Why does ionization energy increase as we go from left to right in a period?

Question

Why does ionization energy increase as we go from left to right in a period?

In my textbook, the explanation is as follows:

"This is consistent with the idea that electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. Thus electrons in the same principal quantum level are generally more strongly bound as we move to the right on the periodic table, and there is a generally more increase in ionization energy values as electrons are added to a given principal quantum level."

This, however, doesn't make any sense to me. From what I understand, in a stepwise ionization process, it is always the highest-energy electron (the one bound least tightly) that is removed first. So when we go more to the right in the periodic table, there are more electrons, and thus much more possibility for these to shield the outer electron from attraction to the nucleus. This would make it easier to remove that outer electron, and so I would say the opposite: the ionization energy should lower when we go to the right.

So can someone explain to me why this is not so? Thanks for any clarifications.

Answer 1

Crash Course on Ionization Energy:

1. As we all know, atoms prefer a full valence shell. So as we go right in a period, we are increasing $e-$. And also ADDING PROTONS. Because we are adding protons, the size of the atom gets smaller because the nuclear charge will be more powerful. Adding protons in a period trumps the addition of electrons. At the end of the day, we have a small atom with many electrons in it's valence shell that does not want to let go of them. Especially the Noble Gasses and Halogens.
2. Ionization energy decreases as we move down a group because:
   • As we move down, a new full energy level is being added. More electrons means more repulsion.
   • This creates the shielding effect where the addition of the shells, shields the outer electron from receiving the nucleic charge. NOTE: Here, however the addition of another energy level trumps the addition of protons.

This is just a piece of the whole picture.

Answer 2

Although there're more shielding as you add more electrons, the nuclear charge also increases as you go from left to right because of the increase number of protons. The nuclear charge's pull on electrons has way more effect on the ionization energy than the gradual shielding does (shielding will have a bigger impact down the periods because of different energy levels). So the IE does increase from left to right.

Answer 3

Each time we move one atom across the period there is one more proton in the nucleus, this means that the nuclear charge is increasing. Also each time we move across the period we gain one electron, this electron goes into the same shell (all though the sub shells are slightly different in energies, the difference can be discarded when considering the attractive strength of the protons). Which means there is NO EXTRA SHIELDING of the valance shell electrons, as the electron is going into the valence shell. This means that the atomic radius decreases across the period as the nucleus is pulling all the shells closer to the nucleus as the nuclear charge is greater.

When considering ionisation energies we know that the first ionisation energy it is the energy required to remove the outer most electron. Regardless of the fact that there are more electrons the valance shell is closer to the nucleus therefore there is a greater attraction between the nucleus and the valance shell electrons, this greater attraction means it requires more energy to remove the outer most electron, hence the general trend of increasing first ionisation energies increasing across the period. Of course there are exceptions when you shift across the blocks and when pairing starts in the sub-shells, but that's a different question. Hope this helps.