Usually covalent bonding stops after a few atoms and you get a molecule. Sometimes it doesn’t stop at all — the bonding just keeps going, atom after atom, through the entire crystal. That gives a giant covalent structure, and it behaves nothing like a normal molecular substance.
📚 What you need to know
A giant covalent structure (covalent network) is a huge lattice of atoms joined by strong covalent bonds throughout.
There are no molecules — the bonding never stops, so we use an empirical formula.
All have very high melting points because melting means breaking millions of covalent bonds.
Diamond: 4 bonds per carbon, tetrahedral, 109.5°, very hard, does not conduct.
Graphite & graphene: 3 bonds per carbon, 120°, the 4th electron is delocalised so they conduct.
Buckminsterfullerene (C60) is not giant — it is a molecule with a fixed formula.
What makes a structure “giant”
Think about the difference between CO2 and SiO2. Both are a group 14 element with oxygen, but carbon forms two double bonds and stops — you get a small molecule that is a gas at room temperature. Silicon instead forms four single bonds to four different oxygens, each of which bonds on to another silicon, and so on forever. The result is sand, which melts at about 1700 °C.
That is the whole idea. In a giant covalent structure, there is no edge to the bonding. To melt it you must break actual covalent bonds, and there are an enormous number of them.
Diamond
In diamond, every carbon atom uses all four of its outer electrons to make four single covalent bonds to four other carbons. The arrangement is tetrahedral with bond angles of 109.5°, repeated in three dimensions.
Every carbon in diamond is at the centre of its own tetrahedron. The dashed bonds show that the pattern keeps repeating in all directions.
Everything about diamond follows from that picture:
Extremely hard — the hardest natural substance. To scratch it you have to break covalent bonds, and they hold firm in every direction.
Very high melting point (over 3500 °C) for the same reason.
Does not conduct electricity — all four outer electrons are locked into bonds, so there is nothing free to move.
Insoluble in everything, because no solvent can supply enough energy to break the network.
Silicon and silicon(IV) oxide are built the same way. In silicon, each Si bonds to four other Si atoms. In SiO2, each silicon bonds to four oxygens and each oxygen bridges two silicons — which is where the 1:2 ratio in the empirical formula comes from. Both are hard, high-melting and non-conducting, exactly like diamond.
Graphite
Graphite is also pure carbon, but each atom bonds to only three others. Those three bonds spread out at 120° into flat sheets of linked hexagons. The sheets then stack on top of each other, held together only by weak forces.
That leaves each carbon with one outer electron spare. It isn’t in a bond, so it becomes delocalised and can move freely between the layers.
Strong hexagonal sheets, weak forces between them, and delocalised electrons (red) free to move in the gaps.
Two of graphite’s famous properties come straight out of that structure:
It conducts electricity — unusual for a non-metal — because of the delocalised electrons moving between the layers.
It is soft and slippery, which is why it works in pencils and as a lubricant. The layers themselves are strong, but the forces between layers are weak, so they slide over each other easily.
Notice the trick here. Graphite is soft and has a very high melting point, which sounds contradictory. It isn’t: sliding layers apart only needs the weak forces broken, but melting means breaking the strong covalent bonds inside the layers.
Graphene and buckminsterfullerene
These two are close relatives of graphite, and the IB likes to test whether you can tell them apart.
Graphene is a single layer of graphite — one atom thick. Each carbon still bonds to three others at 120° and still has a delocalised electron, so it conducts extremely well. It is the thinnest and strongest material known, and it extends in two dimensions only.
Buckminsterfullerene, C60, is a hollow ball of 60 carbons arranged as 20 hexagons and 12 pentagons — a football. Each carbon bonds to three others, and the spare electrons let it act as a semiconductor.
Graphene is one sheet of that same hexagonal pattern; C₆₀ is the pattern curled up into a closed ball of exactly 60 atoms.
💡 C60 is the odd one out
It has a fixed formula — exactly 60 carbon atoms, every time.
That makes it a molecule, not a giant structure, even though it is grouped with them.
So it has a low melting point: the balls are held to each other only by weak intermolecular forces.
Summary table
Diamond
Graphite
Graphene
C60
SiO2
Bonds per atom
4
3
3
3
4 (Si), 2 (O)
Bond angle
109.5°
120°
120°
~120°
109.5°
Melting point
Very high
Very high
Very high
Low
Very high
Conducts?
No
Yes
Yes, very well
Semiconductor
No
Hardness
Hardest known
Soft, slippery
Strong, flexible
Soft solid
Very hard
Giant?
Yes
Yes
Yes (2D)
No
Yes
WORKED EXAMPLE
Diamond and graphite are both made of only carbon atoms. Explain why graphite conducts electricity but diamond does not.
Diamond: each carbon makes 4 covalent bondsAll four outer electrons are used up in bonding.no free electrons → does not conductGraphite: each carbon makes only 3 covalent bondsThe 4th outer electron is not needed for bonding.it is delocalised and free to move → conducts
WORKED EXAMPLE
Explain why silicon(IV) oxide has a melting point of about 1700 °C while carbon dioxide is a gas at room temperature.
SiO₂ is a giant covalent structureMelting means breaking many strong covalent bonds throughout the lattice.a huge amount of energy is neededCO₂ is a simple molecular substanceThe covalent bonds stay intact; only the weak forces BETWEEN molecules have to be overcome.very little energy needed → gas at room temperature
⚠️ Common mix-up
“Graphite has weak bonds” — no. The bonds are strong; the forces between the layers are weak. Say which you mean.
Don’t write a formula like C4 or Si2O4. Giant structures take empirical formulas: C, Si, SiO2.
C60 is not giant. Fixed formula = molecule = low melting point.
Giant covalent is not the same as giant ionic. There are no ions here, so these never conduct when molten or dissolved.
Up next: Intermolecular Forces — the weak attractions between separate molecules, which turn out to explain almost every physical property of simple covalent substances.
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