Question

Carbon forms a number of allotropes, two of which are graphite and diamond. Silicon has a diamond structure. Why is there no allotrope of silicon with a graphite structure?

Short answer

\(Si\) has an empty dorbital in which it can accept electrons, which makes it react as a Lewis acid, whereas a carbon doesn't have an emptydorbital. So, a carbon can form carbon-carbon bonds, while silicon can't.

Step-by-step solution

Definition of allotropes

Allotropes are classified as distinct forms of an element in which atoms are attached indifferent ways. This property of elements is called allotropy and is observed in carbon, sulfur, and phosphorus.

Explanation of no allotrope of silicon with a graphite structure

The electronic configurations of silicon and carbon are as follows:

\(\begin{aligned}{}{\rm{Si:}}\left( {{\rm{Ne}}} \right){\rm{3}}{{\rm{s}}^{\rm{2}}}{\rm{3}}{{\rm{p}}^{\rm{2}}}.\\{\rm{C:}}\left( {{\rm{He}}} \right){\rm{2}}{{\rm{s}}^{\rm{2}}}{\rm{2}}{{\rm{p}}^{\rm{2}}}.\end{aligned}\)

The difference lies in the electron configurations. While silicon has an emptyd orbital in which it can accept electrons, which makes it react as a Lewis acid (electron acceptor), a carbon doesn't have an empty dorbital.

Also, elements in the second period can form strong\(\pi \) bonds which in the third period and other elements cannot. The \(\pi \) bonds are the ones forming multiple bonds. Therefore, for this reason,a carbon can form carbon-carbon bonds, while silicon can't.