A simple molecular substance can have strong covalent bonds inside each molecule and still boil at a low temperature. That is because melting and boiling separate molecules from each other; they do not usually break the covalent bonds inside the molecules.

For Edexcel explanations, keep two levels separate:

The strength and type of attraction between molecules help explain why some simple molecular substances are gases at room temperature while others are liquids or solids. In this lesson, the mark-earning move is usually:

1. identify whether the molecule is non-polar or polar
2. name the strongest relevant intermolecular force
3. explain the electron or dipole interaction causing it
4. link that force to the energy needed to separate molecules

Feynman diagnostic: boiling water

Explain to a younger student why steam still contains H2O molecules, not separate hydrogen and oxygen atoms. A strong answer says that boiling overcomes intermolecular attractions between water molecules; it does not break the covalent O-H bonds within each molecule. The trap answer is "boiling breaks the bonds in water", which is too vague and would lose the bonding-level distinction.

London forces are attractions caused by temporary dipoles. They are present between all atoms and molecules, including non-polar molecules.

The particle story is:

1. Electrons in a molecule are always moving.
2. At one instant, the electron cloud may be unevenly distributed.
3. This creates an instantaneous dipole.
4. The instantaneous dipole distorts the electron cloud in a nearby molecule.
5. This induces a dipole in the neighbouring molecule.
6. The opposite partial charges attract.

This is why Edexcel describes London forces as instantaneous dipole-induced dipole interactions. They are individually weak, but many of them together can matter. London forces tend to be stronger when the electron cloud is larger and more easily distorted, and when molecules can make more surface contact.

Common-error contrast

Non-polar does not mean "no intermolecular forces". A non-polar molecule such as methane, CH4, has no permanent dipole, but it still has London forces because it still has moving electrons.

Permanent dipole-dipole forces act between polar molecules. A polar molecule has a permanent separation of charge, usually because polar bonds are arranged so that their dipoles do not cancel.

The previous lesson's molecular-polarity idea matters here. A bond can be polar without the whole molecule being polar. Carbon dioxide, CO2, has polar C=O bonds, but the molecule is linear and symmetrical, so the bond dipoles cancel. Hydrogen chloride, HCl, is polar because chlorine is more electronegative than hydrogen, producing H(delta+)-Cl(delta-).

Permanent dipole-dipole attractions occur when the delta+ end of one polar molecule is attracted to the delta- end of a neighbouring polar molecule. These attractions are in addition to London forces, not instead of them.

Worked example: identifying the strongest named force

Decide the strongest named intermolecular force in samples of CH4, HCl, and CO2.

The chemical interpretation is that the force depends on the whole molecule, not just the presence of one polar bond.

Guided practice

CH3Cl is polar, while CCl4 is non-polar overall. Complete the explanation:

CH3Cl has permanent dipole-dipole forces because ...

CCl4 has London forces as its strongest named force here because ...

Model completion: CH3Cl has an unsymmetrical arrangement of bond dipoles, so the molecule has a permanent dipole. CCl4 has polar C-Cl bonds arranged symmetrically in a tetrahedral molecule, so the bond dipoles cancel.

Hydrogen bonding is a particularly strong type of intermolecular attraction involving hydrogen. In Edexcel A level chemistry, predict hydrogen bonding when a hydrogen atom is covalently bonded to nitrogen, oxygen, or fluorine, and is attracted to a lone pair on nitrogen, oxygen, or fluorine in a neighbouring molecule.

The key pattern is:

X-H...Y

where X is usually N, O, or F, and Y is a lone-pair-bearing N, O, or F atom in another molecule. The ... represents the hydrogen bond; it is not a covalent bond.

Why those atoms? Nitrogen, oxygen, and fluorine are very electronegative and small. They make the H strongly delta+, and their lone pairs can attract that delta+ hydrogen closely.

The three specification examples

Hydrogen bonds are stronger than ordinary permanent dipole-dipole forces, but they are still intermolecular forces. Covalent bonds within molecules are much stronger.

Feynman diagnostic: fluorine is not enough

Explain why CH3F does not form hydrogen bonds between its own molecules, even though it contains fluorine. The diagnostic trap is to say "it has F, so it hydrogen bonds". The better answer is that CH3F has no hydrogen directly bonded to fluorine, oxygen, or nitrogen; its hydrogens are bonded to carbon, so they are not sufficiently delta+ for this hydrogen-bond pattern.

Use a two-question test when predicting hydrogen bonding between molecules of a pure substance.

1. Is there an H directly bonded to N, O, or F? This is the hydrogen-bond donor.
2. Is there a lone pair on N, O, or F in a neighbouring molecule? This is the hydrogen-bond acceptor.

For hydrogen bonding within a pure liquid, the molecules normally need both features available across the sample. A molecule with only a lone-pair acceptor may accept hydrogen bonds from water, but it may not hydrogen-bond to molecules of itself.

Worked example: applying the two-question test

Predict whether each pure substance can form hydrogen bonds between its molecules.

The mark-earning detail is "hydrogen directly bonded to N, O, or F". Saying "contains oxygen" is not enough, because an ether such as CH3OCH3 contains oxygen but has no O-H bond.

Water behaves unusually for a small simple molecular substance because each molecule can form an extended hydrogen-bonding network. The oxygen atom has two lone pairs and the molecule has two O-H bonds, so a water molecule can donate hydrogen bonds through its hydrogen atoms and accept hydrogen bonds through its oxygen lone pairs.

High melting and boiling temperatures

Water has relatively high melting and boiling temperatures because extra energy is needed to overcome many hydrogen bonds between water molecules. The explanation is not that water has strong covalent O-H bonds; those covalent bonds remain inside the molecules during melting and boiling. The relevant process is separation or increased movement of molecules, so the relevant attractions are intermolecular.

Pearson-style explanation chain:

1. Water molecules form hydrogen bonds with neighbouring water molecules.
2. Hydrogen bonds are stronger than ordinary permanent dipole-dipole forces and London forces between similar small molecules.
3. More energy is needed to overcome these intermolecular forces.
4. Therefore the melting and boiling temperatures are relatively high.

Density of ice compared with water

Ice is less dense than liquid water, so ice floats. In ice, hydrogen bonds hold water molecules in a relatively open arrangement. The molecules are kept further apart than in liquid water, where hydrogen bonds are constantly breaking and reforming and molecules can pack closer together.

Use the density idea explicitly: for the same mass of water molecules, the more open ice structure has a larger volume, so its density is lower because density = mass / volume.

Worked example: explaining an anomaly

Question: Explain why ice floats on liquid water.

Model answer:

Ice floats because it is less dense than liquid water. Hydrogen bonds hold water molecules in an open structure in ice, so the molecules are further apart than in liquid water. The same mass therefore occupies a larger volume, giving ice a lower density than liquid water.

For Edexcel 9CH0, intermolecular-force answers usually need a linked explanation, not just the name of a force. The safest structure is:

molecule feature -> force -> strength/energy -> property or prediction

Example 1: explain

Explain why NH3 can form hydrogen bonds.

Strong response: NH3 has N-H bonds, and nitrogen is highly electronegative, so the hydrogen atoms are delta+. The nitrogen atom has a lone pair. The delta+ hydrogen on one ammonia molecule can be attracted to the lone pair on nitrogen in a neighbouring ammonia molecule, forming a hydrogen bond.

Example 2: predict

Predict whether CH3OCH3 forms hydrogen bonds between molecules of the pure liquid.

Strong response: It does not form hydrogen bonds between its own molecules. It has oxygen lone pairs, but it has no hydrogen directly bonded to oxygen, nitrogen, or fluorine. Its hydrogens are bonded to carbon, so it lacks the hydrogen-bond donor needed for this intermolecular interaction.

Recap

This lesson's core skill is moving from structure to interaction to property. London forces come from instantaneous dipole-induced dipole interactions and occur in all molecules. Permanent dipole-dipole forces require polar molecules. Hydrogen bonding requires H directly bonded to N, O, or F and a lone pair on N, O, or F; this model explains the hydrogen bonding in H2O, liquid NH3, and liquid HF, plus water's high melting and boiling temperatures and the lower density of ice.