The Ionization energy of Oxygen

[sylvie_1D]

Hi! I was answering number 81 on the chapter 2 homework and it asked to explain why oxygen's ionization energy is lower than that of nitrogen. I am just a bit confused because the trend on the periodic table is that ionization energy increases across group, so why would oxygen's be lower than nitrogen's?

[Heung Ching Chia 1E]

The general trend is ionization energy increases across a period due to the increased pull from the nucleus (+ve charge) as you are adding electrons to the same principal shell so it will be harder to remove electrons across the period.
However, oxygen is one of the exceptions. This is because nitrogen has a half filled 2p subshell and is stable, whilst oxygen has one more electron and trying to add one more electron to a half filled subshell would result in electron repulsion and less stability compared to a stable half filled subshell. Therefore, the first electron that can be removed from oxygen (first ionization energy) is lower than that of nitrogen.

[Samantha Draghi 1L]

I am confused about this too because if a half filled subshell causes it to be more stable and have a lower ionization energy than would this be the same as for Sulfur having a lower ionization energy than Phosphorous? and if so then why is the trend of ionization energy across a period increasing?