Part 1: Ionic, covalent, dative & metallic bonding

1. Four types of bond

Every bond is an electrostatic attraction. What changes is which particles attract each other.

2. Ionic bonding and lattices

When a metal reacts with a non-metal, the metal atoms lose electrons and the non-metal atoms gain them. Take magnesium and oxygen. Magnesium (Z = 12, 2,8,2) loses its two outer electrons to become Mg²⁺ (2,8). Oxygen (Z = 8, 2,6) gains two electrons to become O²⁻ (2,8). Both ions now have 10 electrons, the same as neon.

In ionic notation, the ion is written in square brackets with its charge outside: [Mg]²⁺ and [O]²⁻, with the outer-shell electrons shown as dots or crosses to show where they came from.

Ionic compounds do not form molecules. The ions arrange themselves into a giant ionic lattice: a regular 3D arrangement where every ion is surrounded by ions of opposite charge. The formula (MgO, NaCl) gives the ratio of ions, not a molecule.

3. Ion charges and formulae

Simple ion charges follow from the group: Group 1 forms 1+, Group 2 forms 2+, Group 13 (Al) forms 3+, Group 15 forms 3−, Group 16 forms 2− and Group 17 forms 1−. You also need these compound ions:

Method: balance the charges

The total positive charge must equal the total negative charge. Use brackets around a compound ion if you need more than one of it.

4. Ionic bond strength

The attraction between ions is stronger when the ions have larger charges and smaller radii (so their centres can get closer). A stronger attraction needs more energy to overcome.

5. Ionic radii

Down a group, ions with the same charge get bigger, because each has more occupied shells (Li⁺ < Na⁺ < K⁺; F⁻ < Cl⁻ < Br⁻).

Isoelectronic ions have the same number of electrons. N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ all have 10 electrons (1s² 2s² 2p⁶). From N³⁻ (7 protons) to Al³⁺ (13 protons), the nuclear charge rises while the electron arrangement stays the same. The electrons are pulled in more strongly, so the radius decreases: N³⁻ is the largest and Al³⁺ the smallest.

Compared with their atoms: positive ions are smaller and negative ions are larger. For Na → Na⁺ or Mg → Mg²⁺, the whole outer shell is lost, so the ion has one fewer occupied shell. For an anion such as Cl⁻, the extra electron increases repulsion between electrons with the same nuclear charge, so the ion is larger than the atom. Not every cation loses an entire occupied shell; check the electron configuration.

6. Covalent bonding

A covalent bond is a shared pair of electrons between two atoms, usually non-metals. Each nucleus attracts the shared pair, and that attraction holds the atoms together.

7. Dative covalent bonding

In a dative bond, one atom supplies both electrons of the shared pair from a lone pair. The other atom (or ion) needs an empty orbital to accept them.

In the last example, boron in BF₃ has only six outer electrons, so it can accept nitrogen's lone pair. In a displayed formula, a dative bond is drawn as an arrow pointing from the donor atom (the one with the lone pair) to the acceptor: N→B, N→H⁺, O→H⁺.

8. Metallic bonding

A metal is a giant lattice of positive metal ions in a sea of delocalised electrons. The bond is the electrostatic attraction between the positive ions and the delocalised electrons. A complete answer about a metal's structure needs both: the giant lattice, and the attraction between the cations and the delocalised electrons.

Avoid a blanket rule such as "more protons always means a stronger metal". The comparison works through ion charge, number of delocalised electrons and ion size, not proton count alone.

9. Quick checks

Finesse practice: indicative answers to check your reasoning, not official mark allocations.

10. Sources