Question

Why is the first ionization energy of a non-metal significantly higher than that of an alkali metal?

Answer 1

We know that the amount of energy required to remove one or more electrons from a neutral atom to generate a positively charged ion is a physical feature that affects the atom's chemical activity. The energy required to remove the outermost, or highest energy, electron from a neutral atom in the gas phase is known as the first ionization energy of an element. The tendency of a neutral atom to resist the loss of electrons is measured by ionization energies.

Complete answer:
As mentioned before, the energy required to remove the outermost (valence) electron from a neutral atom in the gaseous state to create a cation is known as the element's initial ionization energy.
The atomic radius of an element has an inverse relationship with its first ionization energy. The first ionization energy increases and the atomic radius decreases as the periodic table moves from left to right. The first ionization energy reduces as the group number decreases, but the atomic radius increases.
The radii of alkali metal atoms are larger than those of nonmetals. As a result, the positively charged atomic nucleus's attraction to the valence electron is substantially lower than that of a nonmetal nucleus.
As a result, an alkali metal's first ionization energy is much lower than that of a nonmetal.

Note:
It must be noted that since nonmetal atoms have smaller radii, the positively charged nucleus attracts the valence electrons more strongly. Because the first ionization energy of a nonmetal is substantially higher than that of an alkali metal, nonmetals gain one or more electrons instead of losing them in an ionic bond, forming anions. The first ionization energy of neutral chlorine which is in period $3$ is $1251kJ/mol$