Arrange these elements in order of increasing first ionization energy: Si, F, In, N.

Key Concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. It is a key factor in determining an element's reactivity and is influenced by the atomic size and the effective nuclear charge. Generally, ionization energy increases across a period due to increasing nuclear charge and decreases down a group due to increased distance from the nucleus.

Trends in the Periodic Table

The periodic table exhibits specific trends in ionization energy, where elements in the same group have lower ionization energies than those in the same period. As you move from left to right across a period, ionization energy tends to increase due to the increasing positive charge of the nucleus, which attracts electrons more strongly. Conversely, moving down a group, ionization energy decreases because the outer electrons are further from the nucleus and experience greater shielding.

Comparison of Elements

When comparing elements like Si, F, In, and N, it is essential to consider their positions in the periodic table. Fluorine (F) has the highest ionization energy due to its small atomic size and high electronegativity, while indium (In) has the lowest due to its larger atomic radius and lower effective nuclear charge. Silicon (Si) and nitrogen (N) fall in between, with nitrogen having a higher ionization energy than silicon due to its position in the same period but further to the right.