Ionic Bonding and Ionic Compounds — AQA A-level Chemistry

SPEC 3.1.3.1 & 3.1.3.3 - BONDING

Ionic & Metallic Bonding

Electrostatic attractions in ionic lattices, compound ions, and the delocalised electron model of metals

Why Bonding Matters

Chemical bonding determines the structure and properties of every substance around us. Understanding bonding allows us to predict how materials will behave.

Learning Objectives

By the end of this lesson, you will be able to:

Define ionic bonding as electrostatic attraction between oppositely charged ions in a lattice

Predict ion charges using position in the Periodic Table

Write formulas for ionic compounds including those with compound ions

Describe metallic bonding using the delocalised electron model

Prior Knowledge

This lesson builds on your understanding of electron configuration from Topic 1. Recall that atoms form ions to achieve stable noble gas configurations.

What is Ionic Bonding?

Definition

Ionic bonding involves electrostatic attraction between oppositely charged ions in a lattice.

Ionic bonding occurs between metals (on the left of the periodic table) and non-metals (on the right-hand side). The process involves:

1

Electron transfer - Metal atoms lose electrons to form positive ions (cations)

2

Ion formation - Non-metal atoms gain electrons to form negative ions (anions)

3

Electrostatic attraction - Opposite charges attract, holding ions together

Key Concept

Ionic compounds do NOT exist as individual molecules. The ions are arranged in a giant ionic lattice - a regular 3D arrangement extending in all directions. Each ion is surrounded by ions of opposite charge. This is called non-directional bonding because the attraction occurs in all directions.

Important

The electrostatic attraction in ionic bonding is very strong and requires a lot of energy to overcome. This is why ionic compounds have high melting and boiling points.

Forming Cations

Metals lose electrons from their outer shell to form positively charged ions (cations). Group 1 metals have 1 outer electron, so they lose 1 electron to achieve a stable noble gas configuration.

Key Points

The loss of a negatively charged electron leaves the ion with more protons than electrons, giving it an overall positive charge. Square brackets show the charge is evenly distributed over the entire ion. The sodium ion (Na⁺) now has a stable electron configuration like neon: [2,8]⁺.

Exam Tip

You can use the group number on the periodic table to work out how many electrons an atom will lose. Group 1 metals lose 1 electron (1+ charge), Group 2 metals lose 2 electrons (2+ charge), and Group 3 metals lose 3 electrons (3+ charge).

Forming Anions

Non-metals gain electrons into their outer shell to form negatively charged ions (anions). Group 16 non-metals have 6 outer electrons, so they gain 2 electrons to achieve a full outer shell.

Key Points

The gain of negatively charged electrons (shown in red) from another atom gives the ion more electrons than protons, resulting in an overall negative charge. The oxide ion (O²⁻) now has a stable electron configuration like neon: [2,8]²⁻.

Noble Gas Configurations

Once atoms become ions, their electronic configurations are the same as a noble gas. For example, a chloride ion (Cl⁻) has the configuration [2,8,8]⁻ (same as argon). A potassium ion (K⁺) also has [2,8,8]⁺ (same as argon).

Ionic Bond Formation

When a metal reacts with a non-metal, electrons are transferred from the metal to the non-metal. The resulting cations and anions are attracted to each other by electrostatic forces, forming an ionic compound.

Key Points

The Na⁺ and Cl⁻ ions are attracted through strong electrostatic forces. This forms the ionic compound sodium chloride (NaCl). The ionic bond is a non-directional bond - the attraction occurs in all directions. These forces require a lot of energy to overcome, which explains the high melting points of ionic compounds.

Predicting Ion Charges

You can predict the charge on a simple ion using its position in the Periodic Table:

Why These Charges?

Atoms lose or gain electrons to achieve a full outer shell (noble gas configuration). Group 1 metals have 1 electron to lose; Group 7 non-metals need 1 electron to complete their shell.

The Ionic Lattice Structure

Ions form a lattice structure — an evenly distributed crystalline arrangement. Ions are arranged in a regular repeating pattern so that positive charges cancel out negative charges. The attraction between cations and anions occurs in all directions — each ion is attracted to all of the oppositely charged ions around it.

Key Points

The electrostatic forces of attraction bring cations and anions together to form the lattice structure. Ions are arranged in a regular repeating pattern so that positive charges cancel out negative charges - the lattice has a neutral charge overall. The attraction between cations and anions occurs in all directions - each ion is attracted to all of the oppositely charged ions around it.

Ion Sizes

In an ionic lattice, the negative ions (anions) are larger than the positive ions (cations). This is because anions have gained electrons (more electron-electron repulsion) while cations have lost electrons. The type of lattice formed depends on the relative sizes of the ions, which are arranged in an alternating fashion.

Worked Example01

Describe the formation of magnesium oxide (MgO) from its atoms, including the electron configurations of the ions formed.

1

Identify the atoms: Magnesium (Group 2, electron config: 2,8,2) and Oxygen (Group 6, electron config: 2,6)

2

Electron transfer: Mg loses 2 electrons → Mg²⁺ (config: 2,8 - same as Ne)

3

Electron gain: O gains 2 electrons → O²⁻ (config: 2,8 - same as Ne)

4

Ionic bonding: Strong electrostatic attraction forms between Mg²⁺ and O²⁻ ions in a giant ionic lattice.

Summary: Mg → Mg²⁺ + 2e⁻ and O + 2e⁻ → O²⁻

Compound Ions (Polyatomic Ions)

Compound ions are groups of atoms covalently bonded together that carry an overall charge. You must memorise these for your exams:

NameFormulaCharge

AmmoniumNH₄⁺1+

HydroxideOH⁻1−

NitrateNO₃⁻1−

CarbonateCO₃²⁻2−

SulfateSO₄²⁻2−

Exam Essential

These compound ions appear frequently in exam questions. Note that sulfate and carbonate have a 2− charge, while nitrate and hydroxide have a 1− charge.

Constructing Ionic Formulas

To write the formula of an ionic compound, the total positive charge must equal the total negative charge (the compound must be electrically neutral).

Method

Write the symbols and charges of each ion

Find the ratio that balances the charges

Write the formula (cation first, then anion)

Use brackets for compound ions if more than one is needed

Cross-Over Method

A quick method: "cross over" the numbers from the charges to become subscripts. For Ca²⁺ and Cl⁻: the 2 from calcium becomes the subscript for chlorine, and the 1 from chlorine becomes the subscript for calcium → CaCl₂

Worked Example02

Write the formula for aluminium oxide.

1

Identify the ions: Al³⁺ (Group 3) and O²⁻ (Group 6)

2

Balance charges: Need total positive = total negative

3

Find ratio: 2 × Al³⁺ = +6 and 3 × O²⁻ = −6 ✓

4

LCM method: LCM of 3 and 2 is 6. Need 2 Al³⁺ ions and 3 O²⁻ ions.

Answer: Al₂O₃

Worked Example03

Write the formula for calcium hydroxide and aluminium sulfate.

1

Calcium hydroxide: Ca²⁺ and OH⁻

2

Need 2 hydroxide ions to balance 1 calcium ion: Ca(OH)₂

3

Aluminium sulfate: Al³⁺ and SO₄²⁻

4

LCM of 3 and 2 is 6. Need 2 Al³⁺ (+6) and 3 SO₄²⁻ (−6)

Answers: Ca(OH)₂ and Al₂(SO₄)₃

Note: Brackets around the compound ion indicate more than one is present.

What is Metallic Bonding?

Definition

Metallic bonding involves electrostatic attraction between delocalised electrons and positive ions arranged in a lattice.

In metallic bonding:

Metal atoms lose their outer shell electrons

These electrons become delocalised (free to move throughout the structure)

The electrons form a "sea" of negative charge

Positive metal ions are held in fixed positions in a lattice

The positive charges repel each other and keep the neatly arranged lattice in place

There are very strong electrostatic forces between the positive ions and the delocalised electrons - this holds the structure together

Key Terminology

Delocalised means the electrons are not associated with any particular atom - they are not bound to their original atom and can move freely throughout the entire metal structure. This is sometimes called a "sea of electrons".

The Metallic Structure

Metal atoms with electrons in the outer shell

When metals are closely packed, their outer electrons dissociate forming positive metal ions.

Positive metal ions in a "sea" of delocalised electrons

Once the outer electrons are delocalised, the positive metal ions are suspended in the "sea" of electrons. The delocalised electrons are free to move through the metal. Metallic bonding involves very strong electrostatic forces of attraction between the positive metal ions and the delocalised electrons.

Metallic Bond Strength

The strength of metallic bonding depends on two main factors:

Factor 1

Number of delocalised electrons: More delocalised electrons = stronger metallic bonding. For example, Mg (2 electrons per atom) has stronger metallic bonding than Na (1 electron per atom).

Factor 2

Size of the metal ion: Smaller ions = stronger metallic bonding. The delocalised electrons are closer to the nucleus, so the electrostatic attraction is stronger.

Exam Application

When comparing melting points of metals, consider both factors. Mg has a higher melting point than Na because Mg has more delocalised electrons (2 vs 1) AND smaller ions (higher charge density).

Comparing Ionic and Metallic Bonding

FeatureIonic BondingMetallic Bonding

Particles involvedPositive and negative ionsPositive ions and delocalised electrons

Type of attractionElectrostatic (ion to ion)Electrostatic (ion to electron sea)

StructureGiant ionic latticeGiant metallic lattice

Elements involvedMetal + non-metalMetals only

Electron behaviourTransferred from metal to non-metalDelocalised (free to move)

Common Feature

Both ionic and metallic bonding involve electrostatic attraction and both form giant lattice structures. The key difference is what the positive ions are attracted to.

Practice

Question 1: Write the formulas for the following ionic compounds:

(a) Potassium sulfate

(b) Ammonium carbonate

(c) Iron(III) hydroxide

(a)

K⁺ and SO₄²⁻: Need 2 K⁺ to balance one SO₄²⁻ → K₂SO₄

(b)

NH₄⁺ and CO₃²⁻: Need 2 NH₄⁺ to balance one CO₃²⁻ → (NH₄)₂CO₃

(c)

Fe³⁺ and OH⁻: Need 3 OH⁻ to balance one Fe³⁺ → Fe(OH)₃

Practice

Question 2: Explain why magnesium has a higher melting point than sodium in terms of metallic bonding.

1

Magnesium has more delocalised electrons per atom (2 vs 1 for sodium)

2

The Mg²⁺ ion is smaller than the Na⁺ ion (higher charge density)

3

Therefore there is stronger electrostatic attraction between the positive ions and the delocalised electrons

4

More energy is needed to overcome these stronger metallic bonds → higher melting point

Practice

Question 3: Define the following terms. Include the key words that would appear in an exam mark scheme.

(a) Ionic bonding

(b) Metallic bonding

(a)

Ionic bonding is the electrostatic attraction between oppositely charged ions in a lattice.

(b)

Metallic bonding is the electrostatic attraction between positive ions (or metal ions) and delocalised electrons in a lattice.

Practice

Question 4: Sodium chloride has a melting point of 801°C. Explain, in terms of structure and bonding, why sodium chloride has a high melting point.

1

Sodium chloride has a giant ionic lattice structure

2

There are strong electrostatic attractions between the oppositely charged Na⁺ and Cl⁻ ions

3

These attractions act in all directions throughout the lattice

4

A lot of energy is required to overcome these strong ionic bonds

Summary

Key Definitions

TermDefinition

Ionic bondingElectrostatic attraction between oppositely charged ions in a lattice

Metallic bondingElectrostatic attraction between delocalised electrons and positive ions in a lattice

Delocalised electronsElectrons not associated with any particular atom, free to move throughout the structure

Compound Ions to Memorise

1+ charge: NH₄⁺ (ammonium)

1− charge: OH⁻ (hydroxide), NO₃⁻ (nitrate)

2− charge: CO₃²⁻ (carbonate), SO₄²⁻ (sulfate)

Exam Tips

Always state "electrostatic attraction" when describing bonding

For ionic: specify "between oppositely charged ions"

For metallic: specify "between positive ions and delocalised electrons"

Use brackets for compound ions when more than one is needed in a formula