Giant covalent structures
Definition
Giant covalent structure
A three-dimensional network of a very large number of atoms, each joined to its neighbours by strong covalent bonds.
- Some substances are not made of small, separate molecules but of a huge network of atoms joined together by covalent bonds.
- These substances have giant covalent structures, in which every atom is held in place by strong covalent bonds to its neighbours.
- There are no separate molecules in a giant covalent structure, so a whole crystal can be thought of as one enormous network of atoms.
- Diamond and graphite both have giant covalent structures made only of carbon atoms.
- The carbon atoms in diamond and graphite are bonded in different arrangements, and these different structures give the two substances very different properties.
- To melt either substance, a very large number of strong covalent bonds must be broken, so both diamond and graphite have very high melting points, above 3500 ∘C3500\ ^{\circ}\text{C}3500 ∘C.
The structure of diamond
- In diamond, every carbon atom is joined by covalent bonds to four other carbon atoms.
- The four bonds around each carbon atom point towards the corners of a tetrahedron, which builds up a rigid three-dimensional network.
- Carbon has four outer-shell electrons, and all four are used in the four covalent bonds, so diamond has no free electrons.
- Diamond is extremely hard, because many strong covalent bonds hold every atom firmly in place in all three dimensions.
- Diamond does not conduct electricity, because all its outer-shell electrons are held in covalent bonds and there are no electrons or ions free to move.
- Diamond is used in cutting tools, such as the tips of drill bits and the edges of saws for cutting stone, glass and concrete.
- Diamond is suited to this use because it is harder than the materials it cuts, so its strong network of covalent bonds is not worn away.
The structure of graphite
Definition
Delocalised electrons
Electrons that are not held by one particular atom or bond and are free to move through a structure.
- In graphite, every carbon atom is joined by covalent bonds to only three other carbon atoms.
- These bonds join the carbon atoms into flat layers made of hexagonal rings of six carbon atoms.
- The layers are stacked on top of one another and are held together only by weak forces between the layers.
- Each carbon atom uses three of its four outer-shell electrons in covalent bonds, leaving one electron that is not used in bonding.
- These spare electrons become delocalised, so they are free to move along the layers through the whole structure.
- Graphite is soft and slippery, because the weak forces between the layers let the layers slide over one another easily.
- Graphite conducts electricity, because its delocalised electrons can move along the layers and carry charge.
- Graphite still has a very high melting point, because the strong covalent bonds within the layers must be broken for it to melt.
Uses of graphite
Definition
Lubricant
A substance that reduces friction between surfaces that slide over one another.
Definition
Electrode
A solid electrical conductor through which an electric current enters or leaves a liquid or solution during electrolysis.
- Graphite is used as a lubricant, because the layers slide over one another easily and so reduce friction between moving parts.
- Powdered graphite is used to lubricate locks and machinery that would be damaged by oil or that run too hot for oil to be used.
- Graphite is used as an electrode in electrolysis, because its delocalised electrons allow it to conduct electricity.
- Its very high melting point also means graphite electrodes stay solid in hot molten compounds.
- Pencil leads are made of graphite mixed with clay, because the layers rub off onto paper and leave a grey mark.
Example
Hardness of diamond and graphite
- In diamond, each carbon atom is held by four strong covalent bonds in a rigid three-dimensional network.
- Breaking or scratching diamond would mean breaking many of these strong bonds, so diamond is extremely hard.
- In graphite, each carbon atom is held by three strong covalent bonds within its layer, but the layers are held to one another only by weak forces.
- Only the weak forces between the layers must be overcome for the layers to slide, so graphite is soft and slippery.
Example
Conductivity of diamond and graphite
- In diamond, all four outer-shell electrons of each carbon atom are used in covalent bonds, so no electrons are free to move and diamond does not conduct.
- In graphite, each carbon atom uses only three outer-shell electrons in bonds, so one electron per atom is delocalised.
- The delocalised electrons can move along the layers and carry charge, so graphite conducts electricity.
Common Mistake
- Do not say that graphite is soft because its covalent bonds are weak, because the bonds within the layers are strong and only the forces between the layers are weak.
- Do not say that graphite conducts because it contains ions, because the charge is carried by delocalised electrons.
- Do not describe diamond or graphite as molecules, because both have giant covalent structures.
Exam technique
Linking structure to use
- Build every answer as a chain that goes from structure to property to use.
- For diamond in cutting tools, mention four covalent bonds per carbon atom, a rigid structure and very hard.
- For graphite as a lubricant, mention layers that slide because of weak forces between the layers.
- For graphite as an electrode, name the delocalised electrons and say they can move to carry charge.
Self review
- How many carbon atoms is each carbon atom bonded to in diamond?
- How many carbon atoms is each carbon atom bonded to in graphite?
- Why can graphite conduct electricity?
- Why is graphite used as a lubricant?
- Why is diamond used in cutting tools?
