MCQ Detailed View

Electrical conductivity depends on the presence of free-moving charged particles such as electrons or ions. Among the given options, graphite is the only substance that conducts electricity in solid form.

Graphite is an allotrope of carbon where carbon atoms are arranged in layers of hexagonal lattices. Each carbon atom is bonded to three other carbon atoms via covalent bonds, leaving one electron per atom free to move within the layers. These delocalized electrons act as charge carriers, allowing graphite to conduct electricity efficiently along the layers.

In contrast, diamond, another allotrope of carbon, has a three-dimensional tetrahedral structure where each carbon atom is bonded to four others. There are no free electrons in diamond, so it does not conduct electricity.

Potassium and sodium are metals that conduct electricity in molten or aqueous form, but in solid state, they conduct poorly compared to graphite due to the fixed positions of metal ions in the lattice.

The conductivity of graphite is widely applied in electrodes, batteries, and lubricants. Its layered structure allows electrons to move freely, making it both an excellent conductor and a strong material in specific applications.

Understanding which substances conduct electricity is important in inorganic chemistry, as it helps students relate structure to properties, including metallic, covalent, and network solids. Graphite serves as a clear example of how delocalized electrons in covalent networks enable conduction, differentiating it from non-conducting forms like diamond.