3.2 Ionic compounds  (Page 3/5)

Exceptions to the expected behavior involve elements toward the bottom of the groups. In addition to the expected ions Tl ³⁺ , Sn ⁴⁺ , Pb ⁴⁺ , and Bi ⁵⁺ , a partial loss of these atoms’ valence shell electrons can also lead to the formation of Tl ⁺ , Sn ²⁺ , Pb ²⁺ , and Bi ³⁺ ions. The formation of these 1+, 2+, and 3+ cations is ascribed to the inert pair effect    , which reflects the relatively low energy of the valence s -electron pair for atoms of the heavy elements of groups 13, 14, and 15. Mercury (group 12) also exhibits an unexpected behavior: it forms a diatomic ion, ${\text{Hg}}_2{}^{\text{2+}}$ (an ion formed from two mercury atoms, with an Hg-Hg bond), in addition to the expected monatomic ion Hg ²⁺ (formed from only one mercury atom).

Transition and inner transition metal elements behave differently than main group elements. Most transition metal cations have 2+ or 3+ charges that result from the loss of their outermost s electron(s) first, sometimes followed by the loss of one or two d electrons from the next-to-outermost shell. For example, iron (1 s ² 2 s ² 2 p ⁶ 3 s ² 3 p ⁶ 3 d ⁶ 4 s ² ) forms the ion Fe ²⁺ (1 s ² 2 s ² 2 p ⁶ 3 s ² 3 p ⁶ 3 d ⁶ ) by the loss of the 4 s electron and the ion Fe ³⁺ (1 s ² 2 s ² 2 p ⁶ 3 s ² 3 p ⁶ 3 d ⁵ ) by the loss of the 4 s electron and one of the 3 d electrons. Although the d orbitals of the transition elements are—according to the Aufbau principle—the last to fill when building up electron configurations, the outermost s electrons are the first to be lost when these atoms ionize. When the inner transition metals form ions, they usually have a 3+ charge, resulting from the loss of their outermost s electrons and a d or f electron.

Electronic structures of anions

Most monatomic anions form when a neutral nonmetal atom gains enough electrons to completely fill its outer s and p orbitals, thereby reaching the electron configuration of the next noble gas. Thus, it is simple to determine the charge on such a negative ion: The charge is equal to the number of electrons that must be gained to fill the s and p orbitals of the parent atom. Oxygen, for example, has the electron configuration 1 s ² 2 s ² 2 p ⁴ , whereas the oxygen anion has the electron configuration of the noble gas neon (Ne), 1 s ² 2 s ² 2 p ⁶ . The two additional electrons required to fill the valence orbitals give the oxide ion the charge of 2– (O ^2– ).

Key concepts and summary

Atoms gain or lose electrons to form ions with particularly stable electron configurations. The charges of cations formed by the representative metals may be determined readily because, with few exceptions, the electronic structures of these ions have either a noble gas configuration or a completely filled electron shell. The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells.

Chemistry end of chapter exercises