Ionic compounds are formed when electrons are transferred between a metal and a nonmetal, resulting in ions held together by a strong electrostatic attraction. This transfer creates positively charged ions (cations) and negatively charged ions (anions). Correctly writing the chemical formulas and assigning systematic names to these compounds is essential for communicating in chemistry.
Determining the Charges of Constituent Ions
The first step in constructing an ionic compound formula involves determining the specific charge, or oxidation state, of the individual ions. Atoms engage in electron transfer primarily to achieve a stable electron configuration, which typically mirrors that of a noble gas. Metals, found on the left side of the periodic table, tend to lose electrons, forming positive cations, while nonmetals on the right side generally gain electrons to form negative anions. The periodic table provides a reliable guide for predicting these charges for main group elements.
For example, all elements in Group 1 consistently form ions with a 1+ charge, and Group 2 elements form ions with a 2+ charge. Nonmetals in Group 17 gain one electron to achieve a 1- charge, and elements in Group 16 gain two electrons to form a 2- charge.
Writing Formulas for Binary Ionic Compounds
Formulas for ionic compounds must always reflect electrical neutrality, meaning the total positive charge from the cations must balance the total negative charge from the anions, resulting in a net charge of zero. A practical technique for achieving this balance with simple binary compounds is the criss-cross method, which uses the magnitude of the charge on one ion as the subscript for the other.
For instance, when combining magnesium (2+ ion) with chlorine (1- ion), the “2” from the magnesium charge becomes the subscript for chlorine, and the “1” from the chloride charge is omitted. The resulting formula is written as MgCl₂. When charges have the same magnitude, such as with Mg²⁺ and O²⁻, the charges cancel out in a one-to-one ratio, yielding MgO.
The initial application of the criss-cross method may sometimes yield subscripts that are not in the lowest whole-number ratio. For example, combining the 2+ ion of calcium with the 2- ion of sulfur would initially suggest Ca₂S₂, but this must be reduced to the simplest ratio, CaS.
Naming Simple Ionic Compounds
The name of the compound is constructed by stating the name of the cation first, which is simply the name of the metal element itself. This is followed by the name of the anion, which is derived from the nonmetal element’s root name with the suffix “-ide” replacing the element’s original ending.
For example, a compound formed from the sodium ion (Na⁺) and the chloride ion (Cl⁻) is named sodium chloride. Similarly, the combination of a magnesium cation (Mg²⁺) and an oxide anion (O²⁻) is called magnesium oxide.
A specific rule for ionic compound nomenclature is the absence of numerical prefixes, such as mono- or di-, which are used in other types of chemical naming. Therefore, the compound MgCl₂ is correctly named magnesium chloride, not magnesium dichloride, because the subscripts are implicitly determined by the fixed charges of the ions.
Handling Compounds with Variable Charge Metals and Polyatomic Ions
Naming and writing formulas become more complex when dealing with metals that can form multiple different charges (variable charge metals) or when polyatomic ions are present. Many transition metals and some lower main group metals, like lead and tin, can exist in more than one oxidation state.
To differentiate between these possibilities, the specific charge of the metal ion is indicated using Roman numerals enclosed in parentheses immediately following the metal’s name. For instance, iron can form both a 2+ ion and a 3+ ion, so FeCl₂ is named iron(II) chloride, while FeCl₃ is named iron(III) chloride. The Roman numeral dictates the charge of the metal, which is determined by working backward from the known charge of the anion to ensure the compound is neutral.
Polyatomic ions are distinct groups of two or more covalently bonded atoms that carry an overall positive or negative charge, such as nitrate (NO₃⁻) or sulfate (SO₄²⁻). When writing formulas with these ions, they are treated as a single unit, and their name is used directly. If more than one polyatomic ion is needed to balance the charge, the entire ion must be enclosed in parentheses, with the required subscript placed outside. For example, calcium nitrate, which combines a Ca²⁺ ion with two NO₃⁻ ions, is written as Ca(NO₃)₂.