Chemical bonding

• Ionic — electrons transferred from metal to non-metal, giving oppositely charged ions held in a giant lattice by electrostatic attraction. High melting point; conducts only when molten or dissolved, because only then are the ions free to move.
• Covalent — a shared pair of electrons between two non-metal atoms. A dative (coordinate) bond is a covalent bond where both electrons came from the same atom — shown with an arrow, as in $\mathrm{N}{\mathrm{H}}_4^{+}$ and ${\mathrm{H}}_3{\mathrm{O}}^{+}$ — and once formed it is identical to any other covalent bond.
• Metallic — positive ions in a sea of delocalised electrons. Explains conductivity in the solid, malleability (layers slide without breaking the bonding) and high melting points.

Bond strength rises with charge and falls with size: the lattice energy of $\mathrm{M}\mathrm{g}\mathrm{O}$ far exceeds $\mathrm{N}\mathrm{a}\mathrm{C}\mathrm{l}$ because 2+ and 2− ions attract far more strongly than 1+ and 1−, and smaller ions sit closer together.

The method, every time: count the electron pairs around the central atom (bonding pairs plus lone pairs), arrange them to be as far apart as possible, then name the shape by where the atoms are — not where the lone pairs are.

Look at the middle three rows together. Methane, ammonia and water all have four electron pairs, so all three are based on a tetrahedron — but methane has no lone pairs (109.5°), ammonia has one (107°) and water has two (104.5°). Each lone pair squeezes the angle by roughly 2.5°, because a lone pair is held by only one nucleus and so spreads out more and repels harder. The repulsion order is:

For counting purposes a double or triple bond behaves as one region of electron density, which is why $\mathrm{C}{\mathrm{O}}_2$ — two double bonds, no lone pairs on the carbon — is linear.

Electronegativity is an atom’s ability to attract the electron pair in a covalent bond. It increases across a period (rising nuclear charge, same shell) and decreases down a group (outer shell further away, more shielding), making fluorine the most electronegative element.

A difference in electronegativity makes a bond polar, with partial charges $\delta +$ and $\delta -$. But a molecule with polar bonds is not necessarily polar: if the shape is symmetrical the bond dipoles cancel. Carbon dioxide has two strongly polar C=O bonds and no overall dipole because it is linear; water has the same kind of polar bonds but is bent, so the dipoles add. $\mathrm{C}\mathrm{C}{\mathrm{l}}_4$ is non-polar for the same reason $\mathrm{C}\mathrm{H}\mathrm{C}{\mathrm{l}}_3$ is polar — symmetry, not the bonds.

These act between molecules and are far weaker than the covalent bonds within them. Three, in increasing strength:

• Induced dipole–induced dipole (van der Waals / London) forces. Present between all molecules. They arise from the momentary uneven distribution of electrons, which induces a dipole in a neighbour. They get stronger with more electrons and with a larger surface area of contact — which is why boiling point rises down a homologous series, and why a straight-chain alkane boils higher than its branched isomer.
• Permanent dipole–permanent dipole forces. Between polar molecules, in addition to the above.
• Hydrogen bonding. The strongest, and it needs a specific arrangement: hydrogen bonded directly to N, O or F, and a lone pair on the N, O or F of a neighbouring molecule for it to point at.

Hydrogen bonding explains the anomalies that questions love: water boils far above ${\mathrm{H}}_2\mathrm{S}$ despite being lighter; ice is less dense than liquid water because hydrogen bonds hold the molecules in an open lattice; and ethanol boils far above its isomer methoxymethane, which has the same formula but no O–H bond.

The structure that earns full marks, in order:

1. Name the strongest intermolecular force in each substance.
2. Say which is stronger and why — more electrons, hydrogen bonding present, greater surface contact.
3. Say that more energy is therefore needed to overcome it, so the boiling point is higher.

The mark-losing phrase is “the bonds are stronger”. Boiling a molecular substance breaks intermolecular forces, not covalent bonds — the molecules survive the process intact. Say so explicitly; examiner reports single this out.

Giant structures are a separate case and are compared differently: melting $\mathrm{S}\mathrm{i}{\mathrm{O}}_2$ or diamond means breaking strong covalent bonds throughout a lattice, and melting an ionic solid means overcoming electrostatic attraction between ions — both far higher than any molecular substance.

1. 1.Saying boiling breaks covalent bonds

   It breaks intermolecular forces. The molecules themselves stay whole. This is the most-reported error in the topic.

2. 2.Naming the shape by the electron pairs

   Count all pairs to get the arrangement, then name the shape by where the atoms are. $\mathrm{N}{\mathrm{H}}_3$ has a tetrahedral arrangement but a pyramidal shape.

3. 3.Polar bonds assumed to mean a polar molecule

   Symmetry can cancel the dipoles. $\mathrm{C}{\mathrm{O}}_2$ and $\mathrm{C}\mathrm{C}{\mathrm{l}}_4$ have very polar bonds and no overall dipole.

4. 4.Hydrogen bonding claimed without N, O or F

   Hydrogen must be bonded directly to N, O or F, and there must be a lone pair on a neighbour. HCl has a permanent dipole, not a hydrogen bond.

5. 5.Forgetting the lone pair on the central atom

   Count the central atom’s outer electrons, subtract those used in bonding, and halve the rest. Missing a lone pair gives the wrong shape and the wrong angle.