Why does ionisation energy of elements generally decrease as you move down a group in the periodic table?

Ionisation energy is the energy required to remove the most loosely bound (furthest out) electron from an atom, in the gaseous state.
As you move down a group in the periodic table, the number of protons in each element increases - meaning nuclei of atoms further down the group have a more positive charge. This might lead you to expect the ionisation energy to increase, as the more positive nucleus of atoms further down the group should attract negative electrons more strongly.

However, this effect is offset by the fact that the atoms further down the group have more energy levels (as they have more electrons). This means that the furthest out electron in an atom lower down the group is much further away than the furthest out electron from an atom at the top of the group - meaning the attraction it feels from the nucleus is less strong. Additionally, the electrons in the inner energy levels 'sheild' the outer electrons from the postive nuclues' attraction - further decreasing the attraction felt by the outermost electron.

These two effects combine to overcome the increasing positive charge of the nucleus so that in the atom at the bottom of a group, the outermost electron is much bound much less strongly than in an atom at the top. This means it can be more readily removed from the atom so ionisation energy is lower.