Sign up to access the complete lesson and track your progress!
Unlock This CourseDiamond and graphite are both made of only carbon atoms. Yet diamond is the hardest natural substance, and graphite is soft enough to leave a grey line on paper. The difference is not the atoms. It is how the atoms are bonded together.
Remember, both are giant covalent structures: huge numbers of atoms all joined by strong covalent bonds. In this lesson we look at exactly how the carbon atoms are arranged in each one, and use that to explain their properties.
A carbon atom has 4 electrons in its outer shell, so it can form up to four covalent bonds. Whether it uses all four, or only three, is the key to everything in this lesson.
A diamond's sparkle hides a rigid network where every carbon atom has four strong covalent bonds - that's why it's so hard
In diamond, each carbon atom forms four covalent bonds with other carbon atoms in a giant covalent structure. The bonds point out in three dimensions, so every atom is locked in place by four strong bonds, and the pattern repeats in every direction.
When you draw part of diamond on paper, you are showing a 3D structure in 2D. Look for one carbon in the middle joined to four others, with those four each joined to more carbons. Every carbon in the structure has exactly four bonds.
Every atom is held by four strong covalent bonds in a rigid 3D network. To scratch or dent diamond you would have to break many of these bonds. This is why diamond is used on the tips of drills and cutting tools.
To melt diamond, a huge number of strong covalent bonds must be broken. This needs a great deal of energy, so the melting point is very high.
All four outer electrons of each carbon are used in covalent bonds. There are no delocalised electrons and no ions, so there are no charged particles free to move.
Pencil lead is graphite - its layers of carbon slide apart easily, leaving that grey line on the paper
In graphite, each carbon atom forms three covalent bonds with three other carbon atoms. This makes flat layers of hexagonal rings, a bit like a sheet of chicken wire. There are no covalent bonds between the layers. The layers are held together only by weak forces.
Each carbon has four outer electrons but only uses three of them in bonds. The fourth electron is not stuck to one atom. One electron from each carbon atom is delocalised, which means it is free to move along the layer.
Key terms:
There are no covalent bonds between the layers, only weak forces. The layers can slide over each other easily. This is why graphite is soft and slippery. It is used in pencils, because layers rub off onto the paper, and as a lubricant.
Each carbon atom has one delocalised electron. These electrons can move through the structure and carry charge, so graphite conducts electricity. It also conducts thermal energy.
Inside each layer, every atom is held by three strong covalent bonds. To melt graphite these strong bonds must be broken, not just the weak forces between layers. So graphite also has a very high melting point.
Graphite is similar to metals because it has delocalised electrons. In both, these free electrons are what let the substance conduct electricity. This is unusual, because most non-metals do not conduct.
| Diamond | Graphite | |
|---|---|---|
| Bonds per carbon atom | 4 | 3 |
| Shape | Rigid 3D network | Flat layers of hexagonal rings |
| Delocalised electrons? | No | Yes, one per carbon atom |
| Hardness | Very hard | Soft and slippery |
| Melting point | Very high | Very high |
| Conducts electricity? | No | Yes |
Notice that both have very high melting points, because both are giant covalent structures. The differences in hardness and conduction come from the number of bonds each carbon makes.
Question: Explain why graphite conducts electricity but diamond does not.
Step 1 - bonding in each: In graphite each carbon forms three covalent bonds. In diamond each carbon forms four.
Step 2 - what happens to the spare electron: In graphite, one outer electron from each carbon is not used in bonding and is delocalised. In diamond, all four outer electrons are used in bonds, so none are delocalised.
Step 3 - link to the property: Delocalised electrons can move through graphite and carry charge, so graphite conducts. Diamond has no charged particles free to move, so it does not conduct.
A diagram shows carbon atoms in flat sheets of six-sided rings, stacked on top of each other, with no lines joining one sheet to the next.
Count the bonds on one atom: three. Six-sided rings in flat layers, three bonds per atom and no bonds between layers means this is graphite. If every atom had four bonds going off in different directions, it would be diamond.
Diamond and graphite are both forms of carbon.
(a) Explain why diamond is very hard. [2 marks]
(b) Explain why graphite is soft and slippery. [2 marks]
(c) Give one way graphite is similar to metals. [1 mark]
(a) Each carbon atom forms four covalent bonds with other carbon atoms (1). This makes a rigid giant covalent structure, so many strong bonds would have to be broken (1).
(b) Graphite is made of layers of hexagonal rings with no covalent bonds between the layers (1). So the layers can slide over each other (1).
(c) Graphite has delocalised electrons, like a metal (1).
For any "explain the properties" question on diamond or graphite, always start with the structure: how many bonds each carbon atom makes. Then link it to the property with "so" or "because". Structure, then property, every time.