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Module 1 · L8 of 21 30 min ⚡ +50 XP in Learn · +25 to complete Year 11 · Module 1 · IQ4

Metallic Bonding and Properties

Today's hook, Hammer a piece of copper and it flattens. Hit a salt crystal and it shatters. Both are solids held together by electrical forces, so why do they break so differently?
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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
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Recall, your gut answer first
+5 XP warm-up

Copper is an excellent electrical conductor and can be drawn into thin wires without breaking. Diamond is extremely hard but cannot conduct electricity and will shatter if struck. Both are solids with atoms arranged in a regular lattice. Why do metals behave so differently from covalent network solids like diamond?

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What you'll master, and the words for it

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What you'll master
Know

Key facts

  • The electron sea (delocalised electron) model of metallic bonding
  • The physical properties of metals: conductivity, malleability, ductility, lustre, high MP
  • How alloy formation changes metallic properties
Understand

Concepts

  • Why the electron sea model explains every key metallic property
  • Why metal strength and MP vary across the periodic table
  • Why alloying disrupts the regular lattice and changes mechanical properties
Can do

Skills

  • Explain any metallic property using the electron sea model
  • Compare metals and alloys in terms of structure and properties
  • Predict how alloying would change the properties of a metal
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The Electron Sea Model

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The Electron Sea Model
core concept

What is metallic bonding?

When metal atoms come together, their valence electrons are released from individual atoms and become shared across the entire structure, they are delocalised. The result is a regular lattice of positive metal cations (now missing their valence electrons) immersed in a mobile "sea" of delocalised electrons. The metallic bond is the electrostatic attraction between this cation lattice and the electron sea.

Key insight: Unlike ionic bonds (between specific pairs of ions) or covalent bonds (between specific pairs of atoms), metallic bonding is non-directional there is no preferred direction for the bond. This non-directionality is the root cause of malleability and ductility.

Property-by-Property Explanation

Explanation using the electron sea model
Column B
Positive metal ions sit in a sea of delocalised electrons, explaining conductivity and malleability.

Metallic bonding = electrostatic attraction between a lattice of positive metal cations and a sea of delocalised valence electrons. The electrons are not tied to any atom and are free to move throughout the metal. The bond is non-directional, which is the root cause of malleability, ductility, and electrical and thermal conductivity.

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

Odd one out: which feature does NOT belong in a description of the electron sea model?

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Metallic Properties, Trends and Variations
core concept

Why do different metals have different properties?

Not all metals behave identically. The strength of metallic bonding, and therefore properties like MP and hardness, varies depending on:

  • Number of delocalised electrons per atom: More valence electrons released → stronger electron sea → stronger bonding. Group 1 metals (1 valence e⁻, e.g. Na, K) have weaker metallic bonding than Group 2 (2 valence e⁻, e.g. Mg, Ca) or transition metals (multiple d electrons, e.g. Fe, W).
  • Ionic charge: Higher charge on metal cation → stronger attraction to electron sea → stronger bonding → higher MP.
  • Ion size: Smaller ions → closer to electrons → stronger attraction.
MetalGroupApprox. MP (°C)Delocalised e⁻ per atomTrend explanation
Caesium (Cs)1291Very low MP, weak metallic bonding (1 delocalised e⁻, large ion)
Sodium (Na)1981Low MP, 1 delocalised e⁻, smaller than Cs
Magnesium (Mg)26502Moderate MP, 2 delocalised e⁻
Iron (Fe)Transition15382–3+High MP, multiple d electrons contribute to bonding
Tungsten (W)Transition34226+Highest MP of all metals, very strong metallic bonding

We just saw the electron-sea model and why metallic bonding is non-directional. That raises a question: why do different metals have vastly different melting points if they all have metallic bonding? This card answers it → bond strength depends on how many electrons are delocalised and how tightly the cation lattice holds them.

Metallic bond strength (and MP) depends on: number of delocalised electrons per atom, cation charge, and ion size. More valence electrons → stronger electron sea → higher MP (Group 1 weak, Group 2 stronger, transition metals strongest). Trend: Cs (29°C) < Na (98°C) < Mg (650°C) < Fe (1538°C) < W (3422°C).

Add the highlighted trend to your notes before the check below.

Two truths and a lie: spot the false statement about trends in metallic bonding strength.

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Alloys, Designed Metallic Materials

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Alloys, Designed Metallic Materials
core concept

What is an alloy and why make one?

A pure metal has a regular lattice of identical-sized cations. This regularity means layers slide easily, pure metals are often too soft or too corrosion-prone for engineering applications. An alloy introduces atoms of different sizes into the lattice, disrupting the regular arrangement and making it harder for layers to slide.

Why alloys are harder than pure metals: Foreign atoms (different size) create distortions in the lattice. When a force is applied, the distortions prevent smooth sliding of layers, more force is required to deform the alloy. Result: harder, stronger, less ductile than the pure metal.
AlloyBase metalAdded elementsKey property improvementApplication
SteelIron (Fe)Carbon (C, 0.2–2%)Harder, stronger than pure FeConstruction, tools
Stainless steelIronChromium (Cr, ~18%), NiCorrosion resistant, harderCutlery, surgical tools
BronzeCopper (Cu)Tin (Sn, ~10–12%)Harder, stronger than pure CuBearings, medals, instruments
BrassCopperZinc (Zn, 20–45%)Stronger, corrosion resistant, golden colourPipes, musical instruments
DuraluminAluminium (Al)Cu (~4%), Mg, MnMuch stronger than pure Al, low densityAircraft bodies
Foreign atoms distort the lattice of an alloy, blocking layer sliding so it is harder and stronger than the pure metal.

We just saw how the number of valence electrons and ion size control metallic bond strength. That raises a question: how do engineers make metals even harder and stronger than the pure metal alone? This card answers it → adding foreign atoms of a different size distorts the lattice and prevents layers from sliding.

An alloy is a mixture of a metal with one or more other elements, designed to improve properties. Foreign atoms of a different size distort the regular lattice, preventing layers from sliding → alloys are harder and stronger but less ductile than pure metals. Examples: steel (Fe + C), stainless steel (Fe + Cr + Ni), bronze (Cu + Sn), brass (Cu + Zn).

Pause, write the highlighted alloy definition into your book.

Fill the blanks: drag the right word into each gap.

different distort harder slide

An alloy contains atoms of ___ size to the host metal. These foreign atoms ___ the regular lattice, making it more difficult for layers of cations to ___ past one another. The result is an alloy that is ___ and stronger than the pure metal.

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Short Answer Questions

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Short Answer Questions
core concept

6. Using the electron sea model, explain why metals are good conductors of both electricity and heat. Clearly distinguish the mechanisms for each type of conductivity. 3 MARKS

✏️ Answer in your book

7. Explain why adding carbon atoms to iron produces steel that is harder and less malleable than pure iron. Refer specifically to the effect on the metallic lattice structure. 3 MARKS

✏️ Answer in your book

8. Tungsten (W, Group 6 transition metal, MP 3422°C) has one of the highest melting points of all metals, while caesium (Cs, Group 1, MP 29°C) has one of the lowest. Using the electron sea model, explain this large difference in melting points in terms of the metallic bonding in each metal. 4 MARKS

✏️ Answer in your book

We just saw how alloys work. That raises a question: what is the precise language examiners expect when you explain metallic conductivity, thermal conductivity, and alloy hardness? This card answers it → use specific terms (delocalised electrons drifting, lattice distortion) rather than vague descriptions.

For conductivity answers: delocalised electrons drift under a potential difference, never say "cations move". For thermal conductivity: delocalised electrons transfer kinetic energy rapidly. For trend questions: discuss number of valence electrons contributed, cation charge, and ion size. For alloys: state that foreign atoms disrupt the regular lattice.

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

Odd one out: which phrase does NOT belong in a model answer explaining metallic conduction of electricity?

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Worked examples

Worked examples · reveal as you go

Worked example +5 XP on full reveal

Compare the melting points of sodium (Na, MP 98°C, Group 1) and magnesium (Mg, MP 650°C, Group 2). Explain the difference using the electron sea model.

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Sodium (Na), Group 1
Releases 1 valence electron per atom
Cation: Na⁺ (charge = +1)
Group 1 metals have only 1 valence electron available to contribute to the electron sea.
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Magnesium (Mg), Group 2
Releases 2 valence electrons per atom
Cation: Mg²⁺ (charge = +2)
Group 2 metals have 2 valence electrons, creating a denser electron sea with higher cation charge.
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Electrostatic attraction in Na: weak (low charge Na⁺, sparse electron sea)
Electrostatic attraction in Mg: much stronger (higher charge Mg²⁺, denser electron sea)
Both factors, ionic charge and electron density, increase the strength of metallic bonding in Mg compared to Na.
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Na MP = 98°C (weak bonding, low lattice energy)
Mg MP = 650°C (strong bonding, high lattice energy)
The stronger metallic bonding in Mg requires much more thermal energy to break apart the lattice and allow atoms to move freely, resulting in a much higher melting point.
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Worked example +5 XP on full reveal

Pure copper (Cu) is malleable and relatively soft. Bronze (Cu + ~10% Sn) is significantly harder. Explain this difference using the electron sea model and the structure of alloys.

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Pure Copper (Cu) structure:
Regular lattice of identical copper ion cores (positive, no fixed charge)
Uniform electron sea throughout
Pure metals have a regular, orderly arrangement, all cations are the same size, so the lattice is completely uniform.
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What happens when Cu is hammered:
Layers of Cu ion cores shift past one another into new positions
The delocalised electron sea redistributes and keeps holding the cores together
Because the lattice is uniform and the bonding is non-directional, layers can slide into new positions while the electron sea keeps every core bonded, so the metal bends and deforms rather than shattering. (In an ionic lattice the same shift lines up like charges, they repel, and the crystal cleaves.) There is still real resistance, that is why metals have strength; the point is the bonding survives the deformation.
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Bronze (Cu + ~10% Sn) structure:
Mostly copper ion cores, but ~10% sites occupied by larger Sn atoms
Lattice is distorted where Sn atoms sit
Tin atoms are a different size from Cu atoms, creating local irregularities in the regular lattice, the uniform order is disrupted.
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What happens when Bronze is hammered:
Layer sliding is impeded at Sn sites
Distortions act as obstacles, layers cannot slide smoothly
Result: harder, stronger, less malleable
The distorted lattice blocks dislocation movement. More force is required to deform the metal because layers cannot slip easily past the foreign atoms, this is why alloys are harder than pure metals.
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Misconception to fix

Common errors · the 3 traps that cost marks

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Misconception to fix

Wrong: Alloys are harder than pure metals because they contain more atoms.

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Misconception to fix

Right: Alloys are harder because foreign atoms disrupt the regular metallic lattice, preventing layers from sliding past each other easily. It is the disruption of regularity, not the quantity of atoms, that increases hardness and reduces malleability.

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Saying "metallic bonds break" when metals are hammered

Students often write that hammering a metal "breaks the metallic bonds", but if bonds broke, the metal would shatter like an ionic crystal. The point of the electron sea model is that the bond is non-directional: cations slide past each other and the electron sea re-forms attractions instantly.

Fix: Write that cation layers slide while the electron sea maintains the bonding, bonds re-form, they do not break.

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Drill, then revisit

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

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What two things does the electron sea model say a metal is made of?

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Why is copper ductile (can be drawn into wire)?

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Predict: which has the higher melting point, sodium or magnesium? Justify in one sentence.

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Name the alloy: iron + ~18% chromium + nickel. Give one application.

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Tungsten (MP 3422 °C) has the highest melting point of any metal. Explain in terms of the electron sea model.

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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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Optional prerequisite review, Atomic Structure Builder (Year 9) Open fullscreen ↗

This Year 9 atomic-structure tool is optional revision of a prerequisite idea (atoms, protons, electrons). It is not the core activity for this metallic-bonding lesson, use it only if you need to refresh how atoms are built before looking at how metal atoms bond.

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

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Short answer
ApplyApply3 MARKS

Q1. 6. Using the electron sea model, explain why metals are good conductors of both electricity and heat. Clearly distinguish the mechanisms for each type of conductivity.

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ApplyApply3 MARKS

Q2. 7. Explain why adding carbon atoms to iron produces steel that is harder and less malleable than pure iron. Refer specifically to the effect on the metallic lattice structure.

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ApplyApply4 MARKS

Q3. 8. Tungsten (W, Group 6 transition metal, MP 3422°C) has one of the highest melting points of all metals, while caesium (Cs, Group 1, MP 29°C) has one of the lowest. Using the electron sea model, explain this large difference in melting points in terms of the metallic bonding in each metal.

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📖 Comprehensive answers (click to reveal)

️ Activity 1

A: Calcium has the higher MP. K (Group 1) contributes 1 valence electron and forms K⁺ (charge +1). Ca (Group 2) contributes 2 valence electrons and forms Ca²⁺ (charge +2). Ca²⁺ has a higher ionic charge and contributes more electrons to the electron sea, both factors increase the electrostatic attraction between cations and the electron sea → stronger metallic bonding → higher MP (Ca: 842°C vs K: 63°C).

B: Pure Fe has a regular lattice of identical Fe cations, layers slide relatively smoothly, making it malleable (though less so than Group 1 metals due to stronger bonding). When 0.5% carbon is added, the smaller C atoms occupy interstitial spaces in the Fe lattice, distorting the regular arrangement. These distortions act as obstacles to layer sliding, more force is required to deform the steel. Result: steel is harder and less malleable than pure iron, which is why raw iron is rarely used in structural applications.

Activity 2

Novel Context 1: Tungsten (W) is far more suitable for turbine blades, its MP of 3422°C means it remains solid at operating temperatures of jet engines (~1500°C). Aluminium (MP 660°C) would melt immediately. W has such a dramatically higher MP because it is a transition metal that contributes ~6 valence electrons per atom into the electron sea (compared to Al's 3). Additionally, W⁶⁺ carries a much higher charge than Al³⁺. The combination of many delocalised electrons and high cation charge produces extremely strong metallic bonding requiring enormous energy (very high temperature) to overcome.

Novel Context 2: The chemist is correct, pure 24-carat gold is actually the weakest and softest form of gold jewellery, not the strongest. 18-carat gold (an alloy with 25% other metals such as Ag, Cu, or Pd) is significantly harder and stronger than pure gold. The added atoms have different sizes from Au and disrupt the regular gold lattice, creating distortions that prevent smooth layer sliding, more force is required to scratch or deform the alloy. Pure gold is so soft that rings made from it will deform under normal wear. The trade-off is reduced purity, but the alloy is far more practical for everyday jewellery.

❓ Multiple Choice

Non-directional bonding allows layer sliding while the electron sea maintains cohesion. A, C, D are all incorrect descriptions of metallic structure.

More delocalised electrons + higher charge → stronger bonding → higher MP. This is the correct general principle.

Different-sized Sn atoms distort the regular Cu lattice → impede layer sliding → harder. Not about electron numbers, ionic bonds, or covalent bonds.

Only the electron sea model simultaneously explains conductivity (mobile electrons), malleability (non-directional bonding, layer sliding), and high MP (strong cation–electron attraction).

Stainless steel combines alloying strength with Cr's passive oxide layer for corrosion resistance. Pure iron corrodes easily; pure Al is weak structurally; bronze is for bearings, not structural applications.

Short Answer Model Answers

Q6 (3 marks): Metals conduct electricity because their delocalised valence electrons are free to move throughout the lattice at all times. When a voltage (potential difference) is applied, electrons flow from the negative terminal toward the positive terminal, this directed electron movement constitutes an electric current (1 mark). Metals conduct heat because mobile delocalised electrons can absorb kinetic energy at the hot end of the metal and rapidly transfer this energy through collisions with other electrons and cations throughout the lattice, much faster than vibration-mediated heat transfer in non-metallic solids (1 mark). The mechanisms differ: electrical conductivity is directed electron movement driven by a voltage gradient; thermal conductivity is kinetic energy transfer by electrons moving randomly but carrying energy from hot to cool regions, one is electrical, the other is thermal (1 mark).

Q7 (3 marks): Pure iron has a regular lattice of Fe cations of uniform size, allowing layers to slide past each other relatively easily under applied force, the electron sea redistributes and maintains bonding as layers shift (1 mark). Carbon atoms are much smaller than Fe atoms. When added, they occupy interstitial spaces in the Fe lattice, distorting the regular cubic arrangement at those sites (1 mark). These distortions act as obstacles, when a shear force is applied, layers cannot slide smoothly past the sites where C atoms sit, because the C atom's size difference blocks dislocation movement. Greater force is required to deform the steel → harder; reduced ability to slide → less malleable (1 mark).

Q8 (4 marks): Caesium is a Group 1 metal, each Cs atom contributes only 1 valence electron to the electron sea, and forms a Cs⁺ cation with charge +1 (1 mark). The attraction between Cs⁺ (low charge, very large ion) and the sparse electron sea (1 electron per atom) is very weak → very low lattice energy → low MP of 29°C (1 mark). Tungsten is a Group 6 transition metal, each W atom can contribute up to ~6 valence electrons to the electron sea, and the cation carries a much higher effective charge (1 mark). The electrostatic attraction between the highly charged W cation and the very dense electron sea (~6× more electrons per atom than Cs) is enormously strong → very high lattice energy → highest MP of any metal at 3422°C. The 3393°C difference in melting point reflects this ~6× difference in the number of delocalised electrons and the dramatic difference in cation charge (1 mark).

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