Structures+-+Tom

   __**D****iamond**__

 __**Physical Properties**__

 Diamond  > > __ **The structure of diamond** __ >  >  > Carbon has an electronic arrangement of 2,4. In diamond, each carbon shares electrons with four other carbon atoms - forming four single bonds. Like the diagram below. > > > This is a **giant** covalent structure - it continues on and on in three dimensions. It is not a molecule, because the number of atoms joined up in a real diamond is completely variable - depending on the size of the crystal. > > > **Graphite** > > **The physical properties of graphite** > Graphite >  >>
 * has a very high melting point (almost 4000°C) **because** very strong carbon-to-carbon covalent bonds have to be broken throughout the structure before melting occurs.
 * is very hard. This is again **because** you need to break very strong covalent bonds in 3-dimensions.
 * doesn't conduct electricity **because** the electrons are held tightly between the atoms, and aren't free to move.
 * is insoluble in water and organic solvents **because** there are no possible attractions which could occur between solvent molecules and carbon atoms which could outweigh the attractions between the covalently bound carbon atoms.
 * has a high melting point, similar to that of diamond. In order to melt graphite, it isn't enough to loosen one sheet from another. You have to break the covalent bonding throughout the whole structure.
 * has a soft, slippery feel, and is used in pencils and as a dry lubricant for things like locks. //You can think of graphite rather like a pack of cards - each card is strong, but the cards will slide over each other, or even fall off the pack altogether. When you use a pencil, sheets are rubbed off and stick to the paper. (Wikipedia)//
 * has a lower density than diamond. This is because of the relatively large amount of space that is "wasted" between the sheets.
 * is insoluble in water and organic solvents - for the same reason that diamond is insoluble.

**The bonding in graphite** Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These "spare" electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer. They are no longer associated directly with any particular atom or pair of atoms, but are free to wander throughout the whole sheet. 

  The important thing is that the delocalised electrons are free to move anywhere within the sheet - each electron is no longer fixed to a particular carbon atom. There is, however, no direct contact between the delocalised electrons in one sheet and those in the neighbouring sheets. The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. So what holds the sheets together? In graphite you have the ultimate example of van der Waals dispersion forces. As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal.

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 * The structure of graphite**

<span style="font-family: Helvetica,Arial;">Graphite has a **//layer structure.//** The diagram below shows the arrangement of the atoms in each layer, and the way the layers are spaced. <span style="font-family: Helvetica,Arial;">

<span style="font-family: Helvetica,Arial;">The distance between the layers is about 2.5 times the distance between the atoms within each layer. <span style="font-family: Helvetica,Arial;">The layers, of course, extend over huge numbers of atoms - not just the few shown above. <span style="font-family: Helvetica,Arial;">Carbon has to form 4 bonds because of its 4 unpaired electrons, whereas in this diagram it only seems to be forming 3 bonds to the neighbouring carbons. This diagram is something of a simplification, and shows the arrangement of atoms rather than the bonding.