WHAT IS CONDUCTIVITY?
Conductivity is the ability of a substance to carry an electrical
current. This current is the result of electrically-charged particles
gaining energy and subsequently travelling through the material,
carrying them to such a place as some of their energy can be lost,
usually with the intention of powering a useful process.
To explain this in more detail, I shall use the classic model of a
conductive material--metal:
The conductivity of metals is a direct result of a unique type of bond.
This metallic bond consists approximately of the metal nuclei and most
of their associated electrons being held in a lattice structure.
Holding them in place, however, are the valence electrons of each metal
atom--the electrons which take part in chemical reactions.
The important property of this type of bond is that despite
the valence orbitals having been expanded away from their associated
atom to become a large, ‘gaseous’ collection of molecular orbitals, the
electrons still retain a strong electrostatic attraction to
the metal ions in the lattice. This is due to their negative
charge and the positive charge on the ions--these two charges will
attract each other in an effort to minimise charge separation and hence
energy of the system.
In order to carry a current, however, the electrons must be free to
move throughout the metal, and indeed they are. This is because they
are no longer strongly associated with any particular atom and thus can
move
throughout the lattice of ions under the influence of a potential
difference, yet requiring only a very small amount of energy to do this.
It is important to point out the difference at this point between the
metallic bond and the covalent bond, in which the valence orbitals
overlap in order to allow atoms to share electrons. In the metallic
bond, these orbitals are poorly-defined, having been effectively
replaced by the
delocalised cloud of electrons. This allows for the possibility of the
metal being used as a conduit for energised electrons (e.g. in an
electrical cable), since there is no inherent reason why the metal
would ‘prefer’ its original electrons to those from
an adjacent atom. The covalent bond, however--though
expanded over perhaps several atoms--is very much stronger, forcing the
electrons to stay in place.
