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Chemical Bonds » Ionic Bonding

What you'll learn this session

Study time: 30 minutes

AQA spec: 4.2.1.1, 4.2.1.2

The three types of strong chemical bond and which particles each one involves
How metal atoms pass electrons to non-metal atoms to make ions
How to draw dot and cross diagrams for ionic compounds
How to work out the charge on an ion from its group number

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Three types of strong chemical bond

Table salt is held together by ionic bonds - strong attractions between oppositely charged Na⁺ and Cl⁻ ions

Atoms in substances are held together by chemical bonds. There are three types of strong chemical bond: ionic, covalent and metallic. Each one involves different particles.

Type of bond	The particles	Where it is found
Ionic	Oppositely charged ions	Compounds formed from metals combined with non-metals
Covalent	Atoms which share pairs of electrons	Most non-metallic elements and compounds of non-metals
Metallic	Atoms which share delocalised electrons	Metallic elements and alloys

So you can often tell the type of bonding just by looking at the elements in a substance:

Lithium iodide is a metal plus a non-metal, so it has ionic bonding.
Carbon dioxide is made only of non-metals, so it has covalent bonding.
Iron is a metal element, so it has metallic bonding.

All three bonds come down to two ideas: electrons are either transferred (ionic) or shared (covalent and metallic), and the bond is held by electrostatic forces. You will meet covalent bonding and metallic bonding in their own lessons. This lesson is all about ionic bonding.

Key terms:

Chemical bond: a strong force that holds particles together in a substance.
Electrostatic force: the force of attraction between opposite charges (positive and negative).
Ionic bonding: the strong electrostatic attraction between oppositely charged ions.
Electron transfer: electrons moving from one atom to another during a reaction.

How ionic bonds form

When sodium meets chlorine, each sodium atom hands its one outer electron to a chlorine atom, making Na⁺ and Cl⁻ ions

When a metal atom reacts with a non-metal atom, the electrons in the outer shell of the metal atom are transferred to the non-metal atom.

⊕ The metal atom

It loses its outer electrons. It now has more protons than electrons, so it becomes a positively charged ion.

⊖ The non-metal atom

It gains those electrons. It now has more electrons than protons, so it becomes a negatively charged ion.

The ions made by metals in Groups 1 and 2, and by non-metals in Groups 6 and 7, end up with the electronic structure of a noble gas (Group 0). Each ion has a full outer shell.

The positive and negative ions are then strongly attracted to each other by electrostatic forces. That attraction is the ionic bond. How millions of these ions pack together into a structure is covered in Ionic Compounds and Their Structure.

Dot and cross diagrams

We show electron transfer using a dot and cross diagram. The electrons from one atom are drawn as dots and the electrons from the other atom are drawn as crosses. That way you can see exactly where each electron came from and where it went.

The spec example is sodium chloride:

Draw a sodium atom (2,8,1) with its outer electron as a cross, and a chlorine atom (2,8,7) with its outer electrons as dots.
Draw an arrow to show the one outer electron moving from sodium to chlorine.
Draw the ions. Sodium is now 2,8. Chlorine is now 2,8,8, with seven dots and one cross in its outer shell.
Put each ion in square brackets with its charge at the top right: [Na]⁺ and [Cl]⁻.

In exams you can usually show just the outer shell of each atom. Once sodium loses its electron, its outer shell is the full second shell of 8 electrons. Some diagrams draw this full shell for the ion, and some show the ion with no outer electrons. Both are fine as long as the brackets and charges are right.

Worked example 1: magnesium oxide

Magnesium is in Group 2 (2,8,2). Oxygen is in Group 6 (2,6).

Magnesium loses its two outer electrons to oxygen.
Magnesium becomes [Mg]²⁺ with structure 2,8, like neon.
Oxygen becomes [O]²⁻ with structure 2,8, like neon. Its outer shell has six dots and two crosses.
One magnesium ion and one oxide ion: the formula is MgO.

Worked example 2: potassium sulfide

Potassium is in Group 1 (2,8,8,1). Sulfur is in Group 6 (2,8,6).

Sulfur needs to gain two electrons, but each potassium atom can only give one.
So you need two potassium atoms. Draw both, each giving one cross to sulfur.
You get two [K]⁺ ions (2,8,8, like argon) and one [S]²⁻ ion (2,8,8, like argon).
The total charge is (+1) + (+1) + (−2) = 0. The formula is K₂S.

Working out the charge on an ion

The charge on the ions made by metals in Groups 1 and 2 and non-metals in Groups 6 and 7 relates to the group number. The group number tells you the number of outer electrons, so it tells you how many electrons are lost or gained.

① Groups 1 and 2

Lose 1 or 2 electrons. Group 1 ions are 1+ (e.g. Li⁺). Group 2 ions are 2+ (e.g. Ba²⁺). Charge = group number.

② Group 6

Six outer electrons, so gain 2. Group 6 ions are 2− (e.g. O²⁻). Charge = 8 minus 6.

③ Group 7

Seven outer electrons, so gain 1. Group 7 ions are 1− (e.g. F⁻). Charge = 8 minus 7.

A quick rule: for metals, the charge is positive and equal to the group number. For non-metals in Groups 6 and 7, the charge is negative and equal to 8 minus the group number.

Common mistakes

Saying electrons are shared in ionic bonding. They are transferred. Sharing is covalent.
Getting the sign wrong. Losing negative electrons makes a positive ion. Students often think "lose" means negative.
Saying Group 6 ions are 6−. Oxygen gains 2 electrons, not 6, so the charge is 2−.
Missing brackets or charges on dot and cross diagrams. These are often where the marks are.
Drawing only one atom when two are needed, for example one lithium with one oxygen. Check the charges add up to zero.
Calling the bond a "transfer". The transfer makes the ions. The bond is the attraction between them.

Exam-style question

Magnesium reacts with fluorine to make magnesium fluoride, an ionic compound.

(a) Describe what happens to the electrons when magnesium fluoride forms. [3 marks]

(b) Give the charges on the magnesium ion and the fluoride ion. [2 marks]

Model answer

(a) Each magnesium atom loses its two outer electrons (1). One electron is transferred to each of two fluorine atoms (1). Each fluorine atom gains one electron, so both ions have the electronic structure of a noble gas, 2,8 (1).

(b) Magnesium ion: 2+ (1). Fluoride ion: 1− (1).

Exam tip

Before you draw a dot and cross diagram, work out the charges from the group numbers. Then pick how many of each atom you need so the charges cancel out to zero. Only then start drawing.

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