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Detailed notes on Atoms, Elements and Compounds for Cambridge IGCSE Chemistry, covering key concepts, explanations, examples, and exam-focused revision points.
Diamond, graphite, silicon dioxide and (Extended) graphene. Huge networks of covalently bonded atoms — very high m.p., distinctive properties from the bonding pattern.
Mapped to the Cambridge IGCSE 0620 syllabus (2026-2028).
Each C bonded to 4 others tetrahedrally. Hardest natural substance.
Diamond structure. A giant covalent lattice. Each carbon atom is bonded to FOUR other carbon atoms by single covalent bonds, arranged tetrahedrally.
Properties.
Uses.
Why doesn't diamond conduct? Each carbon's outer electrons are all 'committed' to covalent bonds. No mobile charge carriers.
Each C bonded to 3 others in flat hexagonal layers. The 4th electron is delocalised → conducts.
Graphite structure. Each carbon atom is bonded to THREE other carbon atoms in a flat HEXAGONAL pattern, forming sheets (layers). The fourth outer electron of each carbon is delocalised — free to move within the layer.
The layers are held together by WEAK forces (intermolecular forces, sometimes called van der Waals).
Properties.
Uses.
Compare with diamond. Both are pure carbon (allotropes), but TOTALLY different properties because of bond geometry.
| Feature | Diamond | Graphite |
|---|---|---|
| Bonds per carbon atom | 4 | 3 |
| Structure | Rigid 3D tetrahedral network | Flat hexagonal layers, weak forces between layers |
| Hardness | Very hard | Soft and slippery — layers slide |
| Conducts electricity | No — all 4 outer electrons are in bonds | Yes — the 4th electron is delocalised |
| Density | 3.5 g/cm³ | 2.3 g/cm³ |
| Melting point | Very high (> 3500 °C) | Very high |
Worked qualitative. Why is graphite slippery but diamond not? Diamond's atoms are bonded in a 3D network — no layers to slide. Graphite has 2D layers held weakly together — easy to slide.
Like diamond but with alternating Si and O. Very hard, found as sand and quartz.
Structure. Giant covalent lattice. Each silicon atom is bonded to 4 oxygen atoms; each oxygen to 2 silicon atoms. Repeats in a regular 3D pattern (like diamond, but with two atom types).
Properties.
Found as. Sand, quartz, much of the Earth's crust.
Uses.
Why is glass amorphous, not crystalline? When silicon dioxide is melted and rapidly cooled, the atoms don't have time to arrange into a regular lattice. They freeze in a disordered state. That's why glass is transparent (no grain boundaries to scatter light) and brittle.
Three big bonding categories with very different properties.
| Property | Simple molecular | Giant covalent | Ionic |
|---|---|---|---|
| Bonding | Strong covalent within; weak between | Giant network of covalent bonds | Strong electrostatic in lattice |
| m.p. / b.p. | Low | Very high | High |
| Conducts solid? | No | Only graphite | No |
| Conducts molten/dissolved? | No | (Doesn't usually melt cleanly) | Yes |
| Solubility | Variable | Insoluble | Often water-soluble |
| Hardness | Soft / liquid / gas | Very hard (except graphite is soft) | Hard but brittle |
Choosing the bonding from properties.
Worked qualitative. Substance X has m.p. 1610°C, doesn't conduct in solid or molten state, insoluble. Likely classification: giant covalent.
Verbatim phrases and definitions Cambridge mark schemes credit.
Giant covalent structures appear most years on Paper 4 (5-7 marks): explain a property using bond geometry. Examiner reports flag students attributing diamond's properties to ionic bonds, and saying graphite conducts because of metallic bonds (it conducts because of delocalised electrons in the layers — same idea, different language).
Sources: Cambridge IGCSE Chemistry 0620 syllabus 2026-2028 (2.3); 0620/42 Oct/Nov 2024 — Q6 (diamond vs graphite); 0620 Examiner Reports 2022-2024. Last reviewed 2026-05-06.
Step-by-step solutions to past-paper-style questions on giant covalent structures, written exactly the way a tutor would explain them at the board.
Question
What is meant by a giant covalent (macromolecular) structure? Name three examples.
Step-by-step solution
Step 1
A giant covalent structure is a huge lattice of MANY atoms held together by STRONG covalent bonds throughout the whole structure.
Step 2
There are no separate small molecules — the bonding continues in all directions (hence 'giant' or 'macromolecular').
Step 3
Examples in the 0620 syllabus: diamond, graphite and silicon(IV) oxide (SiO2).
Answer
A giant lattice of atoms joined by strong covalent bonds throughout. Examples: diamond, graphite and silicon(IV) oxide.
Examiner tip
The word 'giant' (or 'macromolecular') and the phrase 'strong covalent bonds throughout' are the marking points.
Question
Diamond and graphite are both made only of carbon atoms. What word describes them, and why do they have such different properties?
Step-by-step solution
Step 1
They are ALLOTROPES — different physical forms of the same element (carbon).
Step 2
Their atoms are arranged differently: in diamond each carbon bonds to 4 others in a tetrahedral lattice; in graphite each carbon bonds to only 3 others in flat layers.
Step 3
Different arrangement and bonding → different properties (e.g. diamond is hard, graphite is soft and conducts).
Answer
They are allotropes of carbon. The atoms are bonded and arranged differently, so the properties differ.
Examiner tip
Don't confuse allotropes (same element, different structure) with isotopes (same element, different number of neutrons).
Question
Explain, in terms of structure and bonding, why diamond is very hard and has a very high melting point.
Step-by-step solution
Step 1
In diamond each carbon atom is covalently bonded to 4 other carbon atoms in a rigid tetrahedral arrangement.
Step 2
These strong covalent bonds extend throughout the whole giant lattice in all directions.
Step 3
To scratch/deform it OR to melt it, a very large number of these strong covalent bonds must be broken, which needs a great deal of energy.
Step 4
Therefore diamond is very hard and has a very high melting point — which is why it is used in cutting tools and drill tips.
Answer
Each carbon forms 4 strong covalent bonds in a rigid 3D lattice. Breaking the many strong bonds requires a lot of energy → very hard and very high melting point.
Examiner tip
Marks come from '4 covalent bonds', 'strong bonds throughout' and 'many bonds must be broken'. Naming a use (cutting tools) is often a separate mark.
Question
Explain why graphite conducts electricity but diamond does not.
Step-by-step solution
Step 1
In graphite each carbon atom forms only 3 covalent bonds, so each atom has one outer electron that is not used in bonding.
Step 2
These spare electrons become DELOCALISED — one delocalised electron per carbon atom — and are free to move through the structure.
Step 3
The delocalised electrons can carry charge, so graphite conducts electricity (used as electrodes).
Step 4
In diamond each carbon uses ALL 4 of its outer electrons in covalent bonds, so there are no free/delocalised electrons → diamond does not conduct.
Answer
Graphite has one delocalised electron per atom (only 3 bonds) that is free to move and carry charge; diamond uses all 4 outer electrons in bonds, so it has no free electrons and cannot conduct.
Examiner tip
Conduction in graphite is by DELOCALISED ELECTRONS, not ions. The contrast '3 bonds vs 4 bonds' is the key explanation.
Question
Graphite is used as a lubricant and in pencil 'lead'. Explain, in terms of its structure, why graphite is soft and slippery, even though the covalent bonds in it are strong.
Step-by-step solution
Step 1
Graphite is arranged in LAYERS, each layer made of carbon atoms joined in hexagonal rings by strong covalent bonds.
Step 2
Between the layers there are only WEAK forces (weak intermolecular/van der Waals forces) — there are no covalent bonds linking the layers.
Step 3
Because these forces are weak, the layers can SLIDE over one another easily when a force is applied.
Step 4
This sliding of layers makes graphite soft and slippery, so it works as a lubricant and rubs off onto paper from a pencil.
Answer
The strong covalent bonds are only WITHIN the layers; the layers are held by weak forces, so they slide over each other — making graphite soft and slippery.
Examiner tip
The classic error is 'the covalent bonds are weak'. They are strong — it is the forces BETWEEN layers that are weak.
Question
Describe the structure of silicon(IV) oxide (SiO2) and explain why its physical properties are similar to those of diamond.
Step-by-step solution
Step 1
Silicon(IV) oxide is a giant covalent structure in which each silicon atom is covalently bonded to 4 oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms.
Step 2
This produces a rigid 3D tetrahedral lattice of strong covalent bonds throughout — the same general arrangement as the carbon atoms in diamond.
Step 3
Because many strong covalent bonds must be broken, SiO2 is very hard and has a very high melting point, like diamond.
Si bonded to 4 O; each O bonded to 2 Si
Step 4
All outer electrons are used in bonding, so there are no free electrons → SiO2 does not conduct electricity, also like diamond. It is found as sand/quartz and used in glass and as a refractory lining.
Answer
SiO2 is a giant covalent lattice (each Si bonded to 4 O, each O to 2 Si). Like diamond it has many strong covalent bonds in a rigid 3D network → very hard, very high melting point, and a non-conductor.
Examiner tip
Examiners reward the link to diamond: 'giant covalent / tetrahedral lattice with strong bonds throughout' explains both the hardness and the high melting point.
High-scoring sample answers for giant covalent structures on the Cambridge IGCSE 0620 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.
Name one allotrope of carbon that has a giant covalent structure. (1 mark)
Model answer
Diamond (or graphite).
Why this scores
One mark for either diamond or graphite. Both are giant covalent allotropes of carbon.
State the number of covalent bonds each carbon atom forms in (i) diamond and (ii) graphite. (2 marks)
Model answer
(i) In diamond each carbon atom forms 4 covalent bonds. (ii) In graphite each carbon atom forms 3 covalent bonds.
Why this scores
One mark each. The 4-bond vs 3-bond difference is the root of nearly every property contrast between the two allotropes.
Explain, in terms of its structure and bonding, why diamond is very hard. (3 marks)
Model answer
In diamond each carbon atom is covalently bonded to four other carbon atoms in a rigid tetrahedral arrangement. These strong covalent bonds extend throughout the whole giant lattice in every direction. To deform or scratch the structure, a very large number of these strong covalent bonds must be broken, which requires a great deal of energy, so diamond is very hard.
Why this scores
Three marks: (1) each carbon bonded to 4 others; (2) strong covalent bonds throughout a giant/rigid lattice; (3) many strong bonds must be broken to deform it.
Explain why graphite conducts electricity but diamond does not. (4 marks)
Model answer
In graphite each carbon atom forms only three covalent bonds, so each atom has one outer electron that is not used in bonding. These spare electrons are delocalised (one per carbon atom) and are free to move through the structure, so they can carry charge and graphite conducts electricity. In diamond, however, each carbon atom uses all four of its outer electrons in covalent bonds, so there are no free (delocalised) electrons to carry charge, and diamond does not conduct.
Why this scores
Four marks: graphite carbon forms 3 bonds; one delocalised/free electron per atom able to move; diamond carbon forms 4 bonds; no free electrons in diamond. Must say electrons, not ions.
Graphite is used as a lubricant. Explain, in terms of its structure and bonding, why graphite is soft and slippery. (5 marks)
Model answer
Graphite has a giant covalent structure arranged in layers, and within each layer the carbon atoms are joined in hexagonal rings by strong covalent bonds. Between the layers there are only weak forces of attraction, and there are no covalent bonds linking one layer to the next. Because these forces between the layers are weak, the layers can slide over one another easily when a force is applied. This sliding of the layers is why graphite is soft and slippery, making it a good lubricant and allowing it to rub off onto paper from a pencil.
Why this scores
Five marks: layered structure; strong covalent bonds within layers (hexagonal rings); weak forces between layers; no covalent bonds between layers; layers slide → soft/slippery. The 'bonds are weak' error scores zero.
Silicon(IV) oxide (SiO2) is used to line furnaces. Describe its structure and explain why its hardness, melting point and electrical conductivity are similar to those of diamond. (6 marks)
Model answer
Silicon(IV) oxide has a giant covalent structure in which each silicon atom is covalently bonded to four oxygen atoms and each oxygen atom is bonded to two silicon atoms, forming a rigid three-dimensional tetrahedral lattice of strong covalent bonds throughout — the same general arrangement as the carbon atoms in diamond. Because there are many strong covalent bonds, a great deal of energy is needed to break them, so SiO2 is very hard and has a very high melting point, just like diamond. Since all the outer electrons are used in the covalent bonds, there are no free (delocalised) electrons to carry charge, so SiO2 does not conduct electricity, again like diamond. These properties make it suitable as a hard, heat-resistant (refractory) lining for furnaces.
Why this scores
Six marks across: giant covalent structure; Si bonded to 4 O; O bonded to 2 Si; many strong covalent bonds needing much energy to break (→ hard, high m.p.); all outer electrons in bonds / no free electrons (→ non-conductor); valid comparison to diamond throughout.
Definitions to memorise and the exact keywords mark schemes credit for giant covalent structures answers — sharpened from recent examiner reports for the 2026 0620 sitting.
A huge lattice of many atoms held together by strong covalent bonds throughout the whole structure, e.g. diamond, graphite and SiO2.
Different physical forms of the SAME element, e.g. diamond and graphite are allotropes of carbon (same element, different structure).
Electrons not bound to a single atom and free to move through a structure to carry charge — graphite has one delocalised electron per carbon atom.
The 3D shape in which a central atom is bonded to four others pointing to the corners of a tetrahedron — found in diamond (C–C) and silicon(IV) oxide (Si–O).
The traps other students keep falling into on giant covalent structures questions — taken from recent Cambridge IGCSE 0620 examiner reports and mark schemes — and how to avoid them.
0620/42 — recurring
Why it happens
Students confuse the weak forces between the layers with the bonds inside the layers.
How to avoid it
The covalent bonds WITHIN each layer are STRONG. It is the FORCES BETWEEN the layers that are weak, letting the layers slide.
Why it happens
Over-generalising from electrolysis, where ions carry charge.
How to avoid it
Graphite conducts via DELOCALISED ELECTRONS (one free electron per carbon atom), NOT ions.
Why it happens
Students assume all carbon allotropes behave the same.
How to avoid it
In diamond all 4 outer electrons of each carbon are used in covalent bonds, so there are NO free electrons → diamond does NOT conduct.
Why it happens
Both involve the idea of 'different forms' of an element.
How to avoid it
Allotropes: same element, different STRUCTURE (diamond vs graphite). Isotopes: same element, different number of NEUTRONS.
The things students keep getting wrong in this sub-topic, answered.