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Module 1 · L7 of 21 30 min ⚡ +50 XP in Learn · +25 to complete Year 11 · Module 1 · IQ4

Ionic Bonding and Properties

Today's hook, Sodium explodes on contact with water. Chlorine has been used as a chemical weapon. Bond them together and you sprinkle the result on your chips.
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Warm up and recall

Warm up first

Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.

Worksheets

Practise this lesson

Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.

Build Foundations & guided practice Apply Application practice Master Mastery challenge Exam Exam-style questions
01
Recall, your gut answer first
+5 XP warm-up

Sodium chloride (NaCl) and magnesium oxide (MgO) both form ionic lattices. NaCl has a melting point of 801°C, while MgO melts at 2852°C. Both are made of positive and negative ions. What could explain why MgO requires so much more energy to melt?

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What you'll master, and the words for it

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What you'll master
Know

Key facts

  • Ionic compounds form giant 3D lattices (not discrete molecules); NaCl has 6:6 coordination in a face-centred cubic structure
  • Lattice energy depends on ion charge (higher → stronger) and ion size (smaller → stronger) according to Coulomb's Law
Understand

Concepts

  • Why ionic compounds often have high melting points (relative to molecular substances), are hard but brittle, and conduct only when molten or dissolved, and why the exact melting point and the solubility still depend on the specific ions (charge, size, lattice and hydration)
  • How Coulomb's Law (F ∝ q₁q₂/r²) explains why MgO has a much higher MP than NaCl
Can do

Skills

  • Compare melting points of ionic compounds using ion charge and ionic radius
  • Explain the conductivity of ionic compounds in solid, molten, and dissolved states using the ionic model
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Ionic Lattice Structure

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Ionic Lattice Structure
core concept

What is an ionic lattice?

Ionic compounds do not form discrete molecules. Instead, ions arrange themselves into a giant regular 3D lattice where every cation is surrounded by anions and every anion is surrounded by cations. The arrangement maximises attractive forces and minimises repulsive forces.

NaCl as the model lattice: In sodium chloride, each Na⁺ ion is surrounded by exactly 6 Cl⁻ ions (and vice versa) in an octahedral arrangement. This is the face-centred cubic (rock salt) structure. The ratio of Na⁺:Cl⁻ = 1:1 reflects the formula NaCl. The entire crystal is one giant "molecule", the formula simply gives the simplest ratio of ions.

Lattice Energy and Melting Point

The lattice energy is the energy holding the ions together. It depends on ion charge and ion size:

  • Higher charge → stronger attraction → higher lattice energy → higher MP. MgO (Mg²⁺ and O²⁻, charges ±2) has a much higher MP (2852°C) than NaCl (Na⁺ and Cl⁻, charges ±1, MP 801°C).
  • Smaller ion → ions are closer together → stronger attraction → higher lattice energy → higher MP. LiF has a higher MP than CsI because Li⁺ and F⁻ are small and can get very close.
Na⁺ and Cl⁻ ions alternate in a giant 3D lattice with 6:6 coordination.

Ionic compounds form giant 3D lattices, not discrete molecules. In NaCl, each Na⁺ is surrounded by 6 Cl⁻ and vice versa (6:6 coordination). Lattice energy increases with higher ion charge and smaller ion size: MgO (Mg²⁺/O²⁻, ±2, MP 2852°C) vs NaCl (Na⁺/Cl⁻, ±1, MP 801°C), higher charge → much stronger electrostatic attraction.

Pause, copy the highlighted definition into your book before moving on.

Fill the blanks: drag each token into the matching gap.

electrostatic six 3D lattice simplest ratio

In NaCl, ions form a giant ___ rather than discrete molecules. Each Na⁺ is surrounded by ___ Cl⁻ ions (and vice versa), held together by ___ attraction. The formula NaCl simply gives the ___ of ions.

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Physical Properties of Ionic Compounds, Explained

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Physical Properties of Ionic Compounds, Explained
core concept

Melting point

Observation: High (hundreds to thousands of °C)
Explanation in terms of lattice structure: Strong electrostatic forces in the lattice require large amounts of energy to overcome. Many bonds must break simultaneously to allow ions to move freely.

Hardness

Observation: Hard (high scratch resistance)
Explanation in terms of lattice structure: The rigid, strongly bonded lattice resists deformation. Large energy needed to displace ions from their equilibrium positions.

Brittleness

Observation: Shatters under impact
Explanation in terms of lattice structure: When a force shifts ion layers, like charges align → strong repulsion → lattice cleaves. No way to redistribute force like metals can.

Conductivity (solid)

Observation: None
Explanation in terms of lattice structure: Ions are fixed in lattice positions, cannot move to carry charge.

Conductivity (molten)

Observation: Excellent
Explanation in terms of lattice structure: Lattice is broken, ions become mobile and can carry charge through the liquid.

Conductivity (dissolved)

Observation: Excellent
Explanation in terms of lattice structure: Water molecules surround and separate ions (hydration), ions become mobile in solution.

Solubility

Observation: Many are soluble in water (polar solvent)
Explanation in terms of lattice structure: Water is polar, its δ+ and δ− ends are attracted to anions and cations respectively, pulling them away from the lattice (hydration energy releases energy).
Key exam phrase, conductivity: "Ionic compounds do not conduct electricity as solids because the ions are fixed in the lattice. When melted or dissolved in water, the ions become mobile and can carry charge, so the substance conducts." This should be memorised and reproduced precisely in exam answers.

We just saw the ionic lattice structure and how lattice energy depends on ion charge and size. That raises a question: how does this structure directly cause the distinctive physical properties of ionic compounds? This card answers it → each property (MP, hardness, conductivity) is explained by whether the lattice holds ions fixed or allows them to move.

Ionic compounds: high MP (strong electrostatic forces require large energy to overcome); hard but brittle (rigid lattice resists deformation but fractures when like-charge layers align); no solid conductivity (ions fixed in lattice); conduct when molten or dissolved (ions become mobile, water molecules hydrate and separate them from the lattice).

Add the highlighted property explanations to your notes before the check below.

Odd one out three of these statements correctly describe an ionic compound. Which one does not?

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Comparing Different Ionic Compounds

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Comparing Different Ionic Compounds
core concept

Using the principles of lattice energy (charge and size), you can predict and explain differences between ionic compounds:

Comparison table migration of NaCl vs MgO lattice energy diagram

NaCl

Ions and charges: Na⁺, Cl⁻ (±1)
MP (°C): 801
Relative lattice energy: Moderate

MgO

Ions and charges: Mg²⁺, O²⁻ (±2)
MP (°C): 2852
Relative lattice energy: Very high

LiF

Ions and charges: Li⁺, F⁻ (±1, small ions)
MP (°C): 848
Relative lattice energy: High (higher than NaCl due to smaller ions)

CsI

Ions and charges: Cs⁺, I⁻ (±1, large ions)
MP (°C): 632
Relative lattice energy: Lower than NaCl (larger ions, more distant)
Predicting MP trend: Compare charges first (higher charge → much higher MP). If charges are equal, compare ion sizes (smaller ions → higher MP). MgO vs NaCl: charge effect dominates, MgO is almost 4× the MP. LiF vs CsI: same charges, size effect, LiF has higher MP.

We just saw why ionic compounds have high MPs, hardness, and conditional conductivity. That raises a question: when comparing two ionic compounds in exam data questions, how do you predict which has a higher MP? This card answers it → compare ion charges first, then ion sizes.

To compare two ionic compounds: check ion charges first (higher charge → much stronger attraction → higher MP); if charges are equal, compare ion sizes (smaller ions → closer together → stronger attraction → higher MP). Example: LiF > CsI (same charges ±1, size effect dominates); MgO > NaCl (charge effect dominates, ±2 vs ±1).

Pause, write the highlighted comparison rule into your book.

Two truths, one lie comparing lattice energies. Pick the lie.

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Short Answer Questions

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Short Answer Questions
core concept

6. Describe the structure of an ionic lattice using sodium chloride (NaCl) as an example. In your answer, explain what holds the lattice together and why no discrete molecules exist in NaCl. 3 MARKS

✏️ Answer in your book

7. Compare the electrical conductivity of solid aluminium oxide (Al₂O₃, an ionic compound) and liquid aluminium (Al, a metal). Explain why both conduct as liquids but only one conducts as a solid, referring to the charge carriers in each case. 4 MARKS

✏️ Answer in your book

8. Magnesium oxide (MgO) is used as a refractory material, a substance that withstands very high temperatures without melting. Using your knowledge of ionic structure and lattice energy, explain why MgO is well-suited to this application. In your answer, compare MgO to NaCl. 4 MARKS

✏️ Answer in your book

We just saw how to compare ionic compounds by charge and ion size. That raises a question: how do you write precise exam answers on ionic conductivity that distinguish it from metallic conductivity? This card answers it → name the charge carrier (ion vs delocalised electron) and state the condition required for it to move.

Solid ionic compounds do not conduct because ions are fixed in the lattice. Molten or dissolved ionic compounds conduct because ions become mobile and can carry charge. Metals conduct in all states via delocalised electrons, a key distinction. For exam answers on conductivity: always state the charge carrier and whether it is free to move.

Pause, copy the highlighted conductivity rule into your book before moving on.

Fill the blanks: complete this exam-style sentence about ionic conductivity.

fixed in the lattice melted or dissolved mobile carry charge

Ionic compounds do not conduct electricity as solids because the ions are ___. When the substance is ___ in water, the ions become ___ and can ___, so the substance conducts.

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Worked examples

Worked examples · reveal as you go

Worked example +5 XP on full reveal

Explain why magnesium oxide (MgO) has a melting point of 2852°C, much higher than that of sodium chloride (NaCl, 801°C). Both are ionic compounds with similar crystal structures.

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Identify the ions: NaCl has Na⁺ (+1) and Cl⁻ (−1); MgO has Mg²⁺ (+2) and O²⁻ (−2)
The key difference is the charge on the ions. Always compare charges first when comparing lattice energies.
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Apply Coulomb's Law: force ∝ (charge₁ × charge₂) / distance². For NaCl: force ∝ 1×1 = 1. For MgO: force ∝ 2×2 = 4
The electrostatic attraction in MgO is approximately 4× stronger than in NaCl, assuming similar ionic radii. This quadrupling comes from doubling both charges.
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Connect to melting point: MgO requires ~4× more energy to break apart compared to NaCl. This is reflected in the ~3.5× higher MP
Melting requires overcoming electrostatic forces. Higher charge → stronger attraction → higher lattice energy → more thermal energy needed → higher MP.
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Worked example +5 XP on full reveal

A student dissolves sodium chloride in water and measures excellent electrical conductivity. She then melts it at 801°C and again measures conductivity. Finally she places probes in solid NaCl and measures no conductivity. Explain all three results using the ionic model.

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Solid NaCl: Na⁺ and Cl⁻ ions are locked in fixed positions in the ionic lattice and cannot move through the solid
Electrical conductivity requires mobile charge carriers. Since no ions can move freely, the solid does not conduct, this is the most common exam question on this topic.
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Molten NaCl: At 801°C, the lattice collapses. Na⁺ and Cl⁻ ions are now free to move independently in the liquid
When a voltage is applied, cations (Na⁺) migrate toward the negative electrode and anions (Cl⁻) migrate toward the positive electrode. Ions themselves carry the charge through the liquid.
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Dissolved NaCl: Water molecules (polar) surround Na⁺ and Cl⁻ ions, pulling them off the lattice surface (hydration)
The hydrated ions are mobile in solution and carry charge just as in the molten state. The mechanism is identical, only the disrupting force is different (water instead of heat).
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Misconception to fix

Common errors · the 3 traps that cost marks

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Misconception to fix

Wrong: Ionic compounds conduct electricity in the solid state because they contain charged ions.

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Misconception to fix

Right: Ionic compounds only conduct electricity when molten or dissolved in water. In the solid state, the ions are locked in a fixed lattice and cannot move. Conductivity requires mobile charge carriers, which are only present when the lattice breaks down.

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The formula NaCl means one sodium atom bonds to one chlorine atom to form a discrete molecule

Look back at the worked examples for the most common slip, units, ratios or sign errors are the usual culprits.

Fix: NaCl is a giant ionic lattice where the formula gives the simplest ratio of ions (1:1), not a molecular formula. There are no discrete molecules in an ionic crystal, each ion is surrounded by oppositely charged ions in all directions.

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Drill, then revisit

Quick-fire practice · 5 reps +2 XP per reveal

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What is the coordination number in sodium chloride (NaCl)?

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Why does MgO have a much higher melting point than NaCl?

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Why are ionic compounds brittle rather than malleable?

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State the standard exam explanation for why solid ionic compounds do not conduct electricity.

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Between LiF and CsI, which has the higher melting point and why?

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Revisit your thinking

Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?

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Interactive Tool, Atomic Structure Builder Open fullscreen ↗
The Atomic Structure tool shows that the atomic number of an element equals…
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Multiple choice

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Multiple choice
+5 XP per correct · +25 XP all-correct

Pick your answer, then rate your confidence. That tells the system what to drill next.

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Short answer
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Q1. 6. Describe the structure of an ionic lattice using sodium chloride (NaCl) as an example. In your answer, explain what holds the lattice together and why no discrete molecules exist in NaCl.

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ApplyApply4 MARKS

Q2. 7. Compare the electrical conductivity of solid aluminium oxide (Al₂O₃, an ionic compound) and liquid aluminium (Al, a metal). Explain why both conduct as liquids but only one conducts as a solid, referring to the charge carriers in each case.

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ApplyApply4 MARKS

Q3. 8. Magnesium oxide (MgO) is used as a refractory material, a substance that withstands very high temperatures without melting. Using your knowledge of ionic structure and lattice energy, explain why MgO is well-suited to this application. In your answer, compare MgO to NaCl.

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📖 Comprehensive answers (click to reveal)

Activity 1

1. CaF₂ will have a higher MP. Ca²⁺ has a charge of +2 while K⁺ has +1; F⁻ and Br⁻ both have −1, but the Ca²⁺/F⁻ combination produces stronger electrostatic attraction (higher charge on Ca²⁺ → higher lattice energy). Additionally, F⁻ is smaller than Br⁻, meaning ions in CaF₂ are closer together, further increasing the attraction. Both factors (higher charge on cation + smaller anion) raise the lattice energy → higher MP for CaF₂.

2. When MgCl₂ melts, the ionic lattice breaks down and Mg²⁺ and Cl⁻ ions become free to move independently in the liquid. When a voltage is applied, Mg²⁺ ions (positive) migrate toward the negative electrode (cathode) and Cl⁻ ions (negative) migrate toward the positive electrode (anode). This movement of charged particles constitutes an electric current, hence excellent conductivity.

3. X: ionic compound, high MP, no solid conductivity but excellent molten conductivity, hard and brittle. Y: covalent molecular compound, low MP (180°C), no conductivity in any state, soft and waxy. Z: metallic element, very high MP (1538°C = iron), excellent conductivity as both solid and liquid, malleable.

Activity 2

A: Copper conducts as a solid because it has a sea of delocalised electrons that are free to move throughout the metallic lattice at all times, no lattice disruption is needed. NaCl does not conduct as a solid because its charge carriers (Na⁺ and Cl⁻ ions) are fixed in the ionic lattice and cannot move. Both conduct as liquids: liquid copper still has delocalised electrons; molten NaCl has freed ions that can now move.

B: Glucose dissolves in water as intact polar molecules (C₆H₁₂O₆), not as ions. There are no charged particles in the glucose solution, water molecules interact with the glucose molecules via hydrogen bonding, but no ionisation occurs. Since conductivity requires mobile charge carriers (ions or free electrons), and dissolved glucose has neither, it does not conduct electricity even in solution.

❓ Multiple Choice

Ions fixed in lattice = correct ionic explanation. A is wrong (there ARE ions; the issue is they can't move). B and D are factually wrong.

NaCl (±1) vs MgO (±2) = massive charge difference → massive MP difference. A, C, D all involve ±1 compounds with small size differences → small MP differences.

High MP + hard/brittle + no solid conductivity + conducts dissolved = all ionic hallmarks.

Higher charges → stronger Coulomb attraction → higher lattice energy → higher MP. The correct causal chain.

Water's polarity attracts and separates ions (hydration). Mobile hydrated ions carry charge. No electrons or neutral atoms are involved.

Short Answer Model Answers

Q6 (3 marks): In NaCl, Na⁺ and Cl⁻ ions are arranged in a regular, repeating 3D pattern, each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ is surrounded by 6 Na⁺ (1 mark). The lattice is held together by strong electrostatic forces (ionic bonds) between oppositely charged ions acting in all directions simultaneously (1 mark). No discrete molecules exist because each ion is attracted to all its nearest neighbours, not to one specific partner, the entire crystal is one giant extended structure in which the formula NaCl simply represents the simplest whole-number ratio of ions (1 mark).

Q7 (4 marks): Solid aluminium (Al) conducts electricity because it has a sea of delocalised electrons that are free to move throughout the metallic lattice, electrons are the charge carriers (1 mark). Solid Al₂O₃ does not conduct because Al³⁺ and O²⁻ ions are fixed in the rigid ionic lattice and cannot move, no mobile charge carriers are present (1 mark). Liquid aluminium conducts because the metallic structure is maintained in the molten state, delocalised electrons remain mobile (1 mark). Molten Al₂O₃ conducts because the lattice has broken down, Al³⁺ and O²⁻ ions are now free to move and carry charge; the charge carriers in this case are ions, not electrons (1 mark).

Q8 (4 marks): MgO has a very high melting point (2852°C) because Mg²⁺ and O²⁻ carry charges of ±2, producing very strong electrostatic attraction between ions and a very high lattice energy (1 mark). In comparison, NaCl (MP 801°C) has ions with charges of only ±1, the electrostatic attraction is roughly four times weaker (applying Coulomb's Law: force ∝ charge₁ × charge₂), so much less energy is needed to disrupt the NaCl lattice (1 mark). To use a substance as a refractory material, it must not melt at operating temperatures, MgO's 2852°C MP means it remains solid in furnaces, kilns, and industrial reactors that operate at temperatures far exceeding those where NaCl would have already melted (1 mark). Additionally, the strong lattice makes MgO chemically and thermally stable under extreme conditions, it does not readily decompose or react with other materials at high temperature (1 mark).