Cations

FIELDS OF STUDY: Inorganic Chemistry; Geochemistry; Metallurgy

ABSTRACT

The basic structure of cations 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 Cations

A cation is any atom or group of atoms that bears a net positive charge due to the absence of one or more valence electrons. The term is a contraction of "cathode ion," which is a reference to the fact that when a direct electric current is applied to an electrolytic solution, positively charged ions are attracted to the cathode, or negative terminal, of the source of the current. By the same token, anions, from "anode ions," have a net negative charge and are attracted to the anode, or positive 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 alkali metals—hydrogen, lithium, sodium, potassium, rubidium, cesium, and francium—are all highly electropositive, meaning that they readily donate their lone valence electron so that their next electron shell becomes the outermost shell, which is ideal because it is full and therefore stable. The resulting cations have an electrical charge of 1+ because they lost one electron and thus one unit of negative electrical charge. Similarly, the alkaline earth metals, such as barium (Ba) and magnesium (Mg), have two electrons in their valence shells, which they readily lose to form cations such as Ba²⁺ and Mg²⁺. Such cations are said to be "divalent," meaning that they can form single bonds with up to two other ions or molecules. They can also be described as multivalent or polyvalent, which simply means that they can form single bonds with more than one other ion or molecule. Multivalent atoms also have the potential to exist in different oxidation states—that is, they can lose or gain different numbers of electrons to form ions of different electrical charges; phosphorous, for example, commonly forms cations with charges of 5+ and 3+ as well as an anion with a charge of 3−.

The Electronics of Cation 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 one or more subshells that contain even more 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 Cations

The rules for naming various chemical species are established and standardized by the International Union of Pure and Applied Chemistry (IUPAC). Monatomic cations are simply called by the name of the element itself; when it is important to note the charge, this can be given in parentheses immediately following the name, so that a sodium cation, for example, is formally rendered as sodium(1+). If a compound contains a multivalent atom that can form ions of different charges (and thus different oxidation states), the particular ion present can be indicated in the compound’s formal name in the same way, as in mercury(2+) chloride (HgCl₂), in which the mercury cation has a charge of 2+. Alternatively, the oxidation number can be used in the place of the charge number; this is given as a roman numeral, so that mercury(2+) chloride is also mercury(II) chloride. It is important to remember that only monatomic ions, such as Hg²⁺ in this example, can be assumed to have oxidation numbers that are equal to their charge numbers, as the oxidation number of a polyatomic ion takes into account a number of other factors.

The parenthetical method of indicating charge also applies to polyatomic cations. For example, a molecule consisting of three hydrogen atoms and having a charge of 1+ is called trihydrogen(1+), while the far more complex cation [Al(POCl₃)₆]³⁺ has the IUPAC name hexakis(trichloridooxidophosphorus)aluminum(3+). The latter is an example of a coordination complex, which consists of a central atom or cation, typically metallic, surrounded by various other molecules or anions called "ligands." Such complexes require their own systematic names that indicate the number and identity of the ligands present. The names of these complexes can become quite complicated, requiring careful attention to the IUPAC naming conventions.

If the polyatomic cation is derived from a hydride (a compound containing a hydrogen anion, H⁻) and gained its positive charge not by losing an electron but by accepting a hydrogen cation (H⁺), a process known as "protonation," it can also be named by adding the suffix -ium to the name of the parent hydride. For example, when ammonia (NH₃) accepts a hydrogen cation, it forms ammonium (NH₄⁺).

PRINCIPAL TERMS

• anion: any chemical species bearing a net negative electrical charge, which causes it to be drawn toward the positive pole, or anode, 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.
• multivalent: describes an atom that has the ability to accept or donate more than one valence electron and thus can exist in more than one oxidation state.
• 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

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