Ionization Energy Trends in the Periodic Table

The ionization energy of an atom is the amount of energy required to remove an electron from the gaseous form of that atom or ion.

1^st ionization energy - The energy required to remove the highest energy electron from a neutral gaseous atom.

For Example:

Notice that the ionization energy is positive. This is because it requires energy to remove an electron.

2^nd ionization energy - The energy required to remove a second electron from a singly charged gaseous cation.

For Example:

The second ionization energy is almost ten times that of the first because the number of electrons causing repulsions is reduced.

3^rd ionization energy - The energy required to remove a third electron from a doubly charged gaseous cation.

For Example:

The third ionization energy is even higher than the second.

Successive ionization energies increase in magnitude because the number of electrons, which cause repulsion, steadily decrease. This is not a smooth curve There is a big jump in ionization energy after the atom has lost its valence electrons. An atom that has the same electronic configuration as a noble gas is really going to hold on to its electrons. So, the amount of energy needed to remove electrons beyond the valence electrons is significantly greater than the energy of chemical reactions and bonding. Thus, only the valence electrons (i.e., electrons outside of the noble gas core) are involved in chemical reactions.

The ionization energies of a particular atom depend on the average electron distance from the nucleus and the effective nuclear charge

These factors can be illustrated by the following trends:

1^st ionization energy decreases down a group.

This is because the highest energy electrons are, on average, farther from the nucleus. As the principal quantum number increases, the size of the orbital increases and the electron is easier to remove.

Examples:

I₁(Na) > I₁(Cs)

I₁(Cl) > I₁(I)

1^st ionization energy increases across a period.

This is because electrons in the same principal quantum shell do not completely shield the increasing nuclear charge of the protons. Thus, electrons are held more tightly and require more energy to be ionized.

Examples:

I₁(Cl) > I₁(Na)

I₁(S) > I₁(Mg)

The graph of ionization energy versus atomic number is not a perfect line because there are exceptions to the rules that are easily explained.

Filled and half-filled subshells show a small increase in stability in the same way that filled shells show increased stability. So, when trying to remove an electron from one of these filled or half-filled subshells, a slightly higher ionization energy is found.

Example 1:

I₁(Be) > I₁(B)

It's harder to ionize an electron from beryllium than boron because beryllium has a filled "s" subshell.

Example 2:

I₁(N) > I₁(O)

Nitrogen has a half-filled "2p" subshell so it is harder to ionize an electron from nitrogen than oxygen.