Chemistry – Atoms, elements and compounds - Giant covalent structures | e-Consult

Diamond is a poor conductor of electricity because of its giant covalent structure. In diamond, all valence electrons of the carbon atoms are involved in covalent bonding, forming a completely filled network of shared electrons. There are no free electrons available to carry an electric current. Electrical conductivity requires the presence of mobile charge carriers, typically electrons. Since diamond lacks these mobile electrons, it does not conduct electricity.

Silicon(IV) oxide (SiO₂) is also generally a poor conductor due to its giant covalent structure. In SiO₂, the silicon-oxygen bonds are primarily covalent, and the electrons are tightly bound within the covalent bonds. While SiO₂ does not have free electrons, it can exhibit semiconducting properties under certain conditions, particularly with the presence of impurities (doping). However, in its pure, crystalline form, the lack of mobile charge carriers prevents it from conducting electricity effectively.

The key similarity is that both materials have their valence electrons tightly held within strong covalent bonds, preventing the free flow of electrons necessary for electrical conduction. This is a direct consequence of their giant covalent structures.