# Ionic Bonding

> CIE IGCSE Chemistry · 0620 2026-2028
> Source: https://www.owlsprep.com/study/cie-0620-u2-ionic-bonding/

This guide covers all core and extended ionic bonding content for CIE IGCSE Chemistry 0620 Unit 2, aligned to the 2026-2028 syllabus, including dot-cross diagrams, formula prediction and lattice properties.

**Prerequisites:** [Atomic structure and electron configuration](https://www.owlsprep.com/study/cie-0620-u2-atomic-structure/); [Periodic table group properties](https://www.owlsprep.com/study/cie-0620-u3-periodic-table-groups/)

## Learning objectives

- Define ionic bonding as electrostatic attraction between oppositely charged ions formed by electron transfer from metals to non-metals
- Draw accurate dot-cross diagrams for simple ionic compounds
- Predict neutral formulae of ionic compounds using ion charges
- Explain the effect of ionic charge and radius on lattice strength (Extended only)
- Write balanced half-equations for ion formation (Extended only)

## 1. What is Ionic Bonding (Core)

**Ionic Bond** — The strong electrostatic force of attraction between positively charged cations and negatively charged anions, formed when electrons are transferred from metal atoms to non-metal atoms.

*Example:* Sodium (Group 1 metal) loses 1 outer electron to chlorine (Group 17 non-metal), forming Na⁺ and Cl⁻ ions that are attracted to each other.

Ionic bonding only occurs between metals (left of the periodic table) and non-metals (right of the periodic table). Metals lose electrons to form positive cations, while non-metals gain those electrons to form negative anions. This electron transfer gives both ions a full outer shell of electrons, matching the stable noble gas electron configuration.

**Worked example:** Explain how ionic bonding forms between lithium and fluorine.

1. Lithium is a Group 1 metal with 1 outer electron; fluorine is a Group 17 non-metal with 7 outer electrons.
2. Lithium transfers its single outer electron to fluorine.
3. Lithium now has a full outer shell, forming a +1 cation (Li⁺). Fluorine gains the electron, has a full outer shell, forming a -1 anion (F⁻).
4. The strong electrostatic attraction between Li⁺ and F⁻ is the ionic bond holding the compound together.

> **Exam tip:** Always specify ionic bonding is *electrostatic attraction* between oppositely charged ions in exam answers, not just 'attraction' to get full marks.

## 2. Dot-Cross Diagrams for Ionic Compounds (Core)

**Dot-Cross Diagram** — A simplified diagram showing the arrangement of outer shell electrons in atoms, ions or molecules, where electrons from one atom are drawn as dots and electrons from another as crosses.

For ionic compounds, every ion must be drawn inside square brackets, with the overall charge of the ion written in the top right corner outside the brackets. You do not need to draw lines between ions, as ionic bonds are electrostatic attractions not shared electron pairs.

**Worked example:** Draw the dot-cross diagram for sodium chloride (NaCl).

1. Draw the Na⁺ ion: outer shell is empty (since it lost its 1 outer electron), enclose in square brackets, write +1 outside the top right corner.
2. Draw the Cl⁻ ion: outer shell has 8 electrons: 7 original (drawn as crosses) and 1 gained from Na (drawn as a dot).
3. Enclose the Cl⁻ ion in square brackets, write -1 outside the top right corner.
4. Place the two ions next to each other with no connecting line between them.

> **Exam tip:** Mark the transferred electron as a different symbol (dot vs cross) to clearly show which atom it came from, this guarantees you full marks for diagram questions.

## 3. Predicting Ionic Compound Formulae (Core)

All ionic compounds are electrically neutral, meaning the total positive charge from cations equals the total negative charge from anions. You can use the 'swap and drop' method to find the correct formula quickly.

**Worked example:** Find the formula of aluminium oxide.

1. Identify ion charges: aluminium is Group 13, so charge = +3 (Al³⁺); oxygen is Group 16, so charge = -2 (O²⁻).
2. Swap the numerical values of the charges to become subscripts for the opposite ion: Al gets subscript 2, O gets subscript 3.
3. Write the formula: Al₂O₃.
4. Check neutrality: 2*(+3) + 3*(-2) = 6 - 6 = 0, so the formula is correct.

> **Exam tip:** If your subscripts share a common factor, simplify them: e.g. Mg₂O₂ simplifies to MgO, as ionic formulae are always written in the smallest whole number ratio.

## 4. Extended Only: Advanced Ionic Bonding Concepts

> **Extended Only Content**
>
> This section is only tested in Extended tier papers (Paper 2 and 4). Core tier candidates can skip this section.

**Ionic Lattice** — A giant regular 3D structure of alternating positive and negative ions held together by strong ionic bonds.

*Example:* Sodium chloride forms a cubic lattice with equal numbers of Na⁺ and Cl⁻ ions.

The strength of an ionic lattice depends on two factors: 1) ion charge: higher charge ions create stronger electrostatic attraction, leading to higher melting points; 2) ionic radius: smaller ions pack closer together, so attraction is stronger, leading to higher melting points.

**Worked example:** Explain why magnesium oxide (MgO) has a much higher melting point than sodium chloride (NaCl).

1. Identify ion charges: Mg²⁺ and O²⁻ in MgO; Na⁺ and Cl⁻ in NaCl.
2. Ions in MgO have double the charge of ions in NaCl, so electrostatic attraction between MgO ions is much stronger.
3. More energy is required to overcome these stronger ionic bonds, so MgO has a higher melting point.

Extended candidates also need to write balanced half-equations for ion formation, which show the gain or loss of electrons (written as e⁻) by an atom.

**Worked example:** Write the half-equations for the formation of ions in magnesium chloride (MgCl₂).

1. Magnesium loses 2 electrons to form a Mg²⁺ cation (oxidation):
2. $$Mg \rightarrow Mg^{2+} + 2e^-$$
3. Two chlorine atoms each gain 1 electron to form 2 Cl⁻ anions (reduction):
4. $$Cl_2 + 2e^- \rightarrow 2Cl^-$$
5. Check that total charge is equal on both sides of each half-equation to confirm it is balanced.

> **Exam tip:** For oxidation half-equations (electron loss), electrons are written on the right of the arrow. For reduction half-equations (electron gain), electrons are written on the left of the arrow.

## Common pitfalls

- **Wrong:** Describing ionic bonding as just the transfer of electrons
  - Why it fails: Electron transfer forms the ions, but the ionic bond itself is the electrostatic attraction between the charged ions.
  - Correct: First explain electron transfer to form ions, then define the ionic bond as the electrostatic attraction between oppositely charged ions.
- **Wrong:** Drawing lines between ions in ionic dot-cross diagrams
  - Why it fails: Lines represent shared covalent electrons, not electrostatic ionic attractions.
  - Correct: Draw each ion separately in square brackets with charges, no connecting lines between ions.
- **Wrong:** Writing unbalanced ionic formulae e.g. AlO instead of Al₂O₃
  - Why it fails: Ionic compounds are always neutral, so total positive charge must equal total negative charge.
  - Correct: Use the swap and drop method, then check that the sum of all ion charges equals zero.
- **Wrong:** Writing ion charges inside square brackets in dot-cross diagrams
  - Why it fails: The charge is a property of the whole ion, not just the electrons inside the shell.
  - Correct: Write the charge in the top right corner outside the square brackets surrounding the ion.
- **Wrong:** Ignoring ionic radius when explaining melting point differences (Extended)
  - Why it fails: Both ion charge and ionic radius affect the strength of electrostatic attraction in the lattice.
  - Correct: Mention both factors if relevant: e.g. LiF has a higher melting point than KF because Li⁺ is smaller than K⁺, so attraction is stronger.

## Cheatsheet

| Concept | Core Requirement | Extended Requirement |
| --- | --- | --- |
| Ionic Bond Definition | Electrostatic attraction between oppositely charged ions, formed by electron transfer from metal to non-metal | Same as Core, plus link bond strength to ion charge and radius |
| Dot-Cross Diagrams | Draw outer shell electrons, square brackets, charge outside brackets | Same as Core, plus diagrams for multiply charged ion compounds |
| Ionic Formula | Predict neutral formula using swap and drop method | Same as Core, plus formulae for transition metal ion compounds |
| Ion Formation | Identify which atom loses/gains electrons | Write balanced half-equations for ion formation |
| Lattice Properties | Recall high melting points, conduct when molten/dissolved | Explain effect of ion charge and radius on melting point |

## What's next

Now that you have mastered ionic bonding, move on to covalent bonding, the next key type of chemical bonding in the CIE IGCSE Chemistry 0620 syllabus. Practice past paper questions on ionic bonding to consolidate your knowledge, focusing on dot-cross diagram and formula prediction questions that appear frequently in both Core and Extended papers. For Extended candidates, make sure you practice half-equation writing and lattice strength explanation questions, which are common in Paper 4 structured assessments. You can also revise periodic table group trends to predict ion charges faster and more accurately.

---

From [OwlsPrep](https://www.owlsprep.com) — free study guides for A-Level, IB, AP and IGCSE, written against the official syllabus. Canonical page: https://www.owlsprep.com/study/cie-0620-u2-ionic-bonding/