Here's why:
* Increased Nuclear Charge: As you move across a period, the number of protons in the nucleus increases. This means the positive charge of the nucleus is stronger, pulling the electrons closer and making them more difficult to remove.
* Similar Electron Shielding: The number of electron shells remains the same across a period. While the inner electrons do shield the outer ones from the full nuclear charge, this shielding effect stays relatively constant.
* Decreased Atomic Radius: The atomic radius decreases across a period due to the increasing nuclear charge. This means the outermost electrons are closer to the nucleus, experiencing a stronger attraction and making it harder to remove them.
Exceptions:
There are a few exceptions to this general trend, mainly due to electron configurations:
* Group 13: Ionization energies decrease slightly between groups 2 and 13. This is because the 3rd electron enters a p-orbital, which is higher in energy than the s-orbital. It's easier to remove an electron from a higher energy level.
* Group 16: Ionization energies decrease slightly between groups 15 and 16. This is because the 4th electron in group 16 elements pairs up with an existing electron in a p-orbital. Electron-electron repulsion makes it slightly easier to remove this paired electron.
Overall, the increasing trend in ionization energy across a period reflects the increasing attraction between the nucleus and electrons.
