Allotropes of Carbon: Diamond, Graphite and Amorphous Carbon

1A pencil and a diamond

The "lead" in your pencil and the stone in a diamond ring are made of the same atoms. Both are pure carbon, element number 6, with the electron arrangement 2,4 that you met in S1. Yet a pencil leaves a grey smear when you drag it over paper, while a diamond scratches glass. The black one conducts electricity; the clear one does not. How can one element give two solids that behave so differently?

The answer is not in the atoms themselves but in how they are joined together. This lesson follows the idea that runs through the whole unit: structure decides properties, and properties decide uses.

2Diamond: every electron locked in a bond

In diamond each carbon atom uses all four of its outer electrons to form four covalent bonds with four other carbon atoms. The four neighbours sit at the corners of a tetrahedron, with an angle of 109.5° between any two bonds (Figure 1.1, left). The pattern repeats in all three directions, so a diamond crystal is one enormous molecule: a giant covalent structure.

• Very hard: to scratch or break diamond you must break many strong carbon–carbon bonds at once.
• Very high melting point (above 3500 °C) for the same reason.
• Does not conduct electricity: every outer electron is held in a bond, so none is free to move.
• Dense: about $3.5\mathrm{g}\mathrm{c}{\mathrm{m}}^{-3}3.5\backslash ,\backslash mathrm\{g\backslash ,cm^\{-3\}\}$, because the atoms are packed closely in every direction.

Figure 1.1: The structures of diamond and graphite.

3Graphite: layers that slide

In graphite each carbon atom forms only three covalent bonds, to three neighbours in the same flat plane, at 120°. The atoms form sheets of joined hexagons, like a honeycomb (Figure 1.1, right). The fourth outer electron of every atom is not held in any one bond. These electrons are delocalised: they can move along the whole layer.

• Conducts electricity along the layers, carried by the delocalised electrons. Graphite is unusual: a non-metal that conducts electricity well.
• Soft and slippery: the bonds within a layer are strong, but the layers are held to each other only by weak forces, 0.335 nm apart. The layers slide over each other and flake off, which is why graphite marks paper.
• Very high melting point, like diamond, because melting would mean breaking the strong bonds inside the layers.
• Less dense than diamond, about $2.2\mathrm{g}\mathrm{c}{\mathrm{m}}^{-3}2.2\backslash ,\backslash mathrm\{g\backslash ,cm^\{-3\}\}$, because of the wide gaps between layers.

Table 1.1: Diamond and graphite compared

4Amorphous carbon: charcoal, coke and soot

Most of the carbon you see every day is neither diamond nor graphite. Amorphous carbon ("without shape") is made of tiny, badly ordered pieces of graphite-like layers. It forms when carbon-rich materials are heated with little or no air, so that water and other volatile substances are driven off and carbon is left behind.

• Charcoal is made by heating wood with too little air to let it burn. Many Rwandan kitchens cook on it. Charcoal that has been treated to make it extremely porous is called activated charcoal: one gram has a huge internal surface on which other molecules stick (adsorption). It removes colours, smells and some poisons in water filters, sugar refining and gas masks.
• Coke is made by heating coal without air. It is almost pure carbon and is used to extract iron from its ore (Lesson 2).
• Soot (lampblack or carbon black) forms when fuels burn in too little air. It is used as the black pigment in ink, paint and car tyres.

Two newer allotropes are made of the same hexagons as graphite: graphene is a single layer only one atom thick, and fullerenes such as ${\mathrm{C}}_{60}\backslash mathrm\{C\_\{60\}\}$ are hollow balls of carbon atoms. Both are being studied for faster electronics and stronger materials.