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When thinking about ionization energy comparisons between several elements, you have to take both the period and the group into account rather than just the groups. Therefore the ionization energy trends can’t just be described with one direction, such as down. The ionization energy increases in the right direction of a period because, as a period moves to the right, the elements have more protons resulting in a stronger pull of the electrons. The ionization energy increases up the groups because, as groups go up, they have less and less shells, resulting in less electron shielding.
Another way to think about it is that because ionization energy is the energy required to remove an electron from a gaseous atom/ion, you can understand it in terms of reactivity. The higher the ionization energy, the harder it is to remove an electron. So, as you go across a period, the atomic radius of the elements decrease so there is a greater attraction between the negative electrons and the positive nucleus. Ionization energy decreases going down because shielding occurs, making the outermost (valence) electrons increasingly far from the nucleus.
Instead of thinking about what diagonal direction each trend increases/decreases toward, think about how they change down a group and across a period. IE increases across the period to the right because increasing Z attracts the e- closer to the nucleus and removing an e- from a shell that's almost full takes a lot of E. IE decreases down a group because an increased number of shells shields outside e- from being as attracted to the nucleus, making the outside e- easier to remove.
I think it's easier to think about how ionization increases across a period and decreases down a group. When you go from left to right, the atoms have an increasing amount of protons which results in a higher nuclear charge, pulling in the e- tighter. When you go from top to bottom, you are adding subshells which blocks the attraction of the outer e-, decreasing the ionization energy.
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