Group I elements (alkali metals) have only one valence electron in their outermost shell. They readily lose this electron to achieve a stable octet configuration, forming a +1 ion (e.g., Na⁺, K⁺). Non-metals, particularly those in Group VII (halogens), have a high tendency to gain one electron to achieve a stable octet configuration, forming a -1 ion (e.g., Cl^-, Br^-). The electrostatic attraction between these oppositely charged ions leads to the formation of ionic bonds.

For example, Sodium (Group I) reacts with Chlorine (Group VII) to form Sodium Chloride (NaCl). Sodium readily loses its valence electron to Chlorine, which readily gains that electron. The resulting Na⁺ and Cl^- ions are strongly attracted to each other, forming the ionic compound NaCl.

Group VII elements (halogens) have seven valence electrons. They readily gain one electron to achieve a stable octet configuration, forming a -1 ion (e.g., Cl^-, Br^-). Metals have a tendency to lose electrons to form positive ions. The electrostatic attraction between the metal cations and the halogen anions results in the formation of ionic bonds.

For example, Magnesium (a metal) reacts with Chlorine (a halogen) to form Magnesium Chloride (MgCl₂). Magnesium readily loses two electrons to form Mg²⁺, and Chlorine readily gains one electron to form Cl^-. The resulting Mg²⁺ and Cl^- ions are strongly attracted to each other, forming the ionic compound MgCl₂.

The high melting and boiling points of ionic compounds are a direct consequence of the strong electrostatic forces of attraction between the oppositely charged ions in the ionic lattice. These forces are non-directional and extend throughout the entire crystal structure, requiring a significant amount of energy to overcome. This energy input is manifested as a high melting point.

In the solid state, the ions are fixed in a regular lattice arrangement. They cannot move freely and therefore cannot carry an electric current. This explains why solid ionic compounds are poor conductors of electricity.

When an ionic compound is melted or dissolved in water, the ions are free to move. In a molten state, the ions are not held in a fixed lattice and can move randomly. In an aqueous solution, the ions are surrounded by water molecules, which also allows them to move freely. This mobility of ions enables them to carry an electric current, making molten and aqueous ionic compounds good conductors of electricity. The forces holding the ions together (electrostatic attraction) are overcome by the energy input (heat or solvation), allowing for ion mobility and electrical conductivity.

Ionic compounds exhibit different electrical conductivity depending on their state:

Solid State: In the solid state, ionic compounds are poor conductors of electricity. This is because the ions are held in fixed positions within the crystal lattice and are not free to move. There are no free charge carriers.

Aqueous State (Dissolved in Water): When an ionic compound is dissolved in water, it dissociates into its constituent ions. These ions become mobile and are free to move throughout the solution. The mobile ions can then carry an electric current. Therefore, aqueous solutions of ionic compounds are good conductors of electricity.

Molten State (Liquid): When an ionic compound is heated to its melting point, it becomes a liquid. In the molten state, the ions are still mobile and free to move. This allows them to carry an electric current. Therefore, molten ionic compounds are also good conductors of electricity.

Summary Table:

The ability of ions to move freely is crucial for electrical conductivity. The presence of a solid lattice restricts ion movement, while a liquid or aqueous solution allows for it.