The first ionization energy of the metals follows the order:

Correct option is B

​ Ionization energy (IE) is the minimum energy required to remove the most loosely bound electron of an isolated gaseous atom, positive ion, or molecule. The first ionization energy is quantitatively expressed as

​where X is any atom or molecule, X⁺ is the resultant ion when the original atom was stripped of a single electron, and e^- is the removed electron. Comparison of ionization energies of atoms in the periodic table reveals two periodic trends which follow the rules of Coulombic attraction:

Ionization energy generally increases from left to right within a given period (that is, row). Ionization energy generally decreases from top to bottom in a given group (that is, column).

Ionization energy is also a periodic trend within the periodic table. Moving left to right within a period, or upward within a group, the first ionization energy generally increases. As the nuclear charge of the nucleus increases across the period, the electrostatic attraction increases between electrons and protons, hence the atomic radius decreases, and the electron cloud comes closer to the nucleus because the electrons, especially the outermost one, are held more tightly by the higher effective nuclear charge.

On moving downward within a given group, the electrons are held in higher-energy shells with higher principal quantum number n, further from the nucleus and therefore are more loosely bound so that the ionization energy decreases. The effective nuclear charge increases only slowly so that its effect is outweighed by the increase in n.

​The elements here, zinc (₃₀Zn: 9.4 eV), cadmium (₄₈Cd: 9.0 eV) and mercury (₈₀Hg: 10.4 eV) all record sudden rising IE values. For mercury, it can be extrapolated that the relativistic stabilization of the 6s electrons increases the ionization energy, in addition to poor shielding by 4f electrons that increases the effective nuclear charge on the outer valence electrons.

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