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Module 1 · L6 of 21 25 min ⚡ +50 XP in Learn · +25 to complete Year 11 · Module 1 · IQ4

Chemical Bonding Overview

Today's hook, Diamond and graphite are both pure carbon, nothing else. One cuts steel. The other is soft enough to write with. If the atoms are identical, the bonding cannot be.

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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

Recall, your gut answer first

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Diamond is one of the hardest known substances and does not conduct electricity. Graphite is soft and slippery and conducts electricity well. Both are made of pure carbon. How can two substances with the same element have such different properties?

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What you'll master, and the words for it

What you'll master

Know

Key facts

The five key properties used to classify bonding type: melting point, electrical conductivity, hardness/malleability, solubility in water, and thermal conductivity
Examples of each structural type: ionic (NaCl), covalent molecular (H₂O, CO₂), covalent network (diamond, SiO₂), metallic (Cu, Fe)

Understand

Concepts

How the type of bonding determines physical properties, strong bonds give high MP, mobile charge carriers give conductivity
How the diagnostic trio (MP, conductivity, hardness), used together as converging evidence rather than any single property, lets you distinguish the structural types

Can do

Skills

Classify an unknown substance using melting point, conductivity, and hardness data
Compare and explain the properties of ionic, covalent molecular, covalent network, and metallic substances

The Five Key Properties for IQ2 Classification

core concept

In IQ1, physical properties told you pure vs mixture. In IQ2, the same properties, measured more carefully, reveal which type of bonding holds the substance together. The five properties are: melting point, electrical conductivity (solid and molten/dissolved), solubility in water, hardness/malleability, and thermal conductivity.

Structural type	MP/BP	Conductivity (solid)	Conductivity (molten/aq)	Hardness/Malleability	Solubility in water
Ionic compound	High (hundreds–thousands °C)	None (ions fixed in lattice)	Excellent (mobile ions)	Hard, brittle	Often soluble
Covalent molecular	Low (often <200°C)	None	None	Soft, crumbly	Variable (polar dissolves in water; non-polar doesn't)
Covalent network/lattice	Very high (>1000°C typically)	None (except graphite)	None (except graphite)	Extremely hard (except graphite)	Insoluble
Metallic (element)	Variable (Hg −39°C to W 3422°C)	Excellent	Excellent	Malleable, ductile	Generally insoluble (or reacts)

The diagnostic trio: To classify an unknown substance, focus on three questions: (1) What is the MP, low, high, or very high? (2) Does it conduct as a solid? As a liquid/solution? (3) Is it hard/brittle or malleable? Together, these three usually give a unique answer.

The five diagnostic properties are: melting point (MP), electrical conductivity, hardness/malleability, solubility in water, and thermal conductivity. Ionic: high MP, conduct when molten/dissolved, hard and brittle, often soluble. Covalent molecular: low MP (<200°C), no conductivity, soft. Covalent network: very high MP (>1000°C), extremely hard, insoluble. Metallic: high MP, generally conductive.

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

Lock-in task: A solid has a high melting point, does not conduct as a solid, but conducts well when melted, and is hard but shatters when struck. In one or two sentences, classify its structural type and justify using the diagnostic trio (MP, conductivity, hardness).

Bonding Diagram

core concept

Particle models of the four bonding types: ionic lattice (NaCl, MgO — high melting point, conducts only when molten or dissolved), covalent molecular (H2O, CO2 — low melting point, non-conducting), covalent network (diamond, SiO2 — very high melting point, non-conducting except graphite), and metallic (Cu, Fe, Mg — variable melting point, conducts as a solid via delocalised electrons).

We just saw the five diagnostic properties for classifying bonding types. That raises a question: what underlying structural difference explains why each bonding type shows those property patterns? This card answers it → each bonding type has a distinct particle arrangement that directly determines its properties.

Ionic lattice: alternating cations and anions in a giant 3D array held by electrostatic attraction. Covalent molecular: discrete molecules with strong intramolecular bonds but weak IMFs between molecules. Covalent network: continuous 3D network of covalent bonds. Metallic: positive cations in a sea of delocalised electrons. Melting point reflects the strength of forces that must be overcome to separate particles.

Add the highlighted structure summary to your notes before the check below.

Did you get this? True or false: covalent molecular substances like H₂O and CO₂ have low melting points because the bonds inside the molecule are weak.

Why These Properties Arise, Structure to Property

core concept

Melting Point: Energy to Overcome Bonds

Melting point reflects the strength of attractive forces that must be overcome to separate particles. Ionic bonds and covalent network bonds are very strong → very high MPs. Intermolecular forces in covalent molecular substances are weak (dispersion forces, dipole-dipole, hydrogen bonding) → low MPs. Metallic bond strength varies with metal type and structure.

Electrical Conductivity: Mobile Charge Carriers

Conduction requires particles that can move and carry charge. Metals have a sea of delocalised electrons that are free to move, so metals generally conduct in both solid and molten states. Ionic compounds have ions, but they are only mobile when the lattice is broken (molten or dissolved) → conduct then, not as solid. Most covalent substances have no free electrons or ions, so they generally do not conduct (graphite is a notable exception).

Graphite, the famous exception: Graphite is a covalent network solid where carbon atoms form layers. Within each layer, each C has one delocalised electron that can move freely → graphite conducts electricity. Between layers, only weak dispersion forces → layers slide easily → graphite is soft and slippery despite being a covalent network solid. Diamond, by contrast, has all four bonds fully covalent, no delocalised electrons, no conductivity, very hard.

Hardness vs Malleability

Ionic solids are hard because the lattice is rigid, but brittle when force is applied, like-charged ions align and repel, shattering the lattice. Metals are malleable because metal atom layers can slide past each other while the electron sea maintains bonding. Covalent network solids are hard because every bond must be broken to deform the structure.

We just saw the structural basis for each bonding type. That raises a question: the property that examiners test most is conductivity, why do some substances conduct and others not? This card answers it → conductivity requires mobile charge carriers, and only certain bonding types provide them.

Conductivity requires mobile charge carriers: metals generally conduct via delocalised electrons (in solid and molten states); ionic compounds conduct only when molten or dissolved (ions become mobile); most covalent substances do not conduct. Graphite is the exception, delocalised π-electrons within layers give in-plane conductivity. Diamond is the hardest natural substance due to its 3D covalent network.

Pause, write the highlighted conductivity rule into your book.

Four structure-to-property chains: ionic ions are fixed until molten or dissolved; molecular substances melt by overcoming intermolecular forces; network solids require many covalent bonds to break; and metals conduct through mobile delocalised electrons while ion layers can slide.

Inspect the middle cause in each panel: identify what can move for conductivity, then identify the attraction that must be overcome for melting or deformation.

Match each structural feature on the left to the property it explains on the right.

Delocalised electrons in a metal
Weak intermolecular forces in covalent molecular
Rigid lattice of oppositely-charged ions
Continuous 3D network of covalent bonds

Very high melting point and extreme hardness in diamond and SiO₂.
Low melting points in substances such as H₂O and CO₂.
Excellent conduction of electricity in both solid and molten metals.
Hard but brittle solids that shatter when like-charges align after a force is applied.

Short Answer Questions

core concept

6. Describe the relationship between electrical conductivity and structural type for each of the following: ionic compound, metallic element, and covalent molecular compound. In each case, explain why the substance does or does not conduct electricity. 3 MARKS

✏️ Answer in your book

7. A student is given three unknown solid substances (X, Y, Z) and measures: X, MP 801°C, no solid conductivity, conducts when dissolved; Y, MP 1085°C, excellent solid conductivity, malleable; Z, MP −115°C, no conductivity in any state, soft. Classify each structural type and justify. 4 MARKS

✏️ Answer in your book

8. Using your knowledge of structure and bonding, explain why diamond is extremely hard while graphite is soft enough to be used as a pencil lead, even though both consist entirely of carbon atoms. 4 MARKS

✏️ Answer in your book

We just saw the conductivity rules for each bonding type. That raises a question: how do you write a full-mark structure-to-property exam answer? This card answers it → always chain: bonding type → structural feature → property outcome.

For exam answers linking structure to property: (1) identify the bonding type; (2) describe the relevant structural feature (e.g. delocalised electrons, mobile ions, IMF strength); (3) explain the property outcome. For conductivity specifically: state whether mobile charge carriers are present and why. Graphite (conducts) and diamond (does not) demonstrate that the same element can have opposite properties depending on structure.

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

Lock-in task: In one or two sentences, explain why graphite conducts electricity but diamond does not, even though both are made entirely of carbon atoms.

Worked examples

Worked examples · reveal as you go

Worked example +5 XP on full reveal

An unknown substance has the following properties: MP = 1713°C (sharp), does not conduct electricity as a solid or when melted, insoluble in water, extremely hard. Classify the structural type and identify a likely substance.

Check the melting point: MP = 1713°C, very high

This rules out covalent molecular (low MP) and most metals. Ionic or covalent network are both possible at this stage.

Check conductivity: Does not conduct as solid or molten liquid

This rules out metals (which generally conduct as solids) and ionic compounds (conduct when molten). Only covalent network solids fit.

Confirm with hardness and solubility: Extremely hard + insoluble in water

This is consistent with covalent network solid. With very high MP, no conductivity in any state, extreme hardness and insolubility all agreeing, the converging evidence now points clearly to one type (no single property would have been enough on its own).

Identify likely substance: MP 1713°C, covalent network, extremely hard, insoluble

Silicon dioxide (SiO₂, quartz). Diamond melts at ~3550°C; SiO₂ at 1713°C matches exactly.

Worked example +5 XP on full reveal

Explain why sodium chloride (NaCl) is hard and brittle, while copper (Cu) is malleable and ductile, even though both have high melting points.

Identify the bonding in each: NaCl = ionic; Cu = metallic with delocalised electrons

NaCl has Na⁺ and Cl⁻ ions in a rigid lattice. Copper has positive copper ion cores (no fixed oxidation state) in a delocalised electron sea (non-directional bonding).

Explain why NaCl is brittle: When force is applied, ion layers shift and like-charged ions repel

Na⁺ aligns with Na⁺, Cl⁻ with Cl⁻ → strong electrostatic repulsion → lattice fractures along cleavage planes.

Explain why Cu is malleable: When force is applied, ion layers slide past each other

The electron sea moves with the layers, maintaining bonding throughout deformation. No directional bond breaks → metal deforms without fracturing.

Models, not rigid boxes. Ionic, covalent and metallic are the three school models you use to explain and predict properties, but real bonding character is a continuum: many substances sit between the ideal types (a polar covalent bond, for instance, has partial ionic character, and some compounds are best described as intermediate). Choose the model that best explains the evidence in front of you, and treat the category boundaries as useful approximations rather than hard walls.

Misconception to fix

Common errors · the 3 traps that cost marks

Misconception to fix

Wrong: Hardness and strength mean the same thing when describing materials.

Misconception to fix

Right: Hardness refers to resistance to scratching or indentation. Strength refers to resistance to breaking under force. Diamond is extremely hard but brittle, it can be scratched by almost nothing yet shatters under impact. These are independent properties.

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

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

Fix: Ionic compounds only conduct when molten or dissolved in water. In the solid state, the ions are locked in fixed positions in the lattice and cannot move to carry charge, conductivity requires mobile charge carriers.

Drill, then revisit

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

What are the three diagnostic properties used to classify an unknown substance?

Why do covalent molecular substances like H₂O and CO₂ have low melting points?

Why does sodium chloride conduct electricity when dissolved in water but not as a solid?

What is the key structural difference between diamond and graphite that explains their different properties?

Why are metals malleable and ductile?

Working · 1

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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 bonding lesson, use it only if you need to refresh how atoms are built before looking at how they bond.

The Atomic Structure tool shows that the atomic number of an element equals…

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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.

Short answer

ApplyApply3 MARKS

Q1. 6. Describe the relationship between electrical conductivity and structural type for each of the following: ionic compound, metallic element, and covalent molecular compound. In each case, explain why the substance does or does not conduct electricity.

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ApplyApply4 MARKS

Q2. 7. A student is given three unknown solid substances (X, Y, Z) and measures: X, MP 801°C, no solid conductivity, conducts when dissolved; Y, MP 1085°C, excellent solid conductivity, malleable; Z, MP −115°C, no conductivity in any state, soft. Classify each structural type and justify.

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ApplyApply4 MARKS

Q3. 8. Using your knowledge of structure and bonding, explain why diamond is extremely hard while graphite is soft enough to be used as a pencil lead, even though both consist entirely of carbon atoms.

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📖 Comprehensive answers (click to reveal)

Activity 1

1. Metallic element (sodium, Na). MP 98°C is within the metallic range; excellent conductivity as both solid and liquid indicates free delocalised electrons (metallic bonding); malleability and ductility are exclusive to metals.

2. Covalent network solid (diamond, C). MP 3550°C is extremely high, only covalent network solids reach this level. No conductivity in any state confirms absence of free electrons or mobile ions. Extreme hardness is characteristic of a continuous covalent bond network.

3. Covalent molecular compound (dry ice/CO₂ or similar). MP −78°C is very low, characteristic of weak intermolecular forces in a covalent molecular substance. No conductivity in any state confirms no free electrons or ions. Softness is consistent with weak van der Waals forces between discrete molecules.

Activity 2

Response 1, Error: Student A's answer is incomplete, NaCl does not conduct as a solid, but the claim that it "does not conduct electricity" is incorrect as a general statement. Correct answer: NaCl does not conduct as a solid (ions are fixed in the lattice and cannot move). However, when melted or dissolved in water, the ions become mobile and NaCl conducts electricity well. The reason for non-conduction as a solid is immobile ions, not the absence of free electrons, that reasoning applies to covalent substances, not ionic ones.

Response 2, Error: Student B incorrectly applied the general rules for covalent network solids to graphite. Graphite is a special case: within each carbon layer, one electron per carbon atom is delocalised and free to move → graphite conducts electricity. Between layers, only weak dispersion forces act → layers slide easily → graphite is soft. Graphite is used as an electrode precisely because it conducts; as a lubricant because its layers slide. Student B should have noted graphite as an exception to the general covalent network rules.

Response 3, Error: Student C is incorrect to say "definitely ionic". Both ionic compounds AND covalent network solids can have very high MPs and be very hard. The key distinguishing property is conductivity: ionic compounds conduct when molten or dissolved; covalent network solids do not. Without conductivity data, the classification cannot be definitive, it is "ionic compound or covalent network solid".

❓ Multiple Choice

Metallic: variable MP + conducts as solid and liquid + malleable/ductile. A = ionic (but "dissolved" is wrong word for metals). B = ionic compound. D = covalent molecular.

No solid conductivity but conducts dissolved = ionic compound. Metals generally conduct as solids; most covalent molecular substances do not conduct; covalent networks generally do not conduct (graphite is the exception).

Layer shift in ionic → like charges align → repulsion → fracture. Metal layers slide with electron sea intact → no fracture → malleable.

S: MP 2030°C (very high) + no conductivity in solid or molten state + extremely hard = covalent network. P = covalent molecular. Q = ionic. R = metal.

Graphite has delocalised electrons within layers (→ conducts) and weak interlayer forces (→ soft). It is a covalent network solid, not a metal; it doesn't dissolve in water; its melting point is very high.

Short Answer Model Answers

Q6 (3 marks): Ionic compound: does not conduct as a solid (ions are fixed in the rigid lattice and cannot move to carry charge), but conducts when molten or dissolved (ions become mobile and free to carry charge) (1 mark). Metallic element: conducts in all states because delocalised electrons are free to move throughout the metal structure, carrying charge in the solid and molten states (1 mark). Covalent molecular compound: does not conduct in any state because there are no free electrons and no ions, all electrons are localised in covalent bonds between specific atoms (1 mark).

Q7 (4 marks): X is an ionic compound, MP 801°C is high but not extreme; no conductivity as a solid indicates fixed ions in a lattice; conducts when dissolved confirms ionic character (mobile ions in solution) (1 mark + 1 justification). Y is a metallic element, MP 1085°C (consistent with copper); excellent conductivity as a solid indicates delocalised electrons; malleability is exclusively metallic (1 mark + 1 justification). Z is a covalent molecular compound, MP −115°C is very low, indicating only weak intermolecular forces between discrete molecules; no conductivity in any state confirms no free electrons or ions; softness is consistent with weak van der Waals forces (1 mark). [Note: 4 marks allocated across the three classifications and their justifications]

Q8 (4 marks): In diamond, each carbon atom forms four strong covalent bonds to four other carbon atoms in a continuous 3D tetrahedral network (1 mark). Every electron is localised in a covalent bond, there are no free electrons or weak points in the structure. To scratch or deform diamond, covalent bonds must be broken; the energy required is extremely high, making diamond the hardest natural substance (1 mark). In graphite, each carbon atom forms three covalent bonds within a flat hexagonal layer, with one remaining electron delocalised within the layer (1 mark). Between layers, only weak dispersion (van der Waals) forces act. These weak interlayer forces are easily overcome, layers slide past each other under small forces. The layers detach and transfer to surfaces (paper), which is why graphite writes. The low force needed to slide layers apart is what makes graphite soft (1 mark).

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