Intramolecular vs Intermolecular Forces

“Intra” means within; “inter” means between. That one difference in prefix separates two completely different kinds of force in chemistry. Intramolecular forces are the chemical bonds that hold atoms together within a molecule. Intermolecular forces are the much weaker attractions between separate molecules. Mixing them up is one of the most common mistakes in chemistry, and it leads to wrong explanations of melting, boiling and dissolving. This comparison sets out the differences clearly.

The difference in one picture

Imagine a room full of people holding hands in pairs.

• The grip within each pair is strong — hard to pull apart. That’s like an intramolecular force: a covalent bond holding a molecule together.
• Pairs also brush against and lean on neighbouring pairs, a much weaker contact. That’s like an intermolecular force.

If the room gets crowded or empties out (like a substance melting or boiling), the pairs move relative to each other — but each pair stays holding hands. Only something far more drastic (a chemical reaction) would split a pair.

Side by side

Feature	Intramolecular forces	Intermolecular forces
Where they act	Within a molecule (or throughout a lattice)	Between separate molecules
What they are	Chemical bonds	Attractions between molecules
Types	Covalent (incl. polar covalent and dative); in a broader sense ionic and metallic bonding	London (dispersion) forces, permanent dipole–dipole forces, hydrogen bonds (and ion–dipole in solutions)
Typical strength	About 150–1,000 kJ mol⁻¹	About 1–40 kJ mol⁻¹ for small molecules
Broken or overcome by	Chemical reactions	Physical changes: melting, boiling, dissolving, sublimation
Determine	Chemical properties: reactivity, bond energies, molecular shape	Physical properties: melting and boiling points, volatility, viscosity, surface tension, solubility
Changes identity of substance?	Yes — new substances form	No — same molecules, different state

Intramolecular forces: bonds

The main intramolecular force in molecules is the covalent bond: a shared pair of electrons between two atoms (see covalent bonding explained). Covalent bonds can be:

• non-polar (H–H, Cl–Cl),
• polar (H–Cl, O–H),
• single, double or triple,
• dative (both electrons from one atom).

Ionic and metallic bonding are also “intra” forces in the sense that they hold a substance’s particles together within a single continuous lattice — there are no separate molecules to be “between”. Breaking any of these bonds requires a lot of energy and changes the substance chemically.

Intermolecular forces: attractions between molecules

There are three main types between neutral molecules (see intermolecular forces):

1. London (dispersion) forces: between all molecules, from temporary and induced dipoles; stronger for bigger molecules (see London dispersion forces).
2. Permanent dipole–dipole forces: between polar molecules (see dipole–dipole forces).
3. Hydrogen bonds: a strong special case, between H bonded to N, O or F and a lone pair on N, O or F (see hydrogen bonding).

Despite the name, a hydrogen bond is usually an intermolecular force, not a chemical bond like a covalent bond. (Hydrogen bonds can also form between different parts of one large molecule, such as a protein — then they’re acting within the molecule, but they’re still the same weak type of attraction.)

Worked example 1: water

Water has both kinds of force:

• Intramolecular: two O–H covalent bonds in each molecule, each about 463 kJ mol⁻¹.
• Intermolecular: hydrogen bonds between molecules, about 20 kJ mol⁻¹ each, plus London and dipole–dipole forces.

What happens when water boils? Only intermolecular forces are overcome. Steam is still H₂O molecules. The energy needed — the enthalpy of vaporisation — is about 41 kJ mol⁻¹.

What would it take to break the O–H bonds? About 2 × 463 = 926 kJ per mole of water — more than twenty times the energy of boiling. This happens only in chemical reactions, such as electrolysis, which splits water into hydrogen and oxygen gas:

2H₂O(l) → 2H₂(g) + O₂(g)

That’s a chemical change; boiling is a physical change (see chemical vs physical changes).

Worked example 2: iodine

Iodine, I₂, is a shiny grey-black solid at room temperature. When warmed, it sublimes to a purple vapour.

• Intramolecular: one I–I covalent bond, about 151 kJ mol⁻¹.
• Intermolecular: London forces between I₂ molecules — quite strong for a non-polar molecule, because each molecule has 106 electrons.

Subliming iodine separates I₂ molecules from each other (enthalpy of sublimation about 62 kJ mol⁻¹); the purple vapour is still I₂ molecules. Only at much higher temperatures do I₂ molecules start to split into iodine atoms.

Worked example 3: comparing chlorine and sodium chloride

• Chlorine, Cl₂: strong Cl–Cl covalent bonds within molecules, weak London forces between them. It’s a gas (boils at −34 °C) because only weak intermolecular forces hold molecules together.
• Sodium chloride, NaCl: there are no molecules; strong ionic bonding holds the whole lattice together. Melting it means overcoming the ionic bonding itself, which is why it melts at 801 °C (see properties of ionic compounds).

The contrast shows why the question “which forces must be overcome?” is the key to explaining melting and boiling points.

Where the two meet: large molecules

In very large molecules, such as proteins, DNA and synthetic polymers, the line between the two kinds of force needs a little care. A protein chain is held together along its length by strong covalent peptide bonds — intramolecular in the strict sense. But the chain also folds back on itself, and different parts of the same molecule attract each other through hydrogen bonds, London forces and ionic attractions. These are the same weak kinds of attraction that act between separate small molecules; they just happen to act between distant parts of one long chain.

This is why gentle heating can denature a protein, changing its shape and destroying its function, without breaking a single covalent bond (see denaturation). It’s also why polymers soften on warming: the weak attractions between neighbouring chains loosen, so the chains can slide past each other, while the chains themselves remain intact. When you explain the properties of big molecules, always make clear whether you mean the strong bonds along the chain or the weak attractions between chains (or between parts of one chain).

A check you can use

When you’re explaining a property, ask:

• Is a new substance formed? If yes, bonds (intramolecular forces) are broken or formed.
• Is it just a change of state or dissolving? If yes, only intermolecular forces are involved (for simple molecular substances).

Common mistakes

• “When ice melts, the bonds break.” Hydrogen bonds between molecules break; the O–H covalent bonds don’t.
• “Methane boils easily because its C–H bonds are weak.” C–H bonds are strong (413 kJ mol⁻¹); methane boils easily because London forces between its molecules are weak.
• “Hydrogen bonds are covalent bonds to hydrogen.” They’re attractions between molecules.
• “Stronger bonds mean higher boiling point.” For simple molecular substances, boiling point depends on intermolecular forces, not the bonds within molecules.

Key takeaways

• Intramolecular forces are bonds within molecules (covalent, and ionic/metallic in lattices) — strong, broken only in chemical reactions.
• Intermolecular forces are attractions between molecules (London, dipole–dipole, hydrogen bonds) — weak, overcome in physical changes.
• Intramolecular forces set chemical properties; intermolecular forces set physical properties like boiling point and solubility.
• Boiling water overcomes hydrogen bonds (about 41 kJ mol⁻¹), not O–H bonds (about 926 kJ per mole of water).