Why ionization energy increases from left to right

Since going from right to left on the periodic table, the atomic radius increases, and the ionization energy increases from left to right in the periods and up the groups. Exceptions to this trend is observed for alkaline earth metals group 2 and nitrogen group elements group Typically, group 2 elements have ionization energy greater than group 13 elements and group 15 elements have greater ionization energy than group 16 elements.

Groups 2 and 15 have completely and half-filled electronic configuration respectively, thus, it requires more energy to remove an electron from completely filled orbitals than incompletely filled orbitals. Alkali metals IA group have small ionization energies, especially when compared to halogens or VII A group see diagram 1.

In addition to the radius distance between nucleus and the electrons in outermost orbital , the number of electrons between the nucleus and the electron s you're looking at in the outermost shell have an effect on the ionization energy as well.

This effect, where the full positive charge of the nucleus is not felt by outer electrons due to the negative charges of inner electrons partially canceling out the positive charge, is called shielding. The more electrons shielding the outer electron shell from the nucleus, the less energy required to expel an electron from said atom.

The higher the shielding effect the lower the ionization energy see diagram 2. It is because of the shielding effect that the ionization energy decreases from top to bottom within a group.

From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy with the exception of Helium and Neon. Each succeeding ionization energy is larger than the preceding energy. Electron orbitals are separated into various shells which have strong impacts on the ionization energies of the various electrons.

For instance, let us look at aluminum. Aluminum is the first element of its period with electrons in the 3p shell.

This makes the first ionization energy comparably low to the other elements in the same period, because it only has to get rid of one electron to make a stable 3s shell, the new valence electron shell. However, once you've moved past the first ionization energy into the second ionization energy, there is a large jump in the amount of energy required to expel another electron. This is because you now are trying to take an electron from a fairly stable and full 3s electron shell. Electron shells are also responsible for the shielding that was explained above.

Both ionization energy and electron affinity have similar trend in the periodic table. This is called electron shielding. If we plot the first ionization energies vs. Moving from left to right across the periodic table, the ionization energy for an atom increases. We can explain this by considering the nuclear charge of the atom. The more protons in the nucleus, the stronger the attraction of the nucleus to electrons. This stronger attraction makes it more difficult to remove electrons.

Within a group, the ionization energy decreases as the size of the atom gets larger. On the graph, we see that the ionization energy increases as we go up the group to smaller atoms. In this situation, the first electron removed is farther from the nucleus as the atomic number number of protons increases.

Being farther away from the positive attraction makes it easier for that electron to be pulled off. Use the following link to answer the questions below:. Skip to main content. The Periodic Table. Search for:. Describe factors affecting ionization energy. Question Question fbdfe. Question 1a23b. Question d7de9. See all questions in Periodic Trends in Ionization Energy. Impact of this question views around the world. You can reuse this answer Creative Commons License.