Ionic Compounds: Formulas and Nomenclature

An element composed of atoms that readily lose electrons (a metal) can react with an element composed of atoms that readily gain electrons (a nonmetal) to produce ions through complete electron transfer. The compound formed by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the oppositely charged ions.

Moving from the far right to the left on the periodic table, nonmetal elements often gain electrons to form anions with the same number of electrons as an atom of the next noble gas in the periodic table, and a negative charge equal to the number of groups moved left from the noble gases. That is, atoms of group 17 gain one electron and form anions with a 1− charge; atoms of group 16 gain two electrons and form ions with a 2− charge, and so on. For example, a neutral oxygen atom, with 8 protons and 8 electrons, readily gains two electrons. This results in an anion with 8 protons, 10 electrons, and a 2− charge, and is symbolized as O²⁻. The anion, O²⁻, has the same number of electrons as the next noble gas – neon. The name of the anions is the name of the nonmetallic element with its ending replaced by the suffix -ide, so O²⁻ is called oxide.

Transition metals and some other metals often exhibit variable charges that are not predictable by their location in the table. For example, copper can form ions with a 1+ or 2+ charge, and iron can form ions with a 2+ or 3+ charge.

In every ionic compound, the total number of positive charges of the cations equals the total number of negative charges of the anions. Thus, ionic compounds are electrically neutral, even though they contain positive and negative ions. The formula of an ionic compound must have a ratio of ions such that the numbers of positive and negative charges are equal.

For example, if a compound contains aluminum and oxygen in the form of  Al³⁺ and O²⁻, the formula of the compound would be Al₂O₃. Two aluminum ions, each with a charge of 3+, would give us six positive charges, and three oxide ions, each with a charge of 2−, would give us six negative charges. Thus, the compound will be electrically neutral, with the same number of positive and negative charges.

Many ionic compounds contain polyatomic ions as the cation, the anion, or both. Polyatomic ions are a group of bonded atoms that act as discrete units, carrying an overall charge. As with simple ionic compounds, these compounds must also be electrically neutral, so their formulas can be predicted by treating the polyatomic ions as discrete units. Parentheses are used in the formula to indicate polyatomic ions that behave as a unit. For example, the formula for calcium phosphate, one of the minerals in our bones, is Ca₃(PO₄)₂. The compound contains the polyatomic ion PO₄³⁻, consisting of one phosphorus atom and four oxygen atoms, and having an overall charge of 3−. This formula indicates that there are three Ca²⁺ ions (total six positive charges) for every two PO₄³⁻ groups (total six negative charges). The compound is electrically neutral, and its formula shows a total count of three Ca, two P, and eight O atoms.

Ionic compounds are symbolized by formula, indicating the relative numbers of its constituent ions. For compounds containing only monatomic ions (such as NaCl) and for many compounds containing polyatomic ions (such as CaSO₄), these formulas are just the empirical formulas. However, the formulas for some ionic compounds containing polyatomic ions are not empirical formulas. For example, the ionic compound sodium oxalate is composed of Na⁺ and C₂O₄²⁻ ions combined in a 2:1 ratio, and its formula is written as Na₂C₂O₄. 

Nomenclature of Ionic Compounds

The name of a binary compound containing monatomic ions consists of the name of the cation (the name of the metal) followed by the name of the anion (the name of the nonmetallic element with its ending replaced by the suffix -ide). For example, the name for Na₂O is sodium oxide.

Compounds containing polyatomic ions are also named similarly to those containing only monatomic ions, i.e., by naming first the cation and then the anion. For example, the name for CaSO₄ is calcium sulfate.

Most of the transition metals and some main group metals can form two or more cations with different charges. Compounds of these metals with nonmetals are named with the same method as binary compounds, except the charge of the metal ion is specified by a Roman numeral in parentheses after the name of the metal. 

The charge of a metal ion is determined from the formula of the compound and the charge of anion. For example, in a binary ionic compound of iron and chlorine, iron typically exhibits a charge of either 2+ or 3+, and the two corresponding compound formulas are FeCl₂ and FeCl₃. In such cases, the charge of the metal ion is included as a Roman numeral in parentheses immediately following the metal name. These two compounds are hence named iron(II) chloride and iron(III) chloride, respectively. 

Ionic compounds that contain water molecules as integral components of their crystals are called hydrates. The name for an ionic hydrate is derived by adding a term to the name for the anhydrous (meaning “not hydrated”) compound that indicates the number of water molecules associated with each formula unit of the compound. The added word begins with a Greek prefix denoting the number of water molecules and ends with “hydrate.” For example, the anhydrous compound copper(II) sulfate also exists as a hydrate containing five water molecules and named copper(II) sulfate pentahydrate (penta = 5). Washing soda is the common name for a hydrate of sodium carbonate containing ten water molecules; the systematic name is sodium carbonate decahydrate (deca = 10).

Formulas for ionic hydrates are written by appending a vertically centered dot, a coefficient representing the number of water molecules, and the formula for water. For example, copper(II) sulfate pentahydrate is written as CuSO₄∙5H₂O.

This text is adapted from Openstax, Chemistry 2e, Section 2.6: Molecular and Ionic Compounds and Openstax, Chemistry 2e, Section 2.7: Chemical Nomenclature.