What trend is observed in ionization energy across the periodic table?

Ionization energy is defined as the energy required to remove an electron from an atom in its gaseous state. Across the periodic table, there is a clear trend regarding how ionization energy varies.

As you move from left to right across a period, the ionization energy tends to increase. This is primarily due to the increasing nuclear charge, which results from the addition of protons to the nucleus as you move across a period. The higher positive charge effectively pulls the electrons closer to the nucleus, making them more difficult to remove.

Conversely, as you move down a group, the ionization energy generally decreases. This trend occurs because the electrons being added to each successive element are located in higher energy levels that are further away from the nucleus. Additionally, the effect of electron shielding comes into play; the inner electrons shield the outermost electrons from the full attractive force of the nucleus, reducing the energy required to remove an outer electron.

Thus, the observation that ionization energy decreases as you go down a group and increases as you move across a period aligns with the fundamental principles of atomic structure and electron interactions.