point, solubility and the anomalous behaviour of water and ice.
The three intermolecular forces
Van der Waals forces arise because electrons move randomly, creating an instantaneous dipole that induces a dipole in a neighbouring molecule. They get stronger as the number of electrons (and surface area of contact) increases, which is why boiling points rise down a homologous series.
Permanent dipole-dipole forces act in addition to van der Waals forces whenever a molecule has a permanent dipole. The slightly positive end of one molecule is attracted to the slightly negative end of its neighbour. This is why a polar molecule such as has a higher boiling point than a non-polar molecule of similar electron count such as : both have van der Waals forces, but only has the extra dipole-dipole attraction. The strength of these forces depends on how polar the molecule is, which in turn depends on the electronegativity difference and the molecular shape.
Polarity and electronegativity
So has polar bonds but is non-polar overall because its linear shape makes the dipoles cancel, while is polar because its bent shape means they do not.
Hydrogen bonding and the anomaly of water
A hydrogen bond forms between a lone pair on N, O or F and a hydrogen atom bonded to N, O or F in another molecule. It is the strongest intermolecular force and explains why water, ammonia and hydrogen fluoride have anomalously high boiling points compared with other hydrides in their groups.
Effects on properties
- Boiling point: stronger intermolecular forces mean more energy is needed to separate molecules, so higher boiling points.
- Solubility: substances dissolve best in solvents with similar polarity ("like dissolves like"); polar and hydrogen-bonding molecules dissolve in water.
Examples in context
Example 1. Why detergents lift grease. A detergent molecule has a polar (hydrogen-bonding) head and a long non-polar hydrocarbon tail. The non-polar tail mixes with grease through van der Waals forces, while the polar head hydrogen-bonds to water. The molecule therefore bridges two phases that would not otherwise mix, lifting grease into the wash water. This is a direct, everyday application of "like dissolves like" reasoning built on the three intermolecular forces.
Example 2. The high surface tension of water. Water striders walk on ponds because water has an unusually high surface tension, caused by extensive hydrogen bonding pulling the surface molecules inward. The same hydrogen bonding gives water its high specific heat capacity, making it an effective coolant in car engines and a temperature buffer in living organisms. CCEA expects candidates to trace these macroscopic properties back to the strength and directionality of hydrogen bonds.
Try this
Q1. Name the strongest type of intermolecular force and state the condition for it. [2 marks]
- Cue. Hydrogen bonding; hydrogen must be bonded to N, O or F.
Q2. Explain why ice is less dense than liquid water. [2 marks]
- Cue. Hydrogen bonds hold molecules in an open lattice, spacing them further apart than in the liquid.
Exam-style practice questions
Practice questions written in the style of CCEA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
CCEA 20196 marksExplain why the boiling points of the hydrides of Group VI increase in the order H2S < H2Se < H2Te, yet water (H2O) has a boiling point far higher than all three.Show worked answer →
Markers want the van der Waals trend explained first, then the anomaly of water explained separately.
For , and the dominant intermolecular force is van der Waals (London) forces. As the molecules get larger down the group they have more electrons, so the instantaneous and induced dipoles are larger and the van der Waals forces are stronger. More energy is needed to separate the molecules, so the boiling point rises in that order.
Water breaks the trend because oxygen is small and highly electronegative, so water molecules form hydrogen bonds (O bonded to H, with a lone pair to accept). Hydrogen bonds are much stronger than van der Waals forces, so far more energy is needed to separate water molecules, giving it an anomalously high boiling point.
Markers reward the link between number of electrons and van der Waals strength for the trend, and the specific condition for hydrogen bonding (H bonded to O and a lone pair) for the anomaly.
CCEA 20213 marksAmmonia is very soluble in water. Explain this solubility in terms of intermolecular forces, and draw a diagram showing a hydrogen bond between an ammonia molecule and a water molecule.Show worked answer →
The marks are for identifying hydrogen bonding and showing it correctly.
Ammonia dissolves readily in water because both molecules can form hydrogen bonds with each other: the lone pair on the nitrogen of ammonia accepts a hydrogen from water, and the lone pairs on water oxygen accept the slightly positive hydrogens of ammonia. These hydrogen bonds between solute and solvent release energy that offsets the energy needed to separate the molecules, so dissolving is favourable.
The diagram should show , with the hydrogen bond drawn as a dashed line from a lone pair on the nitrogen to the hydrogen of water (or from a lone pair on the oxygen to a hydrogen of ammonia). The bond angle at the hydrogen bond should be roughly linear.
Markers reward naming hydrogen bonding, explaining that solute-solvent hydrogen bonds form, and drawing the dashed bond from a lone pair to a slightly positive hydrogen.
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Sources & how we know this
- CCEA GCE Chemistry specification — CCEA (2016)