Overview
What you'll learn
Form the three bonds. Describe how ionic bonds transfer electrons, covalent bonds share them, and metallic bonds pool them — all in the same chase for a full outer shell.
Link structure to properties. Relate the giant ionic lattice, simple molecular, giant covalent and metallic structures to their melting points, conductivity and hardness.
Tell diamond from graphite. Explain why two forms of pure carbon behave so differently, purely because of how the atoms are arranged.
Classify materials. Distinguish elements, compounds and mixtures, and explain why an alloy is harder and stronger than the pure metal.
Tutor's Insight
"Get this and every property makes sense."
There are four subtopics you can't skip and three bond types you must know cold. Structure decides every physical property — that's the rule examiners test. Master the link from bonding → structure → properties (and back), and Topic 3 questions stop feeling like surprises. It's always tested across Paper 1 and Paper 2, and it pairs with atomic structure and chemical equations everywhere else in the paper.
Topic 3 · Chemical Bonding and Structure
Atoms bond to reach a stable arrangement
Ionic
Transfer electrons
Between a metal and a non-metal. Electrons move across — forming oppositely charged ions.
Covalent
Share electrons
Between non-metals. Atoms share pairs of electrons so each gets a full outer shell.
Metallic
Pool electrons
Between metal atoms. Cations sit in a shared 'sea' of delocalised electrons.
Every atom is chasing the same goal — a full, noble-gas outer shell. Three bond types, three routes to it.
3.1 Ionic Bonding
3.1 Ionic Bonding
What this subtopic asks of you
- Describe how ionic bonds form through the transfer of electrons between metals and non-metals.
- Relate the physical properties of ionic compounds to their giant lattice structure.
3.1 Ionic Bonding
Metal gives, non-metal takes
How it forms
Electron transfer, then attraction
A metal atom hands its valence electrons to a non-metal atom — both reach a noble-gas configuration. The oppositely charged ions attract, and that electrostatic attraction is the ionic bond. The ions pack into a giant ionic lattice.
What makes it stronger
Charge up, radius down
The bond is stronger when ions carry a higher charge, and when their radii are smaller — both pull harder. Stronger bonds → higher melting points.
e.g. MgO melts far higher than NaCl — its ions are 2+ and 2−.
e.g. MgO melts far higher than NaCl — its ions are 2+ and 2−.
3.1 Ionic Bonding
The lattice dictates the properties
Property 01
High melting & boiling points
Takes a lot of energy to break the strong electrostatic forces across the whole lattice.
Property 02
Conducts only when molten or dissolved
Ions are locked in place when solid. Once molten or dissolved they move freely and carry charge.
Property 03
Hard, and usually soluble in water
Ions are held firmly in fixed positions. Many ionic compounds dissolve in water.
All three properties trace back to one thing — a giant lattice of ions held by strong electrostatic forces.
Free Notes · O-Level Pure Chemistry
Read the full chapter
Covalent bonding, simple molecular vs giant covalent, diamond vs graphite, metallic bonding, elements vs compounds vs mixtures, alloys, the structure-decides-properties table and 4 worked exam questions.
3.2 Covalent Bonding
3.2 Covalent Bonding
What this subtopic asks of you
- Describe how covalent bonds form when non-metal atoms share pairs of electrons.
- Tell simple molecular from giant covalent structures — and relate each to its properties.
3.2 Covalent Bonding
Non-metals share to fill their shells
How it forms
A shared pair of electrons
Two non-metal atoms each donate electrons to a shared pair. Both end up with a full noble-gas outer shell. The covalent bond is the strong electrostatic attraction between the two nuclei and the shared pair. A dot-and-cross diagram tracks whose electrons are whose.
Where you find it
Elements and compounds alike
Cl₂ and O₂ are covalent elements; H₂O and CH₄ are covalent compounds. Any bond between non-metal atoms is covalent.
3.2 Covalent Bonding
Same bond, two very different structures
Simple Molecular
Small, separate molecules
Discrete molecules held together by weak intermolecular forces. Low m.p. and b.p. Usually don't conduct. Often volatile.
e.g. iodine, methane
e.g. iodine, methane
Giant Covalent
One endless network
A vast network of atoms joined by strong covalent bonds. Very high m.p. and b.p. Hard. Don't conduct — except graphite.
e.g. diamond, silicon dioxide
e.g. diamond, silicon dioxide
The covalent bonds are strong either way — what differs is whether you break weak forces between molecules, or bonds within a network.
3.2 Covalent Bonding · Giant Covalent
Same element — opposite personalities
Diamond
Rigid 3D network
Each carbon bonded to four others in a rigid tetrahedral network. Extremely hard, very high m.p. All four valence electrons are bonded — so it doesn't conduct.
Graphite
Slippery, conducting layers
Carbon in flat layers. Each atom bonded to three others. Weak forces between layers let them slide — soft, used as a lubricant. The fourth electron is delocalised — so graphite conducts.
Both are pure carbon — the only difference is how the atoms are arranged. Structure is everything.
3.3 Metallic Bonding
3.3 Metallic Bonding
What this subtopic asks of you
- Describe a metal as a lattice of positive ions in a 'sea' of delocalised electrons.
- Explain its properties: high melting point, conducts electricity, malleable, ductile.
3.3 Metallic Bonding
Metals pool their electrons
How it forms
Cations in a sea of electrons
A metal is a giant lattice of positive ions sitting in a 'sea' of delocalised valence electrons. The metallic bond is the strong electrostatic attraction between the cations and that shared sea.
Why metals behave like metals
One structure, three signatures
High melting point — strong attraction across the lattice. Conducts electricity (solid and molten) — the delocalised electrons carry charge. Malleable & ductile — layers of ions slide over one another without breaking the bond.
Exam Habit
Name the delocalised electrons — Whenever you explain why a metal conducts or why graphite conducts, the marks are for saying the electrons are free to move and carry charge. Skip that phrase and you lose the mark.
3.4 Structure & Properties of Materials
3.4 Structure & Properties of Materials
What this subtopic asks of you
- Distinguish elements, compounds and mixtures, and describe an alloy as a mixture of a metal with another element.
- Deduce structure and bonding from physical properties, and physical properties from structure and bonding.
3.4 Structure & Properties
Elements, compounds, mixtures — and alloys
Element
One kind of atom
A single type of atom — it can't be broken down into anything simpler.
Compound
Chemically combined
Different elements chemically joined in a fixed ratio. Fixed melting & boiling points.
Mixture
Physically combined
Substances physically together in any ratio — separable by physical methods.
Alloys — stronger by disruption
An alloy is a mixture of a metal with another element. Different-sized atoms disrupt the regular layers — they can't slide as easily. That's why an alloy is harder and stronger than the pure metal. e.g. brass, steel.
3.4 Structure & Properties
Structure decides everything
| Structure type | Melting point | Conducts electricity? | Example |
|---|---|---|---|
| Ionic — giant lattice | High | Only molten / aqueous | NaCl |
| Simple molecular | Low | No | iodine, CH₄ |
| Giant covalent | Very high | No (except graphite) | diamond, SiO₂ |
| Metallic — giant lattice | High | Yes — solid & molten | copper, iron |
The exam move: read off the properties, match the row, name the structure and bonding. Works in reverse too.
Practice
Exam-style questions
Question 01
The Question
Structured
Name that structure
Substance X has a high melting point. It does not conduct electricity as a solid, but it does conduct when molten. What type of structure and bonding does X have? Explain.
Hint: take the two clues one at a time.
Worked Answer
A giant ionic lattice — ionic bonding
- High melting point → a giant structure with strong forces throughout.
- Doesn't conduct as a solid, but does when molten → it contains ions, locked in place when solid but free to move when molten.
- Both clues point to a giant ionic lattice held by ionic bonding.
Question 02
The Question
MCQ
Why graphite conducts
Both diamond and graphite are giant covalent forms of carbon, yet only graphite conducts electricity. Why?
Worked Answer
B
- In diamond every carbon uses all four valence electrons in bonding to four other atoms, so there are no free electrons — it doesn't conduct.
- In graphite each carbon bonds to only three others, so the fourth valence electron becomes delocalised.
- These delocalised electrons are free to move and carry charge, so graphite conducts. Answer: B.
Question 03
The Question
Structured
Why MgO melts higher than NaCl
Both magnesium oxide (MgO) and sodium chloride (NaCl) are giant ionic lattices, yet MgO has a much higher melting point. Explain why.
Worked Answer
Higher ionic charge → stronger attraction
- In MgO the ions carry charges of 2+ (Mg²⁺) and 2− (O²⁻); in NaCl they carry only 1+ (Na⁺) and 1− (Cl⁻).
- The higher charges create a stronger electrostatic attraction between the ions in the lattice.
- More energy is needed to overcome these stronger forces, so MgO has the higher melting point.
Question 04
The Question
Structured
Why alloys are harder than pure metals
Brass is an alloy of copper and zinc, and is harder than pure copper. Using ideas about structure, explain why an alloy is harder and stronger than the pure metal.
Worked Answer
Different-sized atoms disrupt the layers
- A pure metal is a lattice of ions arranged in regular layers that can slide over one another, which is why pure metals are soft and malleable.
- An alloy mixes in atoms of a different size (here, zinc among copper), which disrupt the regular layers.
- The layers can no longer slide over each other as easily, so the alloy is harder and stronger than the pure metal.
Frequently Asked Questions
Chemical Bonding and Structure — FAQ
What is Chemical Bonding and Structure in the O-Level Chemistry syllabus (6092)?
Chemical Bonding and Structure (Topic 3 of Syllabus 6092) covers four subtopics: 3.1 Ionic Bonding — electron transfer between metals and non-metals forming a giant ionic lattice; 3.2 Covalent Bonding — non-metals sharing electron pairs, and simple molecular versus giant covalent structures; 3.3 Metallic Bonding — a lattice of cations in a sea of delocalised electrons; and 3.4 Structure & Properties of Materials — how the four structure types, plus alloys, decide physical properties. Structure decides every physical property, and this topic is tested across Paper 1 and Paper 2.
How does an ionic bond form?
An ionic bond forms between a metal and a non-metal by the transfer of electrons. The metal atom hands its valence electrons to the non-metal atom, so both reach a stable noble-gas configuration. This produces oppositely charged ions — a positive cation and a negative anion — and the ionic bond is the strong electrostatic attraction between them. The ions then pack into a giant ionic lattice.
What is the difference between simple molecular and giant covalent structures?
Both use strong covalent bonds, but the arrangement differs. A simple molecular structure is made of small, discrete molecules held to each other by weak intermolecular forces, so it has low melting and boiling points and usually does not conduct (e.g. iodine, methane). A giant covalent structure is one endless network of atoms joined by strong covalent bonds throughout, so it has very high melting and boiling points and is hard (e.g. diamond, silicon dioxide). Melting a simple molecular solid only breaks the weak forces between molecules; melting a giant covalent solid means breaking strong bonds within the network.
Why do diamond and graphite have such different properties if both are carbon?
Diamond and graphite are both pure carbon giant covalent structures — the only difference is how the atoms are arranged. In diamond each carbon is bonded to four others in a rigid tetrahedral network, so it is extremely hard, has a very high melting point, and does not conduct because all four valence electrons are used in bonding. In graphite each carbon is bonded to only three others in flat layers held together by weak forces, so the layers slide (making it soft and a lubricant), and the fourth, delocalised electron on each atom lets graphite conduct electricity.
Why do metals conduct electricity and why are they malleable?
A metal is a giant lattice of positive ions in a 'sea' of delocalised electrons. It conducts electricity, both solid and molten, because those delocalised electrons are free to move and carry charge. It is malleable and ductile because layers of ions can slide over one another without breaking the metallic bond, since the sea of electrons simply moves with them. The strong attraction between the cations and the electron sea across the lattice also gives metals a high melting point.
Why is an alloy stronger than the pure metal?
An alloy is a mixture of a metal with another element. The added atoms are a different size, so they disrupt the regular layers of the metal lattice. Because the layers can no longer slide over each other as easily, the alloy is harder and stronger than the pure metal. Common examples are brass and steel.
O-Level Pure Chemistry · Syllabus 6092 · Topic 3 of 12 · © 2026 Overmugged. For personal study use only.