Explain the trend in ionization energy down a group of the periodic table.

Ionization energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. This is the following process;
X _g ---> X+_g + e-
When an atom is ionized we break the electrostatic force of attraction between the nucleus and the outermost electron. The greater the magnitude of this attraction the greater the ionization energy will be. This attraction depends on three factors;
Firstly, the number of protons in the nucleus. More protons in the nucleus results in a greater nuclear charge and therefore a stronger electrostatic attraction between the nucleus and the outermost electrons. We would therefore expect ionization energy to increase as the number of protons in the nucleus increases.
Secondly, the distance between the nucleus and the electron being removed. The greater the distance the smaller the attraction between the nucleus and the outermost electron.
Finally we must consider an effect called shielding. This refers to repulsion between electrons in an atom. Inner electrons repel the outermost electron, reducing the overall attraction between the outermost electron and the nucleus. Greater shielding therefore reduces the magnitude of the ionisation energy.
Going down a group we observe an overall decrease in ionization energy. We consider each of the three effects in turn. The number of protons in the nucleus INCREASES going down the group. The distance between the nucleus and outermost electron INCREASES going down the group, as the electron is in a shell with a greater principle quantum number. The magnitude of shielding also INCREASES going down the group as we have added more electrons to the atom.
The first effect (which acts to increase IE) opposes the second and third (which both act to reduce IE). Overall, the increase in shielding and electron-nucleus distance outweigh the greater number of protons and ionization energy decreases down a group.