# Ionic bonding

> Chemistry · CIE A-Level
> Source: https://www.owlsprep.com/study/cie-9701-u3-ionic-bonding/

We cover ionic bond formation, ion charge balancing, giant ionic lattice structure, physical property links, and ionic radius trends fully aligned to CIE 9701 syllabus requirements.

**Prerequisites:** [Electronic configuration of atoms](https://www.owlsprep.com/study/cie-9701-u2-electronic-configurations/); [Periodic table group trends](https://www.owlsprep.com/study/cie-9701-u4-periodic-trends/)

## Learning objectives

- Explain ionic bond formation via full electron transfer between metals and non-metals
- Deduce correct empirical formulae for ionic compounds from known ion charges
- Relate giant ionic lattice structure to characteristic physical properties of ionic solids
- Predict and explain trends in ionic radii across groups and isoelectronic series

## Formation of Ionic Bonds

Ionic bonds form when atoms with low ionisation energy (typically group 1 and 2 metals) transfer their outermost valence electrons to atoms with high electron affinity (typically group 16 and 17 non-metals). This process produces full outer electron shells for both species, generating stable oppositely charged ions.

**Ionic bond** — Strong non-directional electrostatic attraction between a positively charged cation and negatively charged anion

*Notation:* \text{M} + \text{X} \rightarrow \text{M}^+ + \text{X}^-

**Worked example:** Illustrate the full electron transfer process that forms magnesium oxide (MgO) from neutral magnesium and oxygen atoms

1. Step 1: Neutral Mg atom has electron configuration 1s²2s²2p⁶3s², it loses 2 valence electrons to form Mg²⁺ with full outer shell 1s²2s²2p⁶
2. Step 2: Neutral O atom has electron configuration 1s²2s²2p⁴, it gains the 2 lost electrons to form O²⁻ with full outer shell 1s²2s²2p⁶
3. Step 3: Strong electrostatic attraction forms between Mg²⁺ and O²⁻ to create the ionic compound MgO

> **tip**
>
> For A-Level answers, always specify that electron transfer is complete, not partial, to distinguish ionic bonding from polar covalent bonding.

## Deducing Ionic Compound Formulae

All ionic compounds are electrically neutral overall, so total positive charge from cations must exactly cancel total negative charge from anions. For polyatomic ions, you must treat the entire charged group as a single unit, using brackets if multiple copies are required.

**Worked example:** Deduce the correct empirical formula for aluminium sulfate, given Al forms Al³⁺ and sulfate is SO₄²⁻

1. Step 1: Identify charges: Al = +3, SO₄ = -2
2. Step 2: Find lowest common multiple of 3 and 2 = 6, so total positive charge = +6, total negative charge = -6
3. Step 3: Number of Al³⁺ ions = 6 / 3 = 2, number of SO₄²⁻ ions = 6 / 2 = 3
4. Step 4: Write formula with brackets for multiple polyatomic ions: Al₂(SO₄)₃

**Check your understanding**

Test your understanding of charge balancing

1. What is the correct formula for calcium nitrate?

   - CaNO₃
   - Ca(NO₃)₂
   - Ca₂NO₃
   - CaN₂O₆

   *Why:* Ca forms Ca²⁺, nitrate is NO₃⁻, so 2 nitrate ions balance the +2 charge

## Giant Ionic Lattice Structure

Ionic compounds do not exist as discrete molecules. Instead, ions pack into a regular repeating 3D giant ionic lattice, where each ion is surrounded by a fixed number of oppositely charged neighbouring ions called its coordination number. The exact packing arrangement depends on the relative size of the cations and anions.

**Worked example:** Explain why solid sodium chloride has a 6:6 coordination number

1. Step 1: Radius ratio of Na⁺ / Cl⁻ = ~0.52, which falls in the range for octahedral coordination
2. Step 2: Each Na⁺ ion is surrounded by 6 adjacent Cl⁻ ions at equal distance
3. Step 3: Each Cl⁻ ion is in turn surrounded by 6 adjacent Na⁺ ions, giving the 6:6 ratio

**Exam command terms**

CIE uses specific command terms for structure questions:

- **Describe the structure** — You must name the giant ionic lattice, state alternating ions, and give the coordination number *(Sodium chloride has a giant ionic lattice of alternating Na⁺ and Cl⁻ ions with 6:6 coordination)*

## Physical Properties of Ionic Solids

All properties of ionic compounds are directly derived from their strong electrostatic ionic bonds and fixed lattice arrangement. You will be expected to link each property explicitly to the structure in exam answers.

**Worked example:** Explain why magnesium oxide has a much higher melting point (~2800°C) than sodium chloride (~801°C)

1. Step 1: MgO has 2+ and 2- charged ions, while NaCl only has 1+ and 1- charged ions
2. Step 2: Electrostatic force between ions is proportional to the product of their charges, so MgO ionic bonds are far stronger
3. Step 3: Much more thermal energy is required to overcome the stronger bonds in MgO, leading to a far higher melting point

## Ionic Radii Trends

Ionic radius increases down any group of the periodic table, as new electron shells are added further from the nucleus. For an isoelectronic series (ions with identical electron configuration), ionic radius decreases as nuclear charge increases.

| Ion | Nuclear charge | Ionic radius / nm |
| --- | --- | --- |
| N³⁻ | +7 | 0.171 |
| O²⁻ | +8 | 0.140 |
| F⁻ | +9 | 0.133 |
| Na⁺ | +11 | 0.095 |
| Mg²⁺ | +12 | 0.065 |

## Common pitfalls

- **Wrong:** Describing ionic bonds as attraction between two neutral atoms
  - Why it fails: Ionic bonds only exist after electron transfer, between charged ions not neutral atoms
  - Correct: Explicitly state the bond is electrostatic attraction between oppositely charged cations and anions
- **Wrong:** Forgetting brackets around polyatomic ions when writing formulae
  - Why it fails: This leads to incorrect subscript values for atoms inside the polyatomic ion group
  - Correct: Enclose all polyatomic ions in brackets if you have more than one of them in the formula
- **Wrong:** Stating solid ionic compounds conduct electricity
  - Why it fails: Ions are locked in a fixed lattice and cannot move to carry charge
  - Correct: Specify conductivity only occurs when the ionic compound is molten or dissolved in water
- **Wrong:** Claiming ionic compounds form simple covalent molecules
  - Why it fails: Ionic lattices are giant extended structures with no discrete molecular units
  - Correct: Refer to the structure as a giant ionic lattice, never use the term 'ionic molecule'
- **Wrong:** Saying ionic radius increases across an isoelectronic series
  - Why it fails: Higher nuclear charge pulls the same number of electrons closer to the nucleus
  - Correct: State ionic radius decreases as nuclear charge rises for ions with identical electron configuration

## Cheatsheet

| Property | Ionic solid behaviour | Explanation |
| --- | --- | --- |
| Melting/boiling point | Very high | Strong electrostatic bonds require large energy to break |
| Electrical conductivity | Only conductive molten/dissolved | Mobile free ions act as charge carriers |
| Solubility | Soluble in polar solvents like water | Polar water molecules hydrate separated ions |
| Brittleness | Brittle, shatters on impact | Layers shift, aligning like charges that repel each other |

## What's next

Mastering ionic bonding gives you a foundational framework to compare against other chemical bond types, which is a core 6-8 mark extended question topic in almost every CIE A-Level Chemistry paper. You will next explore covalent bonding, where electrons are shared rather than transferred, before moving on to calculate lattice energy values using Born-Haber cycles, and then compare the properties of metallic bonding to ionic and covalent structures. This full set of bond type comparisons is heavily weighted for Paper 2 and Paper 4 structured questions, so ensure you can clearly distinguish the formation, structure and properties of each bond family.

- [Metallic Bonding](https://www.owlsprep.com/study/cie-9701-u3-metallic-bonding/)

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