Ion and Ionization

In the Bohr model of atomic structure, electrons orbit the nucleus at fixed distances, similar to the orbits of the planets around the sun. For every element, the distances of the electron orbitals are fixed and unique to that element. Normally, the electrons occupy the orbits closest to the nucleus. This is the most stable configuration of the atom and is known as the ground state. To move an electron to an orbital further from the nucleus requires the input of energy. Atoms which have an electron in a higher orbit are said to be in an excited state.

The strength of attraction between a negatively charged electron and the positively charged nucleus is greater the closer together they are. The energy needed to move an electron from one orbit to a higher energy one is equal to the difference in the attraction between the two configurations; it takes increasing amounts of energy to move an electron to orbits further and further from the nucleus. The energy needed to move electrons from one orbit to another can be thought of like the energy needed to move between rungs on a ladder; to move from a lower rung to a higher rung requires the input of energy, and the more rungs you move up, the more energy it takes. However, if the electron is moved too far from the nucleus, the attraction between the electron and the nucleus is too small to hold the electron in its orbit, and, analogous to stepping of the top rung of the ladder, the electron is separated from the atom leaving behind a positively charged atom; the atom has been ionized.

The ionization of an atom can be represented by:

where X is a single atom of any element and e< is the ejected electron. The amount of energy required for this process is called the ionization energy. The ionization energy is a measure of how difficult it is to remove the electron from the atom—the more strongly the electron is attracted to the nucleus, the higher the ionization energy. Although in theory it is possible to remove any of the electrons from an atom, in practice, the electron in the outermost orbit is typically the first to be removed. The energy required to remove the first electron is called the first ionization energy.

For many electron atoms it is possible to remove more than one electron. A second electron can be removed

from a singly charged ion X⁺ to yield a doubly charged ion, X²⁺. This process can be written as:

The energy required to remove the second electron is called the second ionization energy. Following the removal of the first electron, the atom has one more positively charged proton in the nucleus than it has negatively charged orbiting electrons. This charge imbalance causes the remaining electrons to be held even more tightly to the nucleus. Consequently, more energy is required to remove the second electron than was required to remove the first. The removal of subsequent electrons, creating X³⁺, X⁴⁺, and so on, requires ever increasing amounts of energy. This effect is rather like a small child with a collection of toys. The child might be easily persuaded to give the first toy away, but will hold on to each remaining toy with increasing vigor, thereby requiring increasing amounts of persuasion to give away each subsequent toy.

The first and second ionization energies of the elements in the first three rows of the periodic table are listed in Table I.

Note that the second ionization energy in all cases is larger than the first ionization energy. The hydrogen atom, however, having only one electron, only has a first ionization energy. Note that the ionization energy, in general, increases with increasing atomic number for elements within the same row of the periodic table. The ionization energy is smallest for the alkaline earth elements, Li, Na, K, etc., increasing with atomic number and reaching a maximum at the end of each row, corresponding to the noble gases, Ne, Ar, Kr, etc. This effect is related to the way in which atomic orbitals are filled. The noble gases have filled electronic orbitals, which are very stable.

Molecules can be ionized in a manner analogous to atoms. However, because electrons form the bonds that hold molecules together, their removal may result in the bond being weakened, or even broken. The ionization energies of some simple molecules are listed in Table II.

Note that in general, the ionization energies of molecules have values the same order of magnitude as the first ionization energies of isolated atoms. Molecules with only a few atoms, such as N₂, CO₂ and H₂O, tend to have the highest ionization energies. Within a group of similar molecules, such as the alkanes listed in the table, the ionization energy decreases with increasing size. This effect is due to the fact that in larger molecules, there are more electrons available for ionization without disrupting the bonding stability of the molecule. Again, this is analogous to persuading a child to give up its toys; the more toys the child has, the easier it will be to persuade it to give one up.