Question

Assertion (A): Graphite is a good conductor of electricity, however, diamond belongs to the category of insulators.
Reason (R): Graphite is soft in nature on the other hand diamond is very hard and brittle

Answer 1

Hint:We know that diamond and graphite are both different allotropic forms of carbon. They vary in their structural arrangement and bonding and thus have different properties. To answer this question, you must recall their structures.

Complete answer:
In diamond, each carbon atom is $s{p^3}$ hybridized and thus can form four bonds. It is bonded to four other carbon atoms (also $s{p^3}$ hybridized) in a tetrahedral arrangement forming a three- dimensional tetrahedral cage- like network structure. Due to this structure, diamond is a hard and brittle material. All the four electrons of carbon are used up in bond formation and there are no free electrons in the network. Hence, it does not conduct electricity.
Whereas, in graphite, each carbon atom is $s{p^2}$ hybridized and forms three bonds. Each carbon atom is bonded to three other carbon atoms and thus, graphite has a layered sheet- like structure which makes it soft in nature. Only three out of the four valence electrons of carbon are involved in bond formation. The fourth electron remains unpaired and can easily move between the layers of graphite under the influence of applied potential. Thus, it conducts electricity.

We can conclude by saying that both assertion and reason are correct, but the reason is not the correct explanation for assertion.

Note:
We know that the atomic number of carbon is 6 and it carries four valence electrons. Its ground state electronic configuration can be depicted as:
$C:\left[ {He} \right]2{s^2}2{p^2}$
In the excited state, one electron is excited from the $2s$ orbital to the $2p$ orbital. The electronic configuration in the excited state becomes:
$C:\left[ {He} \right]2{s^1}2{p^3}$
Thus, we can see that carbon has four electrons that can take place in bonding and it can easily form four covalent bonds.