Definition

Ionization energy is the minimum energy necessary to remove one mole of electrons from one mole of atoms in the gaseous state. The first ionization energy refers to the energy necessary to rid of one electron, so it can be noted that the "nth" ionization energy is the energy required to rid of the atom's nth electron, after (n-1) electrons before it have been removed. The first ionization energy is a significant factor in chemical reactivity, due to the fact that atoms with low first ionization energies tend to become the cations during reactions, while those with high ionization energies (with noble gases being an exception) tend to convert to anions.

If the ionization energy is high, that means that it takes a lot of energy to remove the outermost electron. If the ionization energy is low, that means that it takes only a small amount of energy to remove the outermost electron.

As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove the outermost electron and that means there is a higher ionization energy. As you go across the periodic table, nuclear charge is the most important consideration. So, going across the periodic table, there should be an increase in ionization energy because of the increasing nuclear charge.

Trends

Ionization energies have an inverse relationship with atomic radius. As size decreases for an atom, more energy would be required to remove the atom.

Across the Period

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Ionization energies tend to increase across the period, due to the fact that there are a greater number of protons that that gain the attraction to the orbiting electrons more effectively, which would increase the energy necessary to remove an electron.

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Across the Group

Ionization energies tend to decrease down a group, due to the fact that it takes less energy to remove an outer-shell electron as opposed to an inner-shell electron. An exception to this pattern occurs in Group 3A (the group after the transition elements) of the periodic table, since it does not increase from aluminum (Al) to thallium (Tl).

As we go down the periodic table, the electrons are further from the nucleus, causing the atoms to be larger and the ionization energies to be lower.

Importance of Ionization Energy and Why it Matters

In most chemical reactions, a clear understanding of ionization energy is crucial in order to comprehend why certain bonds occur and the energy that is associated with them. It also provides for the understanding how of various atoms make covalent or ionic bonds with each other. For example, the alkali metal of sodium has an ionization energy of roughly 5000 kJ/mol. The ionization energy of the chlorine atom is roughly 1200 kJ/mol. When creating the sodium chloride complex, the difference in their ionization energy is so great that chemically they can combine as an ionic bond, the strongest of bonds known to chemistry. However, when you start to see compounds of atoms that have ionization energies that are much closer to each other, then we start seeing weaker and weaker bonds. For example, the ionization energies of both carbon and chloride are fairly close to each other which dictates their bond type as covalent, the second strongest bond in chemistry. The general rule regarding bond energies and bond strengths is that the closer the two elements are on the periodic table, the weaker the bond will be between the two.

References

Silberberg, Martin S. Principles of General Chemistry. Boston: McGraw-Hill Higher Education, 2007. Print.