Giant covalent structures: diamond, graphite and graphene

Science · Matter & Materials · ages 15-16

Name ______________________ Date ____________

Why do graphite's layers slide over each other so easily?
- the bonds inside each layer are unusually weak
- the free electrons push the layers apart
- only weak forces hold one layer to the next
- each layer is a separate molecule that can float away
Diamond and graphite are made of the same element.
Circle one: True False
How many of each carbon atom's four outer electrons are free to move in graphite?
Answer: ______________
How many covalent bonds does each carbon atom make in graphite?
Answer: ______________
In silicon dioxide, how many oxygen atoms is each silicon atom bonded to?
Answer: ______________
Which of these is a giant covalent structure built from two different elements?
- graphene, which is carbon and hydrogen
- silicon dioxide, which is silicon and oxygen
- diamond, which is carbon and nitrogen
- graphite, which is carbon with added oxygen atoms
Diamond and graphite are both pure carbon, yet one cuts glass and the other marks paper. What accounts for the difference?
- diamond's carbon atoms are heavier than graphite's
- diamond is a compound, graphite an element
- the two use a different number of bonds per atom
- graphite has been contaminated during its formation
Diamond grit is bonded to the rim of a cutting disc. Which property makes that work?
- it melts into the cut and seals it
- it is transparent, so the cut can be seen
- it is light enough to spin quickly
- it is harder than whatever is being cut
Graphite conducts, so Priya expects diamond and silicon dioxide to conduct too. What has she missed?
- neither diamond nor silicon dioxide has a free electron
- diamond and silicon dioxide do conduct, just extremely slowly
- conduction depends on the melting point, which differs
- graphite is not really a giant covalent substance at all
Nadia rubs a block of graphite between her fingers, gets a grey smear with almost no effort, and says its covalent bonds must be weak. Where is the slip?
- graphite has no covalent bonds at all, only weak forces
- pencil lines come from the paper, not from the graphite
- graphite really is soft because its bonds are unusually weak
- the strong bonds are inside a sheet, not between sheets

Answer key

For grown-ups. Fold this page away before handing over the rest.

Giant covalent structures: diamond, graphite and graphene W1-mt_YTEj9srwPG-s1

only weak forces hold one layer to the next · Inside a layer the bonding is strong covalent. Between layers there is only a weak pull, and that is the part a fingertip can beat.
True · Both are pure carbon. Everything that differs between them comes from how the atoms are arranged, not from what they are.
1 · Three go into bonds within the sheet, which leaves one. That single delocalised electron per atom is what carries the current.
3 · Graphite's carbons join into flat sheets of hexagons, and a hexagon needs only three bonds at each corner.
4 · Silicon is in group 4, so it makes four bonds, and in silicon dioxide all four go to oxygen atoms.
silicon dioxide, which is silicon and oxygen · Diamond, graphite and graphene are all pure carbon. Silicon dioxide is the one here that alternates two elements through the network.
the two use a different number of bonds per atom · Diamond spends all four bonds on a three-dimensional network. Graphite spends three on flat sheets and leaves the fourth electron free, so the sheets slide and the electron conducts.
it is harder than whatever is being cut · Cutting works by scratching, and only a harder material can scratch a softer one. Diamond's four-bond network makes it the hardest natural substance there is.
neither diamond nor silicon dioxide has a free electron · Graphite conducts because its fourth electron belongs to no bond and can travel along the sheet. Diamond has nothing loose at all, and in silicon dioxide nothing is free to move either.
the strong bonds are inside a sheet, not between sheets · What gives way is the attraction between one sheet and the next, and that was never a covalent bond. Inside a sheet the bonding is as strong as diamond's, which is why graphite also stays solid at thousands of degrees.