Covalent Bonding Flashcards
Covalent Bond
A shared pair of electrons
Describe the structure of graphite
The carbon atoms are arranged in sheets of flat hexagons covalently bonded with three bonds each.
The fourth outer electron of each carbon atom is delocalised.
The sheets of hexagons are bonded together by weak van der Waals forces
Explain why graphite feels slippery and is used as a dry lubricant and in pencils
The weak bonds between the layers in graphite are easily broken.
Therefore, the sheets can slide over each other
Explain why graphite is an electrical conductor
The delocalised electrons in graphite aren’t attached to any particular carbon atoms and are free to move along the sheets carrying a charge
Explain why graphite has a low density and is used to make strong, lightweight sports equipment
The layers are quite far apart compared to the length of the covalent bonds
Explain why graphite has a very high melting point
Because of the strong covalent bonds in the hexagon sheets (sublimes at over 3900 K)
Sublime
Change straight from a solid to a gas
Explain why graphite is insoluble in any solvent
The covalent bonds in the sheets are too strong to break
Describe the structure of diamond
Made up of carbon atoms.
Each carbon atom is covalently bonded to four other carbon atoms.
The atoms arrange themselves in a tetrahedral shape
Describe all the properties of diamond due to its strong covalent bonds
Diamond has a very high melting point- sublimes at over 3900 K.
Diamond is extremely hard- used in diamond tipped drills and saws.
Diamond won’t dissolve in any solvent.
You can cut diamond to form gemstones. Its structure makes it refract light a lot, which is why it sparkles
Explain why diamond is a good thermal conductor
Vibrations travel easily through the stiff lattice
Explain why diamond can’t conduct electricity
All the outer electrons are held in localised bonds
Co-ordinate Bond
Shown by (in a diagram)
A bond formed where both electrons come from one atom.
An arrow pointing away from the donor atom
