One goal, three strategies
Atoms bond to reach a full outer shell. There are only three ways to get there: give electrons away, take them, or share them. Which one happens is decided almost entirely by the electronegativity difference between the two atoms.
| Electronegativity difference | Bond type | Typical pair |
|---|---|---|
| Large (> ~1.8) | Ionic — full transfer | Metal + non-metal (NaCl) |
| Moderate (0.4 – 1.8) | Polar covalent — unequal sharing | H₂O, HCl |
| Small (< 0.4) | Pure covalent — equal sharing | H₂, Cl₂, CH₄ |
| Both low, both metals | Metallic — delocalised sea | Cu, Fe, Al |
Watch the ionic animation: the electron physically leaves sodium and joins chlorine. Then switch to Covalent — the pair stays between both nuclei, belonging to neither.
Why bond type dictates properties
This is the payoff. If you know the bonding, you can predict the substance's behaviour without ever having seen it.
Ionic solids form giant lattices of alternating charges. Breaking that lattice needs enormous energy, so melting points are high. They do not conduct when solid (ions are locked in place) but do conduct when molten or dissolved (ions are free to move). Brittle: shift the layers and like charges meet, and the crystal splits.
Simple covalent substances have strong bonds inside each molecule but weak forces between molecules. Melting only has to overcome the weak intermolecular forces, so melting points are low and they do not conduct — there are no free charges.
Metallic bonding gives free electrons: excellent electrical and thermal conduction, and layers that slide without breaking the bond, which is why metals are malleable rather than brittle.
The exception worth knowing
Giant covalent structures — diamond, graphite, silicon dioxide — are covalent but form one continuous network rather than separate molecules. Melting diamond means breaking actual covalent bonds, so it melts above 3500 °C.
Graphite is the interesting one: each carbon bonds to only three neighbours, leaving one electron delocalised. That is why graphite conducts electricity and diamond does not, despite both being pure carbon.
A common exam trap
Low melting point does not mean weak covalent bonds. The covalent bonds in a water molecule are very strong. What melts easily is the weak hydrogen bonding between molecules. Always be explicit about which force you are breaking.
Covalent bonding up close
A covalent bond is a shared pair of electrons held between two nuclei. Atoms can share one pair (a single bond, as in H–H), two pairs (a double bond, as in O=O) or three (a triple bond, as in N≡N) — the more pairs, the stronger and shorter the bond. Dot-and-cross diagrams track which atom each shared electron came from.
Sometimes both shared electrons come from the same atom: this is a dative (coordinate) bond, shown by an arrow. The ammonium ion, NH₄⁺, forms when the lone pair on ammonia's nitrogen bonds to an H⁺.
Shapes of molecules (A Level)
Electron pairs around a central atom repel each other and settle as far apart as possible — this is VSEPR (valence-shell electron-pair repulsion), and it sets a molecule's shape. Lone pairs repel slightly more strongly than bonding pairs, which squeezes bond angles a little.
| Bonding pairs | Lone pairs | Shape | Angle | Example |
|---|---|---|---|---|
| 2 | 0 | linear | 180° | CO₂ |
| 3 | 0 | trigonal planar | 120° | BF₃ |
| 4 | 0 | tetrahedral | 109.5° | CH₄ |
| 3 | 1 | trigonal pyramidal | 107° | NH₃ |
| 2 | 2 | bent | 104.5° | H₂O |
Intermolecular forces
The covalent bonds inside a molecule are strong, but the forces between molecules are much weaker — and it is those that must be overcome to melt or boil a simple molecular substance. From weakest to strongest they are: van der Waals (London) forces (present in all molecules, stronger the more electrons there are), permanent dipole–dipole forces (between polar molecules), and hydrogen bonding (when H is bonded to N, O or F).
This is why boiling points rise down a group of similar molecules, and why water boils far higher than its size suggests — its hydrogen bonds are unusually strong.
Which force am I breaking?
Melting or boiling a simple molecular substance breaks the weak intermolecular forces, NOT the covalent bonds inside the molecules — a classic exam trap. The molecules stay whole; they just separate.
A worked example
Predict and explain the shape and bond angle of a molecule of ammonia, NH₃.
- Nitrogen has 5 outer electrons; three form bonds to H and two remain as a lone pair.Count the electron pairs around the central atom: 3 bonding + 1 lone = 4 pairs.
- Four electron pairs repel to a tetrahedral arrangement.Four pairs get as far apart as possible — tetrahedral.
- One pair is a lone pair, which repels more strongly, so the shape of the atoms is trigonal pyramidal.We name the shape by the positions of the atoms, not the lone pair.
- The lone pair squeezes the bond angle from 109.5° down to about 107°.Lone-pair repulsion is greater than bonding-pair repulsion.
Trigonal pyramidal, bond angle ≈ 107°, because of one lone pair repelling the three bonding pairs.