How do covalent bonds conduct electricity?
Generally speaking, there are two requirements for a material to be a good electrical conductor:
- It must have a network of overlapping covalent bonds in 3 dimensions that hold the atoms together in a lattice.
- It must have some empty orbitals that are just slightly higher in energy than the filled orbitals that comprise the covalent bonds.
Most metal atoms have partially filled s or d orbitals that can form a network (band) of orbitals. Because the band is only partially filled, it is easy to add electrons to one end of a metal wire and extract other electrons from the opposite end of the wire.
Diamond is an example of a material that has a network of covalent bonds, but all of the low-energy orbitals are filled, making diamond an electrical insulator.
Semiconductors like silicon have a band gap that requires at least a bit of energy (usually less than 0.5 eV) to overcome in order to conduct electrons in the upper band (or holes in the lower band).