3.1.3.7 - Forces Between Molecules

Why these forces matter

Simple molecular substances often have relatively low melting and boiling points compared with ionic compounds or giant covalent structures. That is not because covalent bonds are weak. It is because melting and boiling only require the intermolecular forces between molecules to be overcome, and those forces are much weaker than the covalent bonds within each molecule.

For a useful comparison, identify all the attractions present and then compare similar molecules. London forces occur between all molecules; permanent dipole–dipole attractions occur between polar molecules; hydrogen bonding requires particular donor and lone-pair sites. Hydrogen bonds are relatively strong individual intermolecular attractions, but there is no universal ranking of whole substances by the name of one force. A large, easily polarised non-polar molecule can have stronger overall attractions than a small polar molecule. Crystal packing also affects melting points, so boiling-point comparisons are often simpler.

The diagram below shows the key distinction: during melting or boiling, molecules separate from one another, but the covalent $\ce{O-H}$ bonds inside each water molecule stay intact.

[DIAGRAM: asset_name: 1.3.7 - Forces Between Molecules - Diagram 1; asset_slug: 1.3.7 - Forces Between Molecules - Diagram 1; recommended_method: retained_png; description: Strong O–H covalent bonds within water molecules compared with a weaker hydrogen bond between molecules.]

Induced dipole-dipole forces

Start with the force that every atom and every molecule has. Electrons are always moving, so the electron cloud around a particle is never perfectly fixed. At any instant, the electrons may be slightly more concentrated on one side. That creates a temporary dipole: one side becomes slightly negative and the other slightly positive.

This temporary dipole can distort the electron cloud in a neighboring particle and induce a dipole there as well. The two particles then attract each other. Because the dipoles are constantly appearing and disappearing, these attractions are weak, but they are present all the time. The diagram below shows how a temporary dipole in one iodine molecule can induce a dipole in a neighboring molecule.

[DIAGRAM: asset_name: 1.3.7 - Forces Between Molecules - Diagram 2; asset_slug: 1.3.7 - Forces Between Molecules - Diagram 2; recommended_method: retained_png; description: A temporary dipole in iodine induces a dipole in a neighbouring molecule, producing attraction between opposite partial charges.]

Induced dipole-dipole forces become stronger when electron clouds are larger and more easily distorted. That is why molecules with more electrons usually have stronger London forces and therefore higher melting and boiling points. Shape can matter too: long, less-branched molecules can make contact over a larger surface, so the attractions between them are stronger than in compact, highly branched molecules with a similar $M_r$.

This same pattern explains why boiling points increase down the noble gases and why longer-chain hydrocarbons generally boil at higher temperatures than shorter ones.

Permanent dipole-dipole forces

Some molecules have a permanent uneven distribution of charge. This happens when the molecule is polar overall, not just when it contains a polar bond. That distinction matters. A carbon dioxide molecule contains two polar $\ce{C=O}$ bonds, but the molecule is linear, so the bond dipoles cancel and there is no permanent dipole. Hydrogen chloride, $\ce{HCl}$, has one polar bond and is polar overall, so it does have a permanent dipole.

When polar molecules come close together, the $\delta^+$ region of one molecule attracts the $\delta^-$ region of another. This is a permanent dipole-dipole force. These add to London forces and can raise the boiling point relative to an otherwise comparable non-polar molecule. The total attraction still depends on electron-cloud size and molecular shape.

Remember, however, that polar molecules do not lose their London forces. $\ce{HCl}$ molecules experience both induced dipole-dipole forces and permanent dipole-dipole forces. A non-polar molecule such as $\ce{CO2}$ only has induced dipole-dipole forces. The diagram below compares these two situations directly: in $\ce{CO2}$ the bond dipoles cancel, whereas $\ce{HCl}$ molecules can align so opposite partial charges attract.

[DIAGRAM: asset_name: 1.3.7 - Forces Between Molecules - Diagram 3; asset_slug: 1.3.7 - Forces Between Molecules - Diagram 3; recommended_method: retained_png; description: CO2 bond dipoles cancel while HCl molecules can attract through their permanent dipoles.]

A simple experimental observation is the deflection of a thin stream of water by a charged rod. Polar molecules turn so their nearer ends are attracted more strongly, bending the stream towards the rod. Use a narrow stream, keep the rod out of the liquid and compare liquids under the same conditions; flammable liquids must be kept away from ignition sources. Deflection is evidence about electrical response, not a direct measurement of hydrogen-bond strength.

For an unfamiliar molecule, the crucial question is whether any bond dipoles add up to a polar molecule overall. If the molecule is not polar overall, permanent dipole-dipole forces cannot act between its molecules.

A burette can provide the narrow liquid jet: clamp it securely, collect the liquid below and adjust its tap to give the same flow for each comparison. Keep the rod at the same distance from the stream.

Hydrogen bonding and water's unusual behavior

Hydrogen bonding is a particularly strong type of intermolecular attraction, present under specific conditions. The hydrogen atom must be covalently bonded to a very electronegative atom, and there must be a lone pair on a nitrogen, oxygen, or fluorine atom in a neighboring molecule.

Why is this attraction especially strong? Nitrogen, oxygen, and fluorine are very electronegative, so they pull electron density away from hydrogen very strongly. That leaves the hydrogen atom with a large partial positive charge concentrated in a very small space. A lone pair on a neighboring N, O, or F atom is then strongly attracted to it. The diagram below shows this in water, where a $\delta^+$ hydrogen on one molecule is attracted to a lone pair on the oxygen of another molecule.

[DIAGRAM: asset_name: 1.3.7 - Forces Between Molecules - Diagram 4; asset_slug: 1.3.7 - Forces Between Molecules - Diagram 4; recommended_method: retained_png; description: A water hydrogen bond links a partially positive hydrogen to a lone pair on a neighbouring oxygen.]

This is why $\ce{NH3}$, $\ce{H2O}$, and $\ce{HF}$ have anomalously high boiling points. Their boiling points are higher than predicted by extrapolating the trend of the heavier hydrides in each group. This does not mean ammonia must boil above every heavier Group 15 hydride: a sufficiently large electron cloud can give substantial London attractions. The anomaly is the departure from the expected trend.

Hydrogen bonding also explains why ice is less dense than liquid water. In liquid water, hydrogen bonds are constantly breaking and reforming, so the molecules can move closer together. In ice, the hydrogen bonds hold water molecules in a more open, regular arrangement. The molecules are therefore slightly further apart, so the same mass occupies a larger volume. A larger volume for the same mass means a lower density, which is why ice floats. The diagram below compares the open hydrogen-bonded arrangement in ice with the closer packing in liquid water.

[DIAGRAM: asset_name: 1.3.7 - Forces Between Molecules - Diagram 5; asset_slug: 1.3.7 - Forces Between Molecules - Diagram 5; recommended_method: deterministic_code; description: Equal numbers of water molecules occupy a more open arrangement in ice than in liquid water; selected hydrogen bonds are dashed.]

A sensible order is:

1. Check whether hydrogen is bonded to N, O, or F and whether a neighboring molecule can provide a lone pair on N, O, or F. If yes, hydrogen bonding is present.
2. If not, decide whether the molecule is polar overall. If yes, permanent dipole-dipole forces are present.
3. In every case, include induced dipole-dipole forces, because all atoms and molecules have them.

Once you know which forces are present, link force strength to property: stronger intermolecular forces mean more energy is needed for melting or boiling, so melting points and boiling points are higher.