6.5 Ionization Energy Trend

Periodic trends in the energetics of ion formation

Chemical reactions are based largely on the interactions between the most loosely bound electrons in atoms, so it is not surprising that the tendency of an atom to gain, lose or share electrons should play an important role in determining its chemical properties.

Periodic trends in ionization energy

The term “ionization energy” always refers to removal of electrons of an atom, leading to the formation of positive ions. In order to remove an electron from an atom, work must be done to overcome the electrostatic attraction between the electron and the nucleus; this work is called the ionization energy of the atom and corresponds to the exothermic process

M(g) → M⁺(g) + e^–

in which M(g) stands for any isolated (gaseous) atom.

An excellent series of  videos on ionization energy and electron affinity from UC-Berkeley:

Ionization energy, electron affinity  (6 min)

Quiz on ionization energy 

An atom has as many ionization energies as it has electrons. Electrons are always removed from the highest-energy occupied orbital. An examination of the successive ionization energies of the first ten elements (below) provides experimental confirmation that the binding of the two innermost electrons (1s orbital) is significantly different from that of the n=2 electrons. Successive ionization energies of an atom increase rapidly as reduced electron-electron repulsion causes the electron shells to contract, thus binding the electrons even more tightly to the nucleus.

The table lists the energies, in electron volts, required to remove each successive electron from the first ten elements.

Figure 6.54 Ionization energy Trend

It’s worth taking some time to examine the rather abrupt jumps in the sequence E1 through E10 as the atomic number increases. Note the very large jumps in the energies required to remove electrons from the 1s orbitals of atoms of the second-row elements Li-Ne.

Ionization energies increase with the nuclear charge Z as we move across the periodic table. They decrease as we move down the table because in each period the electron is being removed from a shell one step farther from the nucleus than in the atom immediately above it. This results in the familiar zig-zag lines when the first ionization energies are plotted as a function of Z.

 Figure 6.55 Periodic Trend in first Ionization energy

Figure 6.56 Periodic Trend in Period 2 Ionization energy

This more detailed plot of the ionization energies of the atoms of the first ten elements reveals some interesting irregularities that can be related to the slightly lower energies (greater stabilities) of electrons in half-filled (spin-unpaired) relative to completely-filled subshells.

Depicting the ionization energies of the main group elements in the context of the periodic table offers a more comprehensive view of these trends.

Figure 6.57 Periodic Trend in Ionization energy

Some points to note:

• The noble gases have the highest IE’s of any element in the period. This has nothing to do with any mysterious “special stability” of the s²p⁶ electron configuration; it is simply a matter of the high nuclear charge acting on more contracted orbitals.
• IE’s (as well as many other properties) tend not to vary greatly amongst the d-block elements. This reflects the fact that as the more-compact d orbitals are being filled, they exert a screening effect that partly offsets that increasing nuclear charge on the outermost s orbitals of higher principal quantum number.
• Each of the Group 13 elements has a lower first-IE than that of the element preceding it. The reversal of the IE trend in this group is often attributed to the more easy removal of the single outer-shell p electron compared to that of electrons contained in filled (and thus spin-paired) s– and d-orbitals in the preceding elements.