Name the following ionic compounds: (a) \(\mathrm{NaF},\) (b) \(\mathrm{Mg}_{2} \mathrm{C}\) (c) \(\mathrm{Li}_{3} \mathrm{~N}\) (d) \(\mathrm{Al}_{2} \mathrm{O}_{3}\) (e) \(\mathrm{K}_{2} \mathrm{Se}\)

When it comes to naming monoatomic ionic compounds, the process is straightforward. These compounds consist of single atoms of a metal paired with single atoms of a non-metal. The key is to identify the cation, which is always named first and retains the element's name, and the anion, which takes the root of the element's name and adds the suffix -ide.

For instance, consider \(\mathrm{NaF}\), where \(\textrm{Na}\) is the sodium cation and \(\textrm{F}\) is the fluoride anion. The compound's name is 'Sodium fluoride.' This basic rule provides students with a foundation for understanding the structure and naming of more complex ionic compounds.

Confusion often arises with multivalent metals, which can form cations with more than one possible positive charge. The oxidation state of such a metal in a particular compound is indicated by a Roman numeral in parentheses following the metal's name. However, this step is unnecessary for metals with fixed oxidation states, like magnesium in \(\mathrm{Mg}_2\mathrm{C}\). Here magnesium has a +2 charge, and the number of atoms in the formula represents the simplest ratio of ions needed to balance charge. Therefore, 'Magnesium carbide' accurately defines the combination without needing to specify the oxidation state.

Polyatomic ions are charged entities composed of several atoms bonded together. In naming ionic compounds including polyatomic ions, the rules can differ. For monoatomic ions, such as the \(\textrm{Li}_{3}\textrm{N}\), you only follow the metal-nonmetal naming convention with an 'ide' suffix. But polyatomic ions have specific names like sulfate \(\textrm{SO}_4^{2-}\) or nitrate \(\textrm{NO}_3^{-}\), which are memorized rather than derived from naming rules. These names remain unchanged when forming compounds, making it essential for students to familiarize themselves with common polyatomic ions.

Compounds containing transition metals often require students to determine which oxidation state the metal is taking in a given compound. Transition metals are capable of forming ions with various charges, which complicates their naming. In cases like \(\mathrm{Al}_2\mathrm{O}_3\), the metal, aluminum, has a predictable +3 charge. However, other transition metals may have multiple common oxidation states, such as iron which can be +2 (ferrous) or +3 (ferric). It's imperative to specify the charge in such cases—for example, 'Iron(II) chloride' for \(\mathrm{FeCl}_2\) and 'Iron(III) chloride' for \(\mathrm{IFeCl}_3\). For metals with a single oxidation state, the simple naming method can still apply.

Understanding oxidation states is crucial in naming ionic compounds, as they indicate the degree of oxidation of an atom and help balance the total charge. Oxidation states are denoted by numbers, usually indicated by Roman numerals when naming compounds with multivalent metals. For example, the iron in 'Iron(II) oxide' has an oxidation state of +2. Mastery of oxidation states not only aids in proper naming but also serves as a guide for predicting formulae of compounds and their reactions. In the case of the classroom example \(\mathrm{K}_2\mathrm{Se}\), the +1 and -2 charges of potassium (K) and selenide (Se) explain why two potassium ions bond with one selenide ion to balance the charges in 'Potassium selenide.'