THE PHYSICAL PROPERTIES OF GRAPHITE
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.
- 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.
Attractions between solvent molecules and carbon atoms will never be strong enough to overcome the strong covalent bonds in graphite.
- Conducts electricity. The delocalized electrons are free to move throughout the sheets. If a piece of graphite is connected into a circuit, electrons can fall off one end of the sheet and be replaced with new ones at the other end.
The structure of silicon dioxide, SiO2 Silicon dioxide is also known as silicon (IV) oxide. The giant covalent structure of silicon dioxide There are three different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure. Crystalline silicon has the same structure as diamond.
To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms. Notice that each silicon atom is bridged to its neighbours by an oxygen atom. Don’t forget that this is just a tiny part of a giant structure extending on all 3 dimensions.
THE PHYSICAL PROPERTIES OF DIAMOND
COVALENT BONDING – SINGLE BONDS