The energy required to remove one electron from an isolated, gas-phase atom is the first ionization energy, abbreviated IE. Since the electron is attracted to the positive nucleus, energy must always be provided to complete this process (i.e., the first ionization of an atom is always an endothermic process). This energy could be provided in the form of a photon, h, as shown below. Electromagnetic radiation of known frequency is used, and the kinetic energy of the ejected electron is measured. The law of conservation of energy dictates that the difference
between the energy of the ionizing radiation and the energy of the ejected
electron equals the energy required for ionization. Ionization energies are typically
reported either in kJ/mol or in eV (1 eV = 96.485 kJ/mol).
As might be expected based on the periodic trend for effective nuclear charge, Zeff, ionization energy generally increases to the right across a period and decreases down a given group. A greater Zeff means the valence electrons experience more positive charge from the nucleus. Since coulombic attraction depends (among other things) upon the magnitude of the charges, increasing Zeff leads to a greater ionization energy. Although this trend generally holds true, a plot of IE versus Z (the plot appears on the next page) reveals some noteworthy deviations. In particular, notice that in period two the first ionization energies of boron (Z = 5) and oxygen (Z = 8) are unexpectedly low. The first ionization energies (in eV) for the period one and two elements are given below.
The first ionization energy increases from H to He as expected, then, also as expected, drops considerably on proceeding to Li. The IE again increases progressing across period two from Li to Be, but then decreases as we move from Be to B. Then, as predicted, the IE increases as we move to C and N, but decreases when we move to O, before increasing again as we complete the period with F and Ne. Both of these anomalies can be understood by examining the valence electron configurations and valence orbital diagrams for the elements involved.
Shown below are the valence electron configurations and valence orbital diagrams for Be and B.
The ionization of Be requires removal of an electron from the 2s subshell. Ionization
of B, however, requires removal of an electron from a 2p subshell. The 2p subshell is
higher in energy than the 2s subshell (this why we fill the 2s subshell first when
constructing an electron configuration) resulting in a lower than expected first
ionization energy for boron.
The electron configurations and orbital diagrams for N and O are shown below.
On comparing N and O, we see that oxygen has a fourth electron in its 2p
subshell. Since this subshell contains only three orbitals, this fourth electron must be
"paired"; in other words, it must occupy an orbital that already contains an
electron. Electrons, being negatively charged particles, repel each other and it is this
increased electron-electron repulsion that causes oxygen to have a lower than expected
first ionization energy. Although the electrons in nitrogen's 2p subshell also repel one
another, they reside in separate orbitals and can therefore remain farther apart than the
paired electrons in oxygen's 2p subshell. Notice that if we were to ionize an atom of
oxygen we would form O1+ which has the same electron configuration and orbital
arrangement as an atom of nitrogen.
An examination of the plot of IE versus Z shows that all of the atoms in groups 13 and 16 have first ionization energies that are lower than expected based on the general trend. Since elements are grouped based on electron configurations, these deviations are what we would predict. All of the group 13 elements have a single electron in a p subshell; all of the group16 elements have one pair of electrons in a p subshell.
Atoms other than hydrogen can, of course, lose more than one electron. The process of removing a second electron is always more endothermic than the first ionization energy; to remove a third electron requires still more energy, etc. This greater energy requirement is due to the fact that we are now trying to remove an electron from a system that has a net positive charge. Ionization energies for Li (1st and 2nd), Be (1st, 2nd, and 3rd), and B (1st, 2nd, 3rd and 4th) are given below in eV.
As you can see, second ionization energies, IE2, are always greater than first ionization energies, IE1, for the same element. Notice that IE2 for Li is much, much greater than IE2 for Be, and that IE3 for Be is much, much greater than IE3 for B. We can understand this pattern by considering the electron configurations for these three elements.
Li 1s22s1 Be 1s22s2 B 1s22s22p1
In all cases, the first electron is removed from a valence orbital resulting in the following electron configurations.
Li1+ 1s2 Be1+ 1s22s1 B1+ 1s22s2
In order to form Li2+, we need to take an electron from the core 1s subshell of Li1+. This 1s subshell is much lower in energy than the 2s subshell and, as a result, the second IE2 for lithium is much greater than IE1. To form Be2+ and B2+ we are still removing electrons from the valence shell. For this reason, IE2 is much more endothermic for lithium than for beryllium or boron. Similar reasoning can be used to compare IE3 for Be and B.
A final precautionary note: as pointed out earlier, ionization energies are always endothermic. People may sometimes say that the alkali metals "want to lose one electron" to form monocations, or that the alkaline earth elements "want to lose two electrons" to form dications, but such references are incorrect and can lead to confusion. Even Cs, in the lower left corner of the periodic table has a very endothermic first ionization energy of 376 kJ/mol. The fact that some elements are typically found as cations does not mean that they "want" to lose electrons. We will see that we must consider many factors in order to understand why the alkali metals usually exist as monocations despite the fact that they have quite endothermic ionization energies.