Graphite and diamond as giant covalent allotropes of carbon

Count the bonds

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The amber dots: the fourth outer shell electron of each graphite carbon, delocalised and free to move through the layers.

Both sketches are simplified and not to scale: diamond is really three-dimensional, and both structures carry on far beyond what is drawn. The faint stubs show where.

Predict, then check

You have just met both structures. Commit to an answer before you look.

Diamond and graphite are both pure carbon. Which one conducts electricity, and why?

What you need to know

• Allotropes are different structural forms of the same element — diamond and graphite are both made only of carbon.
• A giant covalent structure is a large, regular arrangement of atoms, all joined together by covalent bonds.
• The number of atoms isn’t fixed, so there’s no particular molecular formula — diamond and graphite are both just C.
• Have a goMaya says: “This diamond has about a million carbon atoms and that one has a billion, so they need different formulas.” Is she right?

  No. Both are just diamond, represented by C.

  In a giant covalent structure the number of atoms depends on the size of the sample, so a bigger piece does not need a new formula.

• In diamond, each carbon atom bonds to four others, making a rigid three-dimensional tetrahedral network.
• In graphite, each carbon bonds to three others, forming flat layers of hexagonal rings with only weak forces between the layers.
• Have a goPush sideways on the top layer of graphite. What gives way first: the bonds inside the layer, or the forces between the layers?

  The weak forces between the layers.

  Only weak forces of attraction hold the layers together, so the layers slide, while each layer is held by covalent bonds.

• Each graphite carbon uses three outer shell electrons for bonding; the fourth is delocalised, free to move through the layers.
• Have a goA flake of graphite contains 100 carbon atoms. How many delocalised electrons does it have?

  100: one for each carbon atom.

  Every carbon atom uses three of its four outer shell electrons in bonding, which leaves one delocalised electron per atom.

• Both have very high melting points, because a large amount of energy is needed to break the many strong covalent bonds.
• Diamond is very hard (rigid lattice, strong covalent bonds); graphite is soft because its layers slide over each other.
• Graphite conducts because its delocalised electrons move and carry charge; diamond doesn’t, as all its outer shell electrons are in bonding.
• Coal is mostly carbon, but it’s a mixture of different molecules with no regular structure, so it isn’t giant covalent.

The big picture

Diamond and graphite are both pure carbon, but each carbon atom bonds to four others in diamond and three in graphite. That one difference decides how hard they are, which one conducts electricity and why graphite’s layers slide, while both keep a very high melting point.

⚠ Watch out

Thinking that because graphite is soft, its bonds must be weak and it must melt easily. The layers slide because only weak forces hold them together, but the covalent bonds inside each layer are strong, and graphite’s melting point is very high.

Memory hook

Four bonds: every electron locked in. Three bonds: one electron left free. Count the bonds, then count the electrons.

Check yourself

Cover the page and sketch it: one diamond carbon with its four bonds, then two graphite layers with a gap between them. Label what is strong, what is weak, and where the delocalised electrons are.