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 1st IE values are not too high, but increase dramatically when any further electrons are to be removed (2nd IE, 3rd IE).
Within the group, an ↑ n value will have a ↓IE comparatively because Zeff is ↓.
Alkali Metals (Group I) vs Alkaline Earth (Group II) Metals: 1st IE for both groups is comparable; however, the 2nd 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 2nd 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 ↑ Zeff values (electrons held tightly by the nucleus and difficult to remove) and ↑ IE.
