Anions

FIELDS OF STUDY: Inorganic Chemistry; Geochemistry; Metallurgy

ABSTRACT

The basic structure of anions is defined, and the development of the modern theory of atomic structure is elaborated. Electronic structure is fundamental to all chemical behaviors and is responsible for the relationships seen in the periodic table.

The Nature of Anions

An anion is any atom or group of atoms that bears a net negative charge due to the presence of one or more extra valence electrons. The term is a contraction of "anode ions," which is a reference to the fact that when a direct electric current is applied to an electrolytic solution, negatively charged ions are attracted to the anode, or positive terminal, of the source of the current. By the same token, cations, from "cathode ions," have a net positive charge and are attracted to the cathode, or negative terminal, of the current source. Anions and cations often combine to form compounds held together with ionic bonds; one common example is sodium chloride (NaCl), better known as table salt, which is created when the sodium cation, Na⁺, bonds to the chloride anion, Cl⁻.

The formation of any ion is the result of an atom or molecule gaining or losing one or more valence electrons. This is most apparent in monatomic (single-atom) ions, in which the electrical charge is equal to the oxidation state. For example, the halogens—fluorine, chlorine, bromine, iodine, and astatine—are all highly electronegative, meaning that they readily accept an extra valence electron so that their outer electron shell is completely full and therefore stable. The resulting anions are called fluoride, chloride, bromide, iodide, and astatinide, respectively, and they have an electrical charge of 1− because they gained one electron and thus one unit of negative charge. Similarly, the chalcogens oxygen (O) and sulfur (S) readily accept two extra electrons to form the oxide ion, O²⁻, and the sulfide ion, S²⁻.

Compound ions in which a central atom is bonded to a number of oxygen atoms are extremely common. These oxoanions tend to form very stable compounds and are the basic materials of many minerals.

The Electronics of Anion Formation

According to the modern theory of atomic structure, each atom contains a very small, extremely dense nucleus that holds at least 99.98 percent of the atom’s mass and all of its positive electrical charge. The nucleus is surrounded by a diffuse and comparatively very large cloud of electrons containing all of the atom’s negative electrical charge. These electrons are allowed to possess only very specific energies. This restricts their movement around the nucleus to specific regions called "electron shells." Within each shell are well-defined regions called "orbitals." The strict geometric arrangement of the orbitals regulates the formation of chemical bonds between atoms.

Each shell and orbital is subject to a number of restrictions that dictate how many electrons it can hold. There are four different types of electron orbitals, designated s, p, d, and f, each of which can contain a specific number of electrons: s orbitals can hold a maximum of two electrons; p orbitals, a maximum of six; d orbitals, a maximum of ten; and f orbitals, a maximum of fourteen. One or more of these orbitals make up an electron shell. The various electron shells are indicated by an integer value known as the "principal quantum number," starting with 1 for the innermost shell, typically referred to as the "n = 1 shell." The standard notation to describe an electron orbital is the principal quantum number, followed by the type of orbital, followed by a superscript number indicating how many electrons it holds. For example, helium has only one s orbital, which is completely full, so its electron configuration is represented as 1s². The first p orbital appears in the n = 2 shell, the first d orbital in the n = 3 shell, and the first f orbital in the n = 4 shell. Due to variances in energy levels, electrons usually fill atomic orbitals in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s . . . rather than in strict numerical order.

The outermost electron shell is the valence shell, and in all elements, except the noble gases (helium, neon, argon, krypton, xenon, and radon), the valence shell is incompletely occupied. Because of this, the noble gases are often called "inert gases," reflecting the fact that they are far less chemically reactive than elements whose valence shells are not filled. All noble gases except for helium have eight electrons in their valence shell, as a shell with eight electrons approximates a stable closed-shell configuration of s²p⁶. This is the basis of the octet rule, which states that atoms of lower atomic numbers tend to achieve the greatest stability when they have eight electrons in their valence shells. The closer an element is to having eight valence electrons, the more likely it is to undergo ionization or a chemical reaction in order to achieve a stable configuration, either by gaining enough electrons to complete the octet or by losing all valence electrons so that the next-highest completed shell becomes the outermost shell. Elements with similar electron distributions in their valence shells typically exhibit similar chemical behaviors, which is the basis on which the periodictable of the elements is arranged.

Naming the Anions

The rules for naming various chemical species are established and standardized by the International Union of Pure and Applied Chemistry (IUPAC). Monatomic anions and polyatomic anions composed of a single element are named by adding the suffix -ide to the name of the element, either instead of or in addition to the existing suffix. Thus a sulfur anion becomes sulfide, a xenon anion becomes xenonide, a potassium anion becomes potasside, and so on. In some cases, the suffix is added to the element’s Latin name instead of its common name; for example, an anion of silver, which has the Latin name argentum, is called argentide.

Oxoanions are generally named for the central atom of the anion and the number of oxygen atoms that surround it, with the charge number given in parentheses at the end. An oxoanion consisting of a sulfur atom surrounded by three oxygen atoms (SO₃²⁻) has the formal name trioxidosulfate(2−), while one with a sulfur atom and four oxygen atoms (SO₄²⁻) is called tetraoxidosulfate(2−). The common (nonsystematic) names of these two oxoanions are sulfite and sulfate, respectively, following the convention that the oxoanion with fewer oxygen atoms takes the suffix -ite and the one with more oxygen atoms takes the suffix -ate. If one element is capable of forming more than two different oxoanions, as is the case with chlorine, the prefixes hypo- and per- are used as well, so that the common name of ClO⁻ is hypochlorite, ClO₂⁻ is chlorite, ClO₃⁻ is chlorate, and ClO₄⁻ is perchlorate. Multiple central atoms are indicated by the appropriate numerative prefix; for example, the anion CrO₃²⁻ is chromate, while the anion Cr₂O₇²⁻ is dichromate.

PRINCIPAL TERMS

• cation: any chemical species bearing a net positive electrical charge, which causes it to be drawn toward the negative pole, or cathode, of an electrochemical cell.
• ionic bond: a type of chemical bond formed by mutual attraction between two ions of opposite charges.
• ionization: the process by which an atom or molecule loses or gains one or more electrons to acquire a net positive or negative electrical charge.
• oxoanion: an ion consisting of one or more central atoms bonded to a number of oxygen atoms and bearing a net negative electrical charge.
• valence electron: an electron that occupies the outermost or valence shell of an atom and participates in chemical processes such as bond formation and ionization.

Bibliography

Johnson, Rebecca L. Atomic Structure. Minneapolis: Twenty-First Century, 2008. Print.

Lide, David R., ed. CRC Handbook of Chemistry and Physics. 94th ed. Taylor and Francis, 2013. Web. 17 Mar. 2014.

Mackay, K. M., R. A. Mackay, and W. Henderson. Introduction to Modern Inorganic Chemistry. 6th ed. Cheltenham: Nelson, 2002. Print.

Miessler, Gary L., Paul J. Fischer, and Donald A. Tarr. Inorganic Chemistry. 5th ed. Boston: Pearson, 2014. Print.

Wehr, M. Russell, James A. Richards Jr., and Thomas W. Adair III. Physics of the Atom. 4th ed. Reading: Addison, 1984. Print.

Winter, Mark J. The Orbitron: A Gallery of Atomic Orbitals and Molecular Orbitals on the WWW. U of Sheffield, 2002. Web. 17 Mar. 2014.