1. Solid State: In a solid compound, the constituent particles (atoms, molecules, or ions) are held together by strong intermolecular forces such as ionic bonds, covalent bonds, or hydrogen bonds. These forces create a rigid structure where particles are fixed in their positions and cannot move easily. As a result, when an electric field is applied, there are no free mobile ions or electrons available to carry the electric current, and the compound behaves as an electrical insulator.
2. Molten State: When a compound is heated to its melting point, it undergoes a phase change from a solid to a liquid state. During melting, the intermolecular forces between the particles are weakened, and the particles gain enough kinetic energy to overcome these forces and move more freely. This increased mobility allows the ions or electrons in the compound to move and respond to an applied electric field. As a result, the compound becomes electrically conductive in the molten state.
For instance, consider sodium chloride (NaCl) as an example. In solid NaCl, the sodium (Na+) and chloride (Cl-) ions are held together by strong ionic bonds, forming a rigid crystal lattice. In this state, NaCl does not conduct electricity because the ions are immobile and cannot move to carry electric current. However, when NaCl is melted, the ionic bonds weaken, and the Na+ and Cl- ions become free to move. In this molten state, NaCl can conduct electricity due to the mobility of its ions.
In summary, the difference in electrical conductivity between solid and molten compounds arises from the mobility of their constituent particles. In the solid state, strong intermolecular forces restrict particle movement, inhibiting electrical conduction. In the molten state, weakened intermolecular forces allow particles to move freely, enabling the passage of electric current and making the compound electrically conductive.
