Ionization Energy

Definition of IE: the minimum amount of energy required to remove an electron form an atom (in kJ/mol) in its gaseous state.

Ex. X (g) → X⁺  (g) +   1e^–     IE₁

Ex. X⁺  (g) → X⁺² (g)  + 1e^–     IE₂

Ex. X⁺²  (g) → X⁺³ (g)  + 1e^–     IE₃

Sample Ionization Energies Table

General Pattern in the Table: for each electron lost, the subsequent IE ↑

Thus IE₁ < IE₂ < IE₃

Alkali Metal (Group I) Comparison: Li, Na, and K all want to lose 1 electron, so the 1^st IE values are not too high, but increase dramatically when any further electrons are to be removed (2^nd IE, 3^rd IE).

Within the group, an ↑ n value will have a ↓IE comparatively because Z_eff is ↓.

Alkali Metals (Group I) vs Alkaline Earth (Group II) Metals: 1st IE for both groups is comparable; however, the 2^nd IE increase is much greater for Group I than Group II because Alkali metals only want to lose 1 electron and will resist any further loses by having ↑IE. Alkaline earth metals want to lose 2 electrons, so they do not resist the second electron loss, hence the lower 2^nd IE value compared to that of alkali metals.

Trend: As you proceed from:

LEFT → RIGHT   IE ↑

BOTTOM → TOP  IE ↑

Why?

Right-most elements are nonmetals and tend to gain electrons so they’re not too keen on losing electrons and will resist the removal of any electrons. Therefore, in order to remove an electron from such an atom, a lot of energy must be expended, hence a lot of ionization energy.

For metals, on the left-most section, who want to lose electrons, not as much energy is required, as they are willing to give up electrons readily, so they have ↓ IE values.

Top-most elements are closer to nucleus ( ↓ n value) and have ↑ Z_eff values (electrons held tightly by the nucleus and difficult to remove) and ↑ IE.