Ionic bonding and ionic compoundsEdexcel A-Level Chemistry: Revision notes

Ionic bonding is the strong electrostatic attraction between oppositely charged ions. It occurs between a metal, which forms positive ions, and a non-metal, which forms negative ions.

In a solid the ions pack into a giant ionic lattice: a regular, repeating three-dimensional arrangement in which every ion is surrounded by ions of opposite charge. The attraction acts in all directions, so there are no separate molecules. The formula of an ionic compound is the simplest ratio of ions, e.g. MgCl₂ is one Mg²⁺ to two Cl⁻.

Metal atoms lose their outer electrons to form cations (positive ions). Non-metal atoms gain electrons to form anions (negative ions). Most simple ions have the electronic configuration of a noble gas.

• Group 1: M → M⁺ + e⁻ (e.g. Na⁺)
• Group 2: M → M²⁺ + 2e⁻ (e.g. Mg²⁺)
• Aluminium: Al → Al³⁺ + 3e⁻
• Group 5: N + 3e⁻ → N³⁻
• Group 6: O + 2e⁻ → O²⁻
• Group 7: Cl + e⁻ → Cl⁻

The total positive charge in the lattice equals the total negative charge, which fixes the formula, e.g. Al³⁺ and O²⁻ give Al₂O₃.

A dot-and-cross diagram shows only the outer-shell electrons. The electrons from one atom are drawn as dots and from the other as crosses, so you can see where they came from. Each ion is placed in square brackets with its charge at the top right.

• Sodium chloride: [Na]⁺ is drawn with the 8 electrons of its new outer shell; [Cl]⁻ is drawn with 8 outer electrons, 7 of its own and 1 transferred from sodium.
• Magnesium oxide: [Mg]²⁺ with 8 outer electrons and [O]²⁻ with 8 outer electrons, two of which came from magnesium.
• Magnesium chloride: one [Mg]²⁺ and two separate [Cl]⁻ ions, each chloride with 8 outer electrons.

A cation that has lost its outer shell shows the full shell beneath it (Li⁺ shows 2 electrons).

The strength of the ionic bond depends on the electrostatic attraction between the ions, which increases with:

• greater ionic charge (a 2+/2− pair attracts far more strongly than a 1+/1− pair)
• smaller ionic radius, because the nuclei of the oppositely charged ions are closer together.

Worked comparison: magnesium oxide (Mg²⁺ and O²⁻) has a much higher melting temperature than sodium chloride (Na⁺ and Cl⁻) because the ions have double the charge and are smaller. Between sodium fluoride and sodium bromide the charges are the same, so the smaller F⁻ ion gives the stronger bond and the higher melting temperature.

Down a group the ions of a given charge get larger, because each ion has an extra electron shell (Li⁺ < Na⁺ < K⁺ and F⁻ < Cl⁻ < Br⁻).

Isoelectronic ions have the same number of electrons (and so the same electronic configuration). N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺ and Al³⁺ all have 10 electrons, 1s²2s²2p⁶. Their radius decreases from N³⁻ to Al³⁺ because the number of protons increases, so the nuclear charge is greater and pulls the same number of electrons more tightly.

A cation is smaller than its atom (it has lost its outer shell, and the same protons now attract fewer electrons). An anion is larger than its atom (extra electrons repel each other).

Because of the strong attractions in the lattice, ionic compounds:

• have high melting and boiling temperatures (much energy is needed to overcome the attractions)
• are hard but brittle: a force shifts a layer so that like ions line up and repel, splitting the crystal
• do not conduct when solid, because the ions are fixed in position
• conduct when molten or dissolved, because the ions are free to move and carry charge.

Evidence for ions: when a coloured ionic compound is placed on damp filter paper with a voltage applied, the colours move apart: blue copper(II) ions migrate towards the negative electrode and yellow chromate(VI) ions towards the positive electrode. Purple manganate(VII) ions also move towards the positive electrode. Charged particles moving to oppositely charged electrodes shows that ions exist.