Atoms, Elements and Compounds – Ions and Ionic Bonds
Learning objective
State that an ionic bond is a strong electrostatic attraction between oppositely‑charged ions.
1. What are ions?
- Ions are atoms that have gained or lost electrons and therefore carry an electrical charge.
- Cations – positively‑charged ions formed when an atom loses one or more electrons (usually a metal).
Example: Na → Na⁺ + e⁻
- Anions – negatively‑charged ions formed when an atom gains one or more electrons (usually a non‑metal).
Example: Cl + e⁻ → Cl⁻
2. Dot‑and‑cross diagrams (Core requirement 3)
Electrons belonging to the atom that loses them are shown as dots (·). Electrons that are taken up are shown as crosses (×). The diagrams illustrate the transfer of electrons from a Group I element to a Group VII element.
| Step |
Diagram |
Result |
| 1 |
Na· → Na⁺ + · |
Sodium becomes Na⁺ (cation) |
| 2 |
Cl + · → Cl× |
Chlorine becomes Cl⁻ (anion) |
| 3 |
Na⁺ Cl⁻ → Na⁺ Cl⁻ |
Electrostatic attraction forms an ionic bond |
3. How an ionic bond is formed
- Electron transfer: a metal atom transfers one or more electrons to a non‑metal atom.
- Ion formation: the metal becomes a cation, the non‑metal becomes an anion.
- Electrostatic attraction: opposite charges attract strongly, producing a stable ionic compound.
4. Giant lattice structure
- Ionic compounds do not consist of discrete molecules; they form a three‑dimensional giant lattice in which each cation is surrounded by several anions and each anion by several cations.
- The lattice maximises attractive forces and minimises repulsive forces, giving ionic solids high melting/boiling points and hardness.
- Because the lattice is continuous, ionic solids are brittle – when a layer slips, like‑charged ions are forced together and repel, causing fracture.
5. Why ionic bonds are strong
- Charge magnitude: larger absolute charges give a stronger attraction (e.g., Mg²⁺ + O²⁻ > Na⁺ + Cl⁻).
- Ionic radii: smaller ions can approach each other more closely, increasing the force.
- Lattice energy: proportional to the product of the charges and inversely proportional to the sum of the ionic radii. Higher lattice energy → higher melting/boiling points and greater hardness.
6. Example – Sodium chloride (NaCl)
| Step |
Process |
Species formed |
| 1 |
Sodium loses one electron (Na· → Na⁺ + ·) |
Na⁺ |
| 2 |
Chlorine gains one electron (Cl + · → Cl×) |
Cl⁻ |
| 3 |
Oppositely charged ions attract → Na⁺ Cl⁻ lattice |
Solid NaCl |
7. Example – Magnesium oxide (MgO) – illustrating the effect of charge
| Step |
Process |
Species formed |
| 1 |
Magnesium loses two electrons (Mg·· → Mg²⁺ + 2·) |
Mg²⁺ |
| 2 |
Oxygen gains two electrons (O + 2· → O××) |
O²⁻ |
| 3 |
Strong attraction between Mg²⁺ and O²⁻ → MgO lattice (higher lattice energy than NaCl) |
Solid MgO |
8. Properties of ionic compounds (derived from the giant lattice)
- Very high melting and boiling points.
- Usually soluble in water; polar water molecules separate the ions.
- Conduct electricity when molten or in aqueous solution (ions are free to move).
- Brittle – slipping of ion layers brings like charges together, causing fracture.
9. Summary
An ionic bond is a strong electrostatic attraction between oppositely‑charged ions that results from the transfer of electrons from a metal (Group I) to a non‑metal (Group VII). The resulting cations and anions arrange themselves in a three‑dimensional giant lattice, giving ionic compounds their characteristic high melting/boiling points, electrical conductivity in the molten or aqueous state, solubility in polar solvents, and brittleness. The bond strength increases with greater ionic charge and smaller ionic radii, as reflected in lattice energy.