You are viewing a free preview of this lesson.
Subscribe to unlock all 10 lessons in this course and every other course on LearningBro.
This lesson covers ionic bonding, one of the three main types of chemical bonding you need to understand for the AQA GCSE Chemistry specification (4.2.1). Ionic bonding occurs between metals and non-metals and involves the transfer of electrons. Understanding how and why ions form is fundamental to explaining the properties of ionic compounds, which are covered in the next lesson.
Ionic bonding is the electrostatic force of attraction between oppositely charged ions. It occurs when atoms transfer electrons from one to another, forming positive and negative ions that are held together by strong electrostatic forces.
Ionic bonds typically form between metals (which lose electrons) and non-metals (which gain electrons). The driving force behind ionic bonding is that atoms want to achieve a full outer shell of electrons — the same electron configuration as a noble gas. This is sometimes called the octet rule (having 8 electrons in the outer shell, or 2 for the first shell).
Exam Tip: The definition of ionic bonding is a common 1-mark question. Learn it precisely: "Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions." Do not say "between a metal and a non-metal" — that describes when it happens, not what it is.
Metal atoms have a small number of electrons in their outer shell (typically 1, 2, or 3). They achieve a full outer shell by losing these outer electrons. When a metal atom loses electrons, it has more protons than electrons, giving it an overall positive charge.
| Metal | Group | Electrons Lost | Ion Formed | Electron Configuration |
|---|---|---|---|---|
| Sodium (Na) | 1 | 1 | Na+ | 2, 8 (same as neon) |
| Magnesium (Mg) | 2 | 2 | Mg2+ | 2, 8 (same as neon) |
| Aluminium (Al) | 3 | 3 | Al3+ | 2, 8 (same as neon) |
| Potassium (K) | 1 | 1 | K+ | 2, 8, 8 (same as argon) |
| Calcium (Ca) | 2 | 2 | Ca2+ | 2, 8, 8 (same as argon) |
Non-metal atoms have outer shells that are close to being full (typically 5, 6, or 7 electrons). They achieve a full outer shell by gaining electrons. When a non-metal atom gains electrons, it has more electrons than protons, giving it an overall negative charge.
| Non-Metal | Group | Electrons Gained | Ion Formed | Electron Configuration |
|---|---|---|---|---|
| Chlorine (Cl) | 7 | 1 | Cl- | 2, 8, 8 (same as argon) |
| Oxygen (O) | 6 | 2 | O2- | 2, 8 (same as neon) |
| Fluorine (F) | 7 | 1 | F- | 2, 8 (same as neon) |
| Sulfur (S) | 6 | 2 | S2- | 2, 8, 8 (same as argon) |
| Nitrogen (N) | 5 | 3 | N3- | 2, 8 (same as neon) |
Exam Tip: The charge on an ion is directly related to the group number. Group 1 metals form 1+ ions, Group 2 form 2+ ions, Group 6 non-metals form 2- ions, and Group 7 form 1- ions. This pattern is a quick way to predict the charge of any main group ion.
Let us consider the formation of sodium chloride (NaCl) as a worked example.
graph LR
A["Na atom<br/>2, 8, 1"] -->|"Loses 1 electron"| B["Na+ ion<br/>2, 8"]
C["Cl atom<br/>2, 8, 7"] -->|"Gains 1 electron"| D["Cl- ion<br/>2, 8, 8"]
B -->|"Electrostatic attraction"| E["NaCl<br/>Ionic compound"]
D -->|"Electrostatic attraction"| E
style A fill:#3498db,color:#fff
style B fill:#e74c3c,color:#fff
style C fill:#27ae60,color:#fff
style D fill:#8e44ad,color:#fff
style E fill:#f39c12,color:#fff
Now consider magnesium oxide (MgO), where two electrons are transferred.
In the exam you will be asked to draw dot-and-cross diagrams to show ionic bonding. These diagrams show only the outer shell electrons, using dots for one atom and crosses for the other.
Exam Tip: When drawing dot-and-cross diagrams for ionic bonding, always include the square brackets and the charge. If you forget the charge notation (e.g. [Na]+ and [Cl]-), you will lose marks even if the rest of the diagram is correct.
The formula of an ionic compound must balance so that the overall charge is zero. This is because the total positive charge from the cations must equal the total negative charge from the anions.
| Compound | Cation | Anion | Formula | Explanation |
|---|---|---|---|---|
| Sodium chloride | Na+ | Cl- | NaCl | 1+ and 1- balance |
| Magnesium oxide | Mg2+ | O2- | MgO | 2+ and 2- balance |
| Magnesium chloride | Mg2+ | Cl- | MgCl2 | 2+ needs two 1- ions |
| Sodium oxide | Na+ | O2- | Na2O | Two 1+ ions balance 2- |
| Calcium chloride | Ca2+ | Cl- | CaCl2 | 2+ needs two 1- ions |
| Aluminium oxide | Al3+ | O2- | Al2O3 | Two 3+ balance three 2- |
Ionic bonding occurs because the resulting ionic compound is more energetically stable than the individual atoms. The process releases energy overall, as the electrostatic attraction between the oppositely charged ions releases a large amount of energy (known as lattice energy). The atoms achieve the stable electron configuration of a noble gas, which is the lowest energy state for their electrons.
Exam Tip: A 6-mark question may ask you to "describe and explain the formation of ionic bonds in sodium chloride." Structure your answer clearly: (1) state what ionic bonding is, (2) describe the electron configurations of Na and Cl, (3) explain the electron transfer, (4) state the charges on the resulting ions, (5) explain the electrostatic attraction. Use a dot-and-cross diagram to support your answer if asked.
Let us work through a second example where both ions carry 2+/2- charges. A magnesium atom (Mg) has the electron configuration 2, 8, 2. An oxygen atom (O) has the electron configuration 2, 6.
The balanced symbol equation for the formation of magnesium oxide from its elements is:
2Mg(s) + O2(g) → 2MgO(s)
Because both ions carry a 2+ or 2- charge, the electrostatic attraction in MgO is much stronger than in NaCl — this is why MgO has a much higher melting point (2,852 °C) than NaCl (801 °C). The same principle explains why calcium oxide (CaO) is used to line furnaces, while sodium chloride would simply melt. At GCSE you should always be ready to link the size of the ionic charges to the energy required to disrupt a giant ionic lattice: the greater the product of the charges, the stronger the electrostatic attraction and the more energy is required to melt the solid.
Common mistake: Students often confuse ionic bonding with covalent bonding at GCSE. Remember: ionic = transfer of electrons (metal + non-metal); covalent = sharing of electrons (non-metal + non-metal). Never describe an ionic bond as a "shared pair".
Common mistake: Writing the formula as "Mg2O2" is wrong. The subscripts in an ionic formula give the simplest whole-number ratio of ions. MgO already has a 1:1 ratio, so no subscripts are needed. Similarly, the formula for calcium chloride is CaCl2 (not Ca2Cl4) because the ratio of Ca2+ to Cl- must be 1:2 for charges to balance.
Exam-style question: Describe how a sodium atom and a chlorine atom form an ionic bond. (4 marks)
Grade 4–5 answer: The sodium atom gives its outer electron to the chlorine atom. Sodium becomes a positive ion and chlorine becomes a negative ion. The two ions are then attracted to each other and this holds them together as sodium chloride.
Grade 8–9 answer: The sodium atom transfers its single outer-shell electron to the chlorine atom so that both atoms achieve a full outer shell with the electron configuration of a noble gas. Sodium becomes an Na+ cation (2, 8) and chlorine becomes a Cl- anion (2, 8, 8). A strong electrostatic attraction then acts between the oppositely charged ions in all directions, locking them into a giant ionic lattice of Na+ and Cl- ions. This electrostatic attraction is the ionic bond itself, and it is not restricted to a single pair of ions — each Na+ is surrounded by six Cl- ions and vice versa.
AQA alignment: This content is aligned with AQA GCSE Chemistry (8462) specification section 5.2 Bonding, structure, and the properties of matter — specifically 5.2.1.1 Chemical bonds, 5.2.1.2 Ionic bonding, and 5.2.1.3 Ionic compounds. Assessed on Paper 1.