Why does ionization energy increase going down a group but decreases going across a period?

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. The trends you've mentioned are due to variations in atomic structure as you move across a period or down a group in the periodic table.

Increasing down a group: As you move down a group, the number of electron shells increases. This means that the outermost electrons are farther away from the nucleus, experiencing weaker attractive forces. Consequently, it becomes easier to remove an electron, requiring less energy, leading to a decrease in ionization energy.

Decreasing across a period: As you move across a period from left to right, the number of protons in the nucleus increases, resulting in stronger attractive forces between the nucleus and the electrons. Additionally, within the same energy level, the effective nuclear charge (the net positive charge experienced by the outermost electrons) increases. This stronger attraction makes it more difficult to remove electrons, requiring more energy, leading to an increase in ionization energy.

In summary, the increase in ionization energy across a period is due to the increasing nuclear charge and the decrease down a group is due to the increasing distance between the outermost electrons and the nucleus.