8.4: Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy (IE₁). The first ionization energy for an element, X, is the energy required to form a cation with 1+ charge:

The energy required to remove the second most loosely bound electron is called the second ionization energy (IE₂).

The energy required to remove the third electron is the third ionization energy, and so on. Energy is always required to remove electrons from atoms or ions, so ionization processes are endothermic and IE values are always positive. For larger atoms, the most loosely bound electron is located farther from the nucleus and so is easier to remove. Thus, as size (atomic radius) increases, the ionization energy should decrease.

Within a period, the IE₁ generally increases with increasing Z. Down a group, the IE₁ value generally decreases with increasing Z. There are some systematic deviations from this trend, however. Note that the ionization energy of boron (atomic number 5) is less than that of beryllium (atomic number 4) even though the nuclear charge of boron is greater by one proton. This can be explained because the energy of the subshells increases as l increases, due to penetration and shielding. Within any one shell, the s electrons are lower in energy than the p electrons. This means that an s electron is harder to remove from an atom than a p electron in the same shell. The electron removed during the ionization of beryllium ([He]2s²) is an s electron, whereas the electron removed during the ionization of boron ([He]2s²2p¹) is a p electron; this results in lower first ionization energy for boron, even though its nuclear charge is greater by one proton. Thus, we see a small deviation from the predicted trend occurring each time a new subshell begins.

Another deviation occurs as orbitals become more than one-half filled. The first ionization energy for oxygen is slightly less than that for nitrogen, despite the trend in increasing IE₁ values across a period. For oxygen, removing one electron will eliminate the electron-electron repulsion caused by pairing the electrons in the 2p orbital and will result in a half-filled orbital (which is energetically favorable). Analogous changes occur in succeeding periods.

Removing an electron from a cation is more difficult than removing an electron from a neutral atom because of the greater electrostatic attraction to the cation. Likewise, removing an electron from a cation with a higher positive charge is more difficult than removing an electron from an ion with a lower charge. Thus, successive ionization energies for one element always increase. As seen in Table 1, there is a large increase in the ionization energies for each element. This jump corresponds to the removal of the core electrons, which are harder to remove than the valence electrons. For example, Sc and Ga both have three valence electrons, so the rapid increase in ionization energy occurs after the third ionization.

Table 1: Successive Ionization Energies for Selected Elements (kJ/mol)

This text is adapted from OpenStax Chemistry 2e, Section 6.5: Periodic Variations in Element Properties.

The chemical behavior of atoms and ions is greatly affected by how easy or difficult it is to remove their electrons, especially the outermost electrons which participate in chemical bond formations. 

The energy required to remove an electron from a gaseous atom in its ground state is called the first ionization energy and is given in kJ/mol. The energy required to remove the next electron is called the second ionization energy, and so on.

Moving down a column, the ionization energies decrease. Recall that the highest principal quantum number of valence electrons increases down the column leading to larger atomic sizes. Thus, the farther the outermost electrons are, the easier they are to remove.

For main-group elements, the ionization energy increases across the period. The reason lies in the increasing atomic number, where valence electrons experience a higher effective nuclear charge making the removal of outermost electrons more difficult. This explains why chlorine has a higher ionization energy than sodium, for example. Generally, ionization energy is a minimum for an alkali metal and rises to a peak with each noble gas.

Transition metals display a small increase in the ionization energy across the period, and the f-block elements show an even smaller change. 

But there are some exceptions to consider.

Boron has a smaller ionization energy than beryllium, even though it is farther to the right on the periodic table. Beryllium has lower energy 2s electrons, whereas boron has a higher energy 2p electron making its removal energetically more favorable. 

Another exception is oxygen, which has lower first ionization energy than nitrogen. Compared to nitrogen, oxygen has four p-electrons, and removing one electron eliminates the electron-electron repulsion. Thus, less energy is required for the ionization. These exceptions are observed in succeeding periods too.

Electron removal from cations is more difficult than from neutral atoms. Generally, the successive ionization energies increase for elements. 

Consider potassium. The second ionization energy is significantly higher, as it involves the removal of a core electron from an ion with a noble gas configuration. 

Similarly, for calcium, there is a high increase from the second to third ionization energy as a core electron is removed from a cation with a noble gas configuration.