Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first or initial ionization energy or Ei of an atom or molecule is the energy required to remove one mole of electrons from one mole of isolated gaseous atoms or ions. You may think of ionization energy as a measure of the difficulty of removing electron or the strength by which an electron is bound. The higher the ionization energy, the more difficult it is to remove an electron. Therefore, ionization energy is in indicator of reactivity. Ionization energy is important because it can be used to help predict the strength of chemical bonds. Also Known As: ionization potential, IE, IP, ΔH° Units: Ionization energy is reported in units of kilojoule per mole (kJ/mol) or electron volts (eV). Ionization, together with atomic and ionic radius, electronegativity, electron affinity, and metallicity, follows a trend on the periodic table of elements.
The energy required to remove the outermost valence electron from a neutral atom is the first ionization energy. The second ionization energy is that required to remove the next electron, and so on. The second ionization energy is always higher than the first ionization energy. Take, for example, an alkali metal atom. Removing the first electron is relatively easy because its loss gives the atom a stable electron shell. Removing the second electron involves a new electron shell that is closer and more tightly bound to the atomic nucleus. The first ionization energy of hydrogen may be represented by the following equation: H(g) → H+(g) + e- ΔH° = -1312.0 kJ/mol If you look at a chart of first ionization energies, two exceptions to the trend are readily apparent. The first ionization energy of boron is less than that of beryllium and the first ionization energy of oxygen is less than that of nitrogen. The reason for the discrepancy is due to the electron configuration of these elements and Hund's rule. For beryllium, the first ionization potential electron comes from the 2s orbital, although ionization of boron involves a 2p electron. For both nitrogen and oxygen, the electron comes from the 2p orbital, but the spin is the same for all 2p nitrogen electrons, while there is a set of paired electrons in one of the 2p oxygen orbitals.
Why does ionization energy increase as we go from left to right in a period? In my textbook, the explanation is as follows: "This is consistent with the idea that electrons added in the same principal quantum level do not completely shield the increasing nuclear charge caused by the added protons. Thus electrons in the same principal quantum level are generally more strongly bound as we move to the right on the periodic table, and there is a generally more increase in ionization energy values as electrons are added to a given principal quantum level." This, however, doesn't make any sense to me. From what I understand, in a stepwise ionization process, it is always the highest-energy electron (the one bound least tightly) that is removed first. So when we go more to the right in the periodic table, there are more electrons, and thus much more possibility for these to shield the outer electron from attraction to the nucleus. This would make it easier to remove that outer electron, and so I would say the opposite: the ionization energy should lower when we go to the right. So can someone explain to me why this is not so? Thanks for any clarifications. |