What are the trends in ionization energy across the periodic table?

Ionisation energy generally increases across a period from left to right and decreases down a group in the periodic table.

Ionisation energy, or the energy required to remove an electron from an atom, is influenced by three main factors: the distance of the electrons from the nucleus, the amount of positive charge in the nucleus, and the extent of electron shielding.

As you move across a period from left to right, the number of protons in the nucleus increases. This means there is a stronger positive charge pulling the electrons closer to the nucleus, making it harder to remove an electron and thus increasing the ionisation energy. The electrons are also added to the same energy level, so the increase in nuclear charge is not offset by an increase in distance or shielding.

On the other hand, as you move down a group, the outer electrons are further from the nucleus and there is more shielding from the inner electrons. Both of these factors make it easier to remove an electron, so the ionisation energy decreases.

However, there are some exceptions to these trends. For example, the ionisation energy decreases slightly from beryllium to boron and from nitrogen to oxygen. This is because in these cases, the added electron goes into an already partially filled orbital, which offers slightly less stability and thus less energy is required to remove it.

Understanding these trends in ionisation energy can help predict the reactivity of elements and their likelihood to form different types of ions. For example, elements with low ionisation energies, like those in Group 1, are likely to form positive ions, while those with high ionisation energies, like those in Group 17, are likely to form negative ions.