Identify the three classes of intermolecular force (dispersion forces, dipole-dipole forces, hydrogen bonding) and use them to explain the physical properties of covalent molecular substances (melting and boiling points, solubility, viscosity, surface tension)

What this dot point is asking

QCAA wants you to identify the three types of intermolecular force (dispersion, dipole-dipole, hydrogen bonding), explain their origin and relative strength, and use them to predict and compare physical properties of covalent molecular substances (melting and boiling points, solubility, viscosity, surface tension). This dot point builds directly on the covalent bonding and polarity dot point.

The answer

Intermolecular forces are the attractions between separate molecules, not the bonds within them. They are significantly weaker than ionic, covalent or metallic bonds, but they govern the physical properties of all molecular substances. Three main types are required at QCE Chemistry Unit 1 level: dispersion forces, dipole-dipole attractions, and hydrogen bonding.

Three types of intermolecular force

Dispersion forces (also called London forces or instantaneous dipole-induced dipole forces). Present between all molecules, polar or non-polar. Caused by random fluctuations in electron density that create an instantaneous dipole; this induces a dipole in a neighbouring molecule; the two attract. Always present, but they are the only intermolecular force in non-polar substances.

Strength scales with:

• Number of electrons (larger, heavier molecules have more electrons and a more polarisable electron cloud, so stronger dispersion forces).
• Surface area / shape (long straight-chain molecules can stack closely, increasing dispersion; branched isomers cannot stack as well, so weaker dispersion).

Dipole-dipole attractions. Present in polar molecules. The delta+ end of one molecule attracts the delta- end of a neighbouring molecule. Stronger than dispersion forces of comparable mass, but typically weaker than hydrogen bonding.

Hydrogen bonding. A special, particularly strong dipole-dipole attraction that occurs when H is bonded directly to N, O or F (the three small, highly electronegative atoms). The H is so strongly delta+ that it attracts a lone pair on N, O or F of a neighbouring molecule. Roughly 5 to 10 times stronger than ordinary dipole-dipole attraction.

Strength order in a single molecule of comparable size: hydrogen bonding > dipole-dipole > dispersion. But for very large non-polar molecules (long alkanes), dispersion can outweigh dipole-dipole in a small polar molecule.

Hydrogen bonding in detail

Three conditions must be met:

1. A hydrogen atom covalently bonded to N, O or F (the donor).
2. A lone pair on N, O or F in a neighbouring molecule (the acceptor).
3. The donor and acceptor must be able to approach close enough (a few angstroms).

Common molecules with hydrogen bonding:

• Water (H_2O). Each O has 2 lone pairs and 2 O-H bonds, so it can participate in up to 4 hydrogen bonds. This is why water has anomalously high melting and boiling points compared with H_2S, H_2Se, H_2Te.
• Ammonia (NH_3). One lone pair on N, three N-H bonds. Average of 1 hydrogen bond per molecule.
• Hydrogen fluoride (HF). Three lone pairs on F, one H-F bond.
• Alcohols (R-OH).
• Amines (R-NH_2 or R_2-NH).
• Carboxylic acids (R-COOH). Strong hydrogen bonding, often dimeric.

C-H groups do not hydrogen bond in any meaningful sense at QCE Chemistry level: C is not electronegative enough to make the H delta+ enough.

Predicting physical properties

Melting point and boiling point. Higher with stronger intermolecular forces. The kinetic energy required to overcome the attractions is proportional to their strength. When comparing substances:

1. Start with the relative molecular mass (dispersion-force proxy).
2. Identify whether each substance is polar (dipole-dipole) and whether it has N-H, O-H or F-H (hydrogen bonding).
3. The substance with the strongest combined intermolecular forces has the highest melting/boiling point.

Trends in homologous series (e.g. alkanes): boiling point increases with chain length, because dispersion forces increase with electron count. Branched isomers boil lower than straight-chain ones of the same formula because they have less surface area for dispersion to act over.

Solubility. "Like dissolves like". Polar solvents (water, ethanol) dissolve polar and ionic solutes; non-polar solvents (hexane, oil) dissolve non-polar solutes. The principle is that the solute-solvent interactions must be of comparable strength to the solute-solute and solvent-solvent interactions.

• Water dissolves NaCl: water solvates the ions via ion-dipole interactions, comparable to the lattice energy lost.
• Water dissolves ethanol completely: O-H of ethanol hydrogen-bonds with water.
• Water does not dissolve hexane: water-water hydrogen bonds are too strong to give up for weak dispersion interactions with hexane.

The alkanols (CH_3OH, C_2H_5OH, C_3H_7OH, ...) become progressively less soluble in water as the carbon chain grows because the non-polar tail begins to dominate.

Viscosity. Higher with stronger intermolecular forces (molecules resist sliding past each other). Glycerol (three O-H groups, extensive hydrogen bonding) is much more viscous than water; water is much more viscous than hexane (only dispersion).

Surface tension. Higher with stronger intermolecular forces (the surface molecules are pulled inward more strongly). Water has high surface tension because of hydrogen bonding; hence water beads up on a non-polar surface.

Vapour pressure. Lower with stronger intermolecular forces. Substances with strong intermolecular forces have fewer molecules escaping to the vapour phase at any given temperature.

Worked comparison: boiling points of the period-2 hydrides

Notice the periodic-table dip down to CH_4 and the spike at H_2O. Water boils higher than HF and NH_3 because each H_2O can engage in 4 hydrogen bonds on average, while HF has 1 H-F donor and NH_3 has 3 donors but only 1 acceptor lone pair.

Worked comparison: structural isomers

Pentane (n-C_5H_12, straight chain) versus 2,2-dimethylpropane (neopentane, branched). Same molar mass (72 g/mol), no polarity (both alkanes). Pentane boils at 36 degrees C; neopentane at 10 degrees C. Difference is dispersion-force surface area: straight chain has more contact between neighbouring molecules; branched, near-spherical neopentane has less.

This shape-driven dispersion argument is a frequent QCAA EA prompt.

Common traps

Calling intermolecular forces "bonds". They are attractions between molecules, not bonds within molecules. Reserve "bond" for ionic, covalent and metallic interactions.

Forgetting dispersion in polar molecules. Polar molecules have dispersion forces too. The total intermolecular force is the sum of all applicable types.

Inventing hydrogen bonding without N, O or F. HCl, H_2S, CHCl_3 do not have hydrogen bonding. They are polar (dipole-dipole) but the H is not bonded to N, O or F.

Mixing up intramolecular and intermolecular. Boiling water at 100 degrees C breaks intermolecular hydrogen bonds, not the O-H covalent bonds (which require around 460 kJ/mol, hundreds of times more energy).

Treating dispersion as negligible. For large molecules, dispersion can dominate. Iodine (I_2) is solid at room temperature with only dispersion forces; bromine (Br_2) is liquid; chlorine (Cl_2) is gas. Mass and electron count matter.

Using "like dissolves like" without naming the forces. The principle is descriptive; the explanation is intermolecular forces. Always name which forces are involved.

In one sentence

Intermolecular forces (dispersion, present in all molecules and scaling with electron count and surface area; dipole-dipole, in polar molecules; and hydrogen bonding, the strongest, when H is bonded to N, O or F) govern the physical properties of covalent molecular substances, with stronger combined forces giving higher melting and boiling points, higher viscosity and surface tension, lower vapour pressure, and the "like-dissolves-like" pattern in solubility.