The big idea: Most covalent substances are small molecules (like CO2 or H2O). But in a few solids the covalent bonds keep going in every direction, joining millions of atoms into one giant lattice — a covalent network (also called a giant covalent structure).
Because there are no separate small molecules, the only way to melt or break the solid is to break strong covalent bonds. That makes these solids extremely strong, with very high melting points.
A giant covalent solid is one continuous, ordered network — every atom is joined to its neighbours by strong covalent bonds, so there are no separate small molecules.
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Key terms: - Giant covalent (network) solid — a continuous 3-D lattice of atoms held together by covalent bonds (no small molecules). - Allotrope — different structural forms of the same element (diamond and graphite are both pure carbon). - Delocalised electron — an outer electron not fixed to one atom, free to move (this is what lets graphite conduct). - Layer — a flat sheet of bonded atoms; layers can be held to each other by weak forces.
The four to know: The IB names four giant covalent solids: diamond and graphite (two allotropes of carbon), silicon (Si), and silicon dioxide (SiO_{2}). All four have very high melting points; their other properties depend on how the atoms are joined up.
In diamond, every carbon atom forms four covalent bonds, pointing to the corners of a tetrahedron. This repeats endlessly to give a rigid 3-D framework. Silicon has the same structure, and silicon dioxide is similar — each silicon is bonded to four oxygens and each oxygen bridges two silicons.
Why diamond is hard and non-conducting: - Very hard: the rigid 3-D network of strong covalent bonds cannot be pushed out of shape. - Very high melting point: melting needs many strong covalent bonds to be broken — this takes a lot of energy. - Does not conduct electricity: all four of each carbon's outer electrons are used in bonds, so there are no free (delocalised) electrons to carry charge.
Same idea, different element: Silicon and silicon dioxide (the main component of sand and quartz) follow the same logic: a giant covalent lattice → very high melting point and hardness, and (like diamond) silicon dioxide does not conduct.
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Graphite is the other allotrope of carbon. Here each carbon forms only three covalent bonds, making flat layers of joined hexagons. The fourth outer electron of every carbon becomes delocalised — free to move along the layer. This one structural difference explains all of graphite's unusual properties.
Diamond
- Each carbon is bonded to 4 others in a rigid 3-D tetrahedral network.
- Very hard — the rigid 3-D framework of strong bonds resists deformation.
- Does not conduct electricity — all 4 outer electrons are in bonds, none are free to move.
- Very high melting point — melting needs many strong covalent bonds broken.
Graphite
- Each carbon is bonded to 3 others in flat layers; the 4th electron is delocalised.
- Soft / slippery — weak forces between layers let the layers slide over each other.
- Conducts electricity — the delocalised electrons move along the layers.
- Very high melting point — the strong covalent bonds within the layers are hard to break.
Why graphite is soft AND conducts: - Soft / a good lubricant: the layers are held to each other only by weak forces, so they slide over one another easily. - Conducts electricity: the delocalised electrons between the layers are free to move and carry charge. - Still a very high melting point: the covalent bonds within each layer are strong, so melting graphite still needs a huge amount of energy.
Contrast: diamond uses all four electrons in bonds (no delocalised electrons → no conduction) and has no layers (rigid → hard).
Common trap: Graphite is soft because layers slide, not because its covalent bonds are weak. The covalent bonds inside a layer are just as strong as in diamond — which is why graphite's melting point is still very high.
How this is tested: S2.2.4 shows up two ways. Paper 1A likes a quick identify/match MCQ — recognising the structures of carbon allotropes (diamond, graphite and the related forms graphene and fullerene).
Paper 2 asks you to explain or discuss a property from the structure — most often a discuss question comparing a giant covalent solid's high melting point with the low melting point of a simple molecular substance, in terms of bonding (covalent bonds vs intermolecular forces).
Score the explain marks: Always name the type of attraction being broken. Giant covalent → strong covalent bonds broken (high m.p.). Simple molecular → only weak intermolecular forces broken (low m.p.) — the covalent bonds inside the molecules stay intact.
IB-style question — silicon vs sulfur
Silicon melts at about 1414 °C, but sulfur (S8) melts at only about 115 °C. Both contain covalent bonds. Explain this large difference in melting points in terms of structure and bonding. [3]
How to score the marks
- Mark 1 — the structures. Silicon is a giant covalent (network) solid, whereas sulfur (S8) is a simple molecular solid made of separate S8 molecules.
- Mark 2 — what is broken in silicon. Melting silicon requires breaking many strong covalent bonds throughout the lattice, which needs a large amount of energy → very high melting point.
- Mark 3 — what is broken in sulfur. Melting sulfur only overcomes the weak intermolecular (London) forces between the S8 molecules; the covalent bonds inside each molecule are not broken → low melting point.
Final answer
Silicon is giant covalent, so melting breaks many strong covalent bonds (high m.p.); sulfur is simple molecular, so melting only overcomes weak intermolecular forces between S8 molecules (low m.p.).