What you'll learn
Ionic bonding explains how metals combine with non-metals to form compounds such as sodium chloride and magnesium oxide. In this guide you will learn how ions form by the transfer of electrons, how to predict the charges of common ions, how to work out the formula of an ionic compound, how to represent ionic bonding with dot-and-cross diagrams, the structure of giant ionic lattices, and how that structure explains the characteristic properties of ionic compounds — high melting points and electrical conductivity when molten or dissolved. These ideas connect atomic structure to the behaviour of real materials.
Key terms and definitions
Ion — a charged particle formed when an atom (or group of atoms) gains or loses electrons.
Cation — a positively charged ion (formed by losing electrons), typically a metal.
Anion — a negatively charged ion (formed by gaining electrons), typically a non-metal.
Ionic bond — the strong electrostatic force of attraction between oppositely charged ions.
Giant ionic lattice — a regular, repeating 3D arrangement of alternating positive and negative ions.
Electrostatic attraction — the force between opposite charges holding the lattice together.
Core concepts
How ions form
When a metal reacts with a non-metal, electrons are transferred from the metal atom to the non-metal atom. Metals lose electrons to form positive ions; non-metals gain electrons to form negative ions. Each atom ends up with a full outer shell — a stable, noble-gas electronic structure. For example, sodium (2,8,1) loses one electron to become Na⁺ (2,8), and chlorine (2,8,7) gains one electron to become Cl⁻ (2,8,8).
Predicting ionic charge
For elements in Groups 1–3, the charge equals the group number (Group 1 = +1, Group 2 = +2, Group 3 = +3) because they lose their outer electrons. For Groups 6 and 7, the charge is negative: Group 7 = −1, Group 6 = −2, because they gain electrons to fill the outer shell. Some common compound ions also have fixed charges, such as sulfate (SO₄²⁻), carbonate (CO₃²⁻), nitrate (NO₃⁻) and hydroxide (OH⁻).
Working out formulae
In an ionic compound the total positive charge equals the total negative charge, so the compound is neutral overall. Balance the charges to find the ratio of ions. For sodium chloride, Na⁺ and Cl⁻ balance one-to-one: NaCl. For magnesium chloride, Mg²⁺ needs two Cl⁻: MgCl₂. For magnesium oxide, Mg²⁺ and O²⁻ balance one-to-one: MgO.
Dot-and-cross diagrams
A dot-and-cross diagram shows electron transfer using dots for one atom's electrons and crosses for the other's. Show the metal losing its outer electron(s) to the non-metal, then draw both ions in square brackets with their charges written outside. Often only the outer shells are shown for clarity.
Giant ionic lattices and properties
Ionic compounds form giant ionic lattices — regular 3D arrangements of alternating ions held by strong electrostatic forces in all directions. This structure explains their properties:
- High melting and boiling points, because a lot of energy is needed to overcome the many strong electrostatic attractions.
- Conduct electricity when molten or dissolved (but not when solid), because the ions are then free to move and carry charge; in the solid the ions are locked in place.
- Many are soluble in water.
Worked examples
Example 1: Forming an ion
Explain how a magnesium atom forms an ion.
Magnesium (2,8,2) loses its two outer electrons to gain a full outer shell, forming Mg²⁺ (2,8). The loss of two negative electrons leaves a 2+ charge.
Example 2: Working out a formula
What is the formula of calcium fluoride?
Calcium is Group 2, so Ca²⁺; fluorine is Group 7, so F⁻. To balance the 2+ charge, two F⁻ are needed: CaF₂.
Example 3: Explaining conductivity
Why does solid sodium chloride not conduct electricity, but molten sodium chloride does?
In the solid, the ions are held in fixed positions in the lattice and cannot move. When molten, the ions are free to move and carry charge, so it conducts.
Common mistakes and how to avoid them
Mixing up which atoms lose or gain electrons. Metals lose (forming positive ions); non-metals gain (forming negative ions).
Forgetting to balance charges in formulae. The overall charge must be zero — adjust the ratio of ions accordingly.
Saying ionic solids conduct electricity. They only conduct when molten or dissolved, when ions are free to move.
Drawing the wrong number of electrons after transfer. Check each ion has a full outer shell and the correct charge in brackets.
Confusing ionic and covalent bonding. Ionic = metal + non-metal with transfer; covalent = non-metals sharing.
Exam technique for Ionic Bonding
Describe electron transfer explicitly, naming which atom loses and which gains, and the resulting ions.
Use group number to predict charges quickly, and learn the common compound ions.
Balance charges to derive formulae, showing the ratio clearly.
Link lattice structure to properties — strong electrostatic forces explain high melting points; free-moving ions explain conductivity.
Draw neat dot-and-cross diagrams with charges and square brackets.
Quick revision summary
Ionic bonding occurs between metals and non-metals: electrons are transferred from the metal (forming positive cations) to the non-metal (forming negative anions), so each ion gains a stable full outer shell. Charges follow the group number — Group 1 = +1, Group 2 = +2, Group 6 = −2, Group 7 = −1 — and formulae are found by balancing total positive and negative charge to give a neutral compound (NaCl, MgCl₂, MgO). The ionic bond is the strong electrostatic attraction between oppositely charged ions, and ionic compounds form giant ionic lattices of alternating ions. This structure means high melting and boiling points (many strong forces to overcome) and electrical conductivity only when molten or dissolved, when ions are free to move — solids do not conduct because ions are fixed. Show electron transfer with dot-and-cross diagrams, predict charges from groups, balance formulae, and always link the giant lattice structure to the observed properties.