Diamond and graphite are both pure carbon, yet one cuts glass and the other slides off onto paper. The difference is how the atoms are bonded: in a giant structure every atom is linked to its neighbours, all the way through the solid.
In a giant covalent structure a huge number of atoms are joined by strong covalent bonds in a continuous network. In diamond each carbon atom is bonded to four others, so it is very hard, has a very high melting point and does not conduct. In graphite each carbon atom is bonded to three others in flat layers of hexagons; there are weak forces between the layers and one delocalised electron per atom, so graphite is soft, slippery and conducts. (Extended) Metallic bonding is the electrostatic attraction between a lattice of positive ions and a sea of delocalised electrons.
Fig. 1: bonding in diamond (left) and graphite (right)
Section 01
Diamond and graphite.
Switch between the two carbon structures in the simulator. Count the bonds on one atom in each, then look at what lies between graphite's layers.
Diamond
Graphite
bonds per C atom
4 (tetrahedral)
3, in flat layers of hexagons
between the atoms
strong covalent bonds in all directions
strong covalent bonds within layers, weak forces between layers
hardness
very hard
soft; layers slide over each other
melting point
very high
very high
electrical conductivity
does not conduct: no free electrons
conducts: one delocalised electron per atom
typical use
cutting tools, drill tips
lubricant, electrodes
Examiner trap
Graphite is soft because the forces between its layers are weak, not because its covalent bonds are weak. The bonds within each layer are very strong, which is why graphite still has a very high melting point. 'Graphite has weak covalent bonds' is a classic lost mark.
Stage 1 · Learn
Check what the sim just showed you
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In diamond, each carbon atom is covalently bonded to:
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Why can graphite be used as a lubricant?
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Why does graphite conduct electricity?
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Which use of diamond depends on its hardness?
Section 02
Silicon(IV) oxide.
Extended. Silicon(IV) oxide, SiO2, the main compound in sand, has a giant covalent structure similar to diamond.
Structure: each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom to two silicon atoms.
Formula: four bonds from each Si and two from each O means there are twice as many O atoms as Si atoms, so the ratio is 1 : 2, SiO2. There are no separate SiO2 molecules.
Properties: like diamond it is hard, has a very high melting point and does not conduct electricity, because it has no free electrons or ions.
Section 03
Metallic bonding.
Extended. In the simulator, apply a voltage and then push on the metal. Watch what the electrons do in each case, and notice that the ions can shift without the bonding breaking.
The bond: metal atoms lose their outer electrons into a shared sea of delocalised electrons. Metallic bonding is the electrostatic attraction between the lattice of positive ions and these delocalised electrons.
Good electrical conductors: the delocalised electrons are free to move through the structure and carry charge.
Malleable and ductile: the layers of positive ions can slide over each other while the sea of electrons still holds them together, so metals can be hammered into shape or drawn into wires.
High melting points: a lot of energy is needed to overcome the strong attraction between the ions and the delocalised electrons.
Examiner trap
Metals conduct because delocalised electrons move, not ions. In a solid metal the positive ions stay in the lattice. Mixing this up with ionic compounds, which conduct only when their ions are free to move, is a common error.
Stage 2 · Exam
Exam-style questions
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Why does diamond have a very high melting point?
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(Extended) Which describes the structure of silicon(IV) oxide?
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(Extended) Metallic bonding is best described as:
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(Extended) Why are metals malleable?
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