IMAT Chemistry · Chemical Bonding

Intermolecular Forces, Boiling Points and Solubility

Weak attractions between molecules (London forces, dipole attractions and hydrogen bonds) decide boiling points and what dissolves in what.

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Why this matters

Asked in 2015, 2016, every year from 2019 to 2022, and again in 2026: which substance has hydrogen bonds, which has the weakest forces, and why one substance boils higher than another. The 2026 question tested the difference between forces between molecules and bonds inside them.

Concept 1 of 4: The three intermolecular forces and where each acts

Molecules attract each other weakly even when they have no overall charge. Electrons are always moving, so at any instant a molecule can be slightly lopsided and make its neighbour lopsided too. Polar molecules add a permanent attraction between their ends, and molecules with H on N, O or F add an especially strong one.

Definition

Intermolecular forces act between separate molecules. They are much weaker than the covalent bonds inside a molecule (intramolecular forces).

London (dispersion) forces, also called instantaneous or induced dipole forces, act between all molecules and atoms. They grow with the number of electrons (bigger molecules) and with contact area (long, unbranched chains).
Dipole-dipole forces act between polar molecules: the $\delta^+$ end of one attracts the $\delta^-$ end of the next.
Hydrogen bonds are the strongest of the three, roughly a tenth of a covalent bond.
Van der Waals forces is the collective name for London and dipole-dipole forces.

Force	Acts between	Cause	Relative strength
London (dispersion)	All molecules and noble gas atoms	Temporary dipole inducing a dipole in a neighbour	Weakest for small molecules; grows with size
Dipole-dipole	Polar molecules	Permanent $\delta^+$ and $\delta^-$ ends attracting	Stronger than London for molecules of similar size
Hydrogen bond	Molecules with H bonded to N, O or F	Very $\delta^+$ H attracted to a lone pair on N, O or F	Strongest intermolecular force
Ion-dipole	An ion and a polar molecule	Ion charge attracting the opposite end of the dipole	Strong; explains salts dissolving in water

Practice this conceptself-check · 3 quick reps

London forces act in polar molecules too

London forces are present in every molecule. In large molecules they can outweigh dipole forces: HI boils higher than HCl even though HCl is more polar, because HI has far more electrons. When asked why a heavier hydrogen halide boils higher, the answer is London forces, not dipoles.

Concept 2 of 4: Hydrogen bonding: which substances have it

N, O and F are small and very electronegative, so an H atom bonded to one of them is left almost bare of electrons. That exposed H is strongly attracted to a lone pair on an N, O or F atom of a neighbouring molecule. Without the N, O or F to strip the H, the effect does not happen.

Definition

A hydrogen bond forms between an H atom covalently bonded to N, O or F and a lone pair on an N, O or F atom of another molecule.

Both parts are needed between molecules of a pure substance: an N to H, O to H or F to H bond, and a lone pair on N, O or F.
H bonded to C (in $\mathrm{CH_3F}$ , ethers, aldehydes, ketones) cannot form hydrogen bonds, even when the molecule contains O or F. These molecules can still accept hydrogen bonds from water.
Water forms up to four hydrogen bonds per molecule, which explains its high boiling point and why ice is less dense than liquid water.
Carboxylic acids (ethanoic acid) hydrogen bond very strongly, pairing up as dimers.
Hydrogen bonds hold together the two strands of DNA and shape proteins.

Substance	H bonded to N, O or F?	Hydrogen bonds between its molecules?
Water, $\mathrm{H_2O}$	Yes, O to H	Yes
Ammonia, $\mathrm{NH_3}$	Yes, N to H	Yes
Hydrogen fluoride, HF	Yes, F to H	Yes
Ethanol, $\mathrm{C_2H_5OH}$	Yes, O to H	Yes
Methylamine, $\mathrm{CH_3NH_2}$	Yes, N to H	Yes
Methane, $\mathrm{CH_4}$	No	No
Hydrogen sulfide, $\mathrm{H_2S}$	No, H on S	No
Dimethyl ether, $\mathrm{CH_3OCH_3}$	No, every H is on C	No
Propanone, $\mathrm{CH_3COCH_3}$	No, every H is on C	No

Practice this conceptself-check · 3 quick reps

Having O or F in the molecule is not enough

A hydrogen bond needs H attached directly to N, O or F. Fluoromethane, ethers and aldehydes contain O or F but all their H atoms are on carbon, so their pure liquids have no hydrogen bonds.

Concept 3 of 4: Ranking boiling points from intermolecular forces

Boiling a molecular substance pulls whole molecules away from each other into the gas. The bonds inside each molecule stay intact. So the boiling point measures only how strongly the molecules attract each other: the stronger those forces, the more energy is needed and the higher the boiling point.

Definition

For simple molecular substances, higher boiling point = stronger intermolecular forces. Boiling breaks no covalent bonds. Method, in order:

1. Giant structures (ionic, metallic, covalent network) boil far higher than any simple molecule.
2. Among molecules of similar size: hydrogen bonding > dipole-dipole > London only.
3. Same kind of forces: more electrons (bigger $M_r$ ) means a higher boiling point.
4. Same formula: the long unbranched chain boils higher than the compact branched isomer (more contact area).
The same reasoning orders melting points and volatility: weaker forces, more volatile.

Worked example

Butane (

M_r = 58

), propanone (58) and propan-1-ol (60) have similar masses. Put them in order of increasing boiling point and explain.

Practice this conceptself-check · 4 quick reps

Boiling breaks intermolecular forces, not covalent bonds

Steam is made of whole

\mathrm{H_2O}

molecules. A stronger O to H or C to H bond inside a molecule does not raise the boiling point. Options that explain a boiling point with the strength of the bonds inside the molecule (intramolecular forces) are wrong.

Isomers with the same mass can boil at different temperatures

Equal

M_r

does not mean equal boiling points. A branched isomer is more compact, touches its neighbours less, and has weaker London forces, so it boils lower than the straight-chain isomer.

Concept 4 of 4: Solubility: like dissolves like

To dissolve, the solute particles must be pulled apart and surrounded by solvent particles. This happens easily when the new attractions between solute and solvent are as strong as the old ones. Polar and ionic solutes find strong partners in water; non-polar solutes only find London forces there, which cannot make up for breaking water's hydrogen bonds.

Definition

Like dissolves like: polar and ionic substances dissolve in polar solvents such as water; non-polar substances dissolve in non-polar solvents such as hexane.

Ionic solids dissolve in water because the water dipoles surround each ion (ion-dipole attraction, hydration).
Small molecules that hydrogen bond (methanol, ethanol, glucose, ammonia) dissolve well in water.
In a series of alcohols, solubility in water falls as the carbon chain grows: the non-polar chain takes over from the single O to H group.
Fats, oils and $\mathrm{I_2}$ dissolve in non-polar solvents, poorly in water.

Substance	Dissolves well in	Reason
Sodium chloride	Water, not hexane	Ion-dipole attraction to water molecules
Ethanol	Water, in any proportion	Hydrogen bonds with water
Glucose	Water	Several O to H groups hydrogen bond with water
Iodine, $\mathrm{I_2}$	Hexane, only slightly in water	Non-polar, London forces only
Cooking oil	Hexane, not water	Long non-polar chains
Hexan-1-ol	Hexane better than water	The long chain outweighs one O to H group

Practice this conceptself-check · 3 quick reps

A molecule with polar bonds may still be non-polar and insoluble in water

\mathrm{CCl_4}

has four polar C to Cl bonds, but its symmetric tetrahedral shape cancels the dipoles, so it hardly dissolves in water. Solubility follows the polarity of the whole molecule, not of its bonds.

Summary — formulas & gotchas at a glance

A revision cheat-sheet for the formulas and gotchas above. Click any concept name to jump back to its full explanation.

Reference tables (3)

The three intermolecular forces and where each acts4 rows

Force	Acts between	Cause	Relative strength
London (dispersion)	All molecules and noble gas atoms	Temporary dipole inducing a dipole in a neighbour	Weakest for small molecules; grows with size
Dipole-dipole	Polar molecules	Permanent $\delta^+$ and $\delta^-$ ends attracting	Stronger than London for molecules of similar size
Hydrogen bond	Molecules with H bonded to N, O or F	Very $\delta^+$ H attracted to a lone pair on N, O or F	Strongest intermolecular force
Ion-dipole	An ion and a polar molecule	Ion charge attracting the opposite end of the dipole	Strong; explains salts dissolving in water

Hydrogen bonding: which substances have it9 rows

Substance	H bonded to N, O or F?	Hydrogen bonds between its molecules?
Water, $\mathrm{H_2O}$	Yes, O to H	Yes
Ammonia, $\mathrm{NH_3}$	Yes, N to H	Yes
Hydrogen fluoride, HF	Yes, F to H	Yes
Ethanol, $\mathrm{C_2H_5OH}$	Yes, O to H	Yes
Methylamine, $\mathrm{CH_3NH_2}$	Yes, N to H	Yes
Methane, $\mathrm{CH_4}$	No	No
Hydrogen sulfide, $\mathrm{H_2S}$	No, H on S	No
Dimethyl ether, $\mathrm{CH_3OCH_3}$	No, every H is on C	No
Propanone, $\mathrm{CH_3COCH_3}$	No, every H is on C	No

Solubility: like dissolves like6 rows

Substance	Dissolves well in	Reason
Sodium chloride	Water, not hexane	Ion-dipole attraction to water molecules
Ethanol	Water, in any proportion	Hydrogen bonds with water
Glucose	Water	Several O to H groups hydrogen bond with water
Iodine, $\mathrm{I_2}$	Hexane, only slightly in water	Non-polar, London forces only
Cooking oil	Hexane, not water	Long non-polar chains
Hexan-1-ol	Hexane better than water	The long chain outweighs one O to H group

Watch out for (5)

London forces act in polar molecules too
→ The three intermolecular forces and where each acts
Having O or F in the molecule is not enough
→ Hydrogen bonding: which substances have it
Boiling breaks intermolecular forces, not covalent bonds
→ Ranking boiling points from intermolecular forces
Isomers with the same mass can boil at different temperatures
→ Ranking boiling points from intermolecular forces
A molecule with polar bonds may still be non-polar and insoluble in water
→ Solubility: like dissolves like

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