Write the mmsanotherstage2019.comical formula for a simple ionic compound. Recognize polyatomic ions in mmsanotherstage2019.comical formulas.

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We have already encountered some mmsanotherstage2019.comical formulas for simple ionic compounds. A mmsanotherstage2019.comical formula is a concise list of the elements in a compound and the ratios of these elements. To better understand what a mmsanotherstage2019.comical formula means, we must consider how an ionic compound is constructed from its ions.

Ionic compounds exist as alternating positive and negative ions in regular, three-dimensional arrays called crystals (Figure \(\PageIndex{1}\)). As you can see, there are no individual \(\ce{NaCl}\) “particles” in the array; instead, there is a continuous lattice of alternating sodium and chloride ions. However, we can use the ratio of sodium ions to chloride ions, expressed in the lowest possible whole numbers, as a way of describing the compound. In the case of sodium chloride, the ratio of sodium ions to chloride ions, expressed in lowest whole numbers, is 1:1, so we use \(\ce{NaCl}\) (one \(\ce{Na}\) symbol and one \(\ce{Cl}\) symbol) to represent the compound. Thus, \(\ce{NaCl}\) is the mmsanotherstage2019.comical formula for sodium chloride, which is a concise way of describing the relative number of different ions in the compound. A macroscopic sample is composed of myriads of NaCl pairs; each individual pair called a formula unit. Although it is convenient to think that \(\ce{NaCl}\) crystals are composed of individual \(\ce{NaCl}\) units, Figure \(\PageIndex{1}\) shows that no single ion is exclusively associated with any other single ion. Each ion is surrounded by ions of opposite charge.


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Figure \(\PageIndex{1}\): A Sodium Chloride Crystal. A crystal contains a three-dimensional array of alternating positive and negative ions. The precise pattern depends on the compound. A crystal of sodium chloride, shown here, is a collection of alternating sodium and chlorine ions.

The formula for an ionic compound follows several conventions. First, the cation is written before the anion. Because most metals form cations and most nonmetals form anions, formulas typically list the metal first and then the nonmetal. Second, charges are not written in a formula. Remember that in an ionic compound, the component species are ions, not neutral atoms, even though the formula does not contain charges. Finally, the proper formula for an ionic compound always has a net zero charge, meaning the total positive charge must equal the total negative charge. To determine the proper formula of any combination of ions, determine how many of each ion is needed to balance the total positive and negative charges in the compound.

This rule is ultimately based on the fact that matter is, overall, electrically neutral.


By convention, assume that there is only one atom if a subscript is not present. We do not use 1 as a subscript.


If we look at the ionic compound consisting of lithium ions and bromide ions, we see that the lithium ion has a 1+ charge and the bromide ion has a 1− charge. Only one ion of each is needed to balance these charges. The formula for lithium bromide is \(\ce{LiBr}\).

When an ionic compound is formed from magnesium and oxygen, the magnesium ion has a 2+ charge, and the oxygen atom has a 2− charge. Although both of these ions have higher charges than the ions in lithium bromide, they still balance each other in a one-to-one ratio. Therefore, the proper formula for this ionic compound is \(\ce{MgO}\).

Now consider the ionic compound formed by magnesium and chlorine. A magnesium ion has a 2+ charge, while a chlorine ion has a 1− charge:

\<\ce{Mg^{2+}Cl^{−}}\>

Combining one ion of each does not completely balance the positive and negative charges. The easiest way to balance these charges is to assume the presence of two chloride ions for each magnesium ion:

\<\ce{Mg^{2+} Cl^{−} Cl^{−}}\>

Now the positive and negative charges are balanced. We could write the mmsanotherstage2019.comical formula for this ionic compound as \(\ce{MgClCl}\), but the convention is to use a numerical subscript when there is more than one ion of a given type—\(\ce{MgCl2}\). This mmsanotherstage2019.comical formula says that there are one magnesium ion and two chloride ions in this formula. (Do not read the “Cl2” part of the formula as a molecule of the diatomic elemental chlorine. Chlorine does not exist as a diatomic element in this compound. Rather, it exists as two individual chloride ions.) By convention, the lowest whole number ratio is used in the formulas of ionic compounds. The formula \(\ce{Mg2Cl4}\) has balanced charges with the ions in a 1:2 ratio, but it is not the lowest whole number ratio.


By convention, the lowest whole-number ratio of the ions is used in ionic formulas. There are exceptions for certain ions, such as \(\ce{Hg2^{2+}}\).


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Table \(\PageIndex{1}\): Some Polyatomic Ions

Polyatomic ions have defined formulas, names, and charges that cannot be modified in any way. Table \(\PageIndex{2}\) lists the ion names and ion formulas of the most common polyatomic ions. For example, \(\ce{NO3^{−}}\) is the nitrate ion; it has one nitrogen atom and three oxygen atoms and an overall 1− charge. Figure \(\PageIndex{2}\) lists the most common polyatomic ions.

Table \(\PageIndex{2}\): Ion Names and Ion Formulas of Common Polyatomic Ions Ion NameIon Formula
ammonium ion NH4+1
hydroxide ion OH−1
cyanide ion CN−1
carbonate ion CO3−2
bicarbonate or hydrogen carbonate HCO3−
acetate ion C2H3O2−1 or CH3CO2−1
nitrate ion NO3−1
nitrite ion NO2−1
sulfate ion SO4−2
sulfite ion SO3−2
phosphate ion PO4−3
phosphite ion PO3−3

Note that only one polyatomic ion in this Table, the ammonium ion (NH4+1), is a cation. This polyatomic ion contains one nitrogen and four hydrogens that collectively bear a +1 charge. The remaining polyatomic ions are all negatively-charged and, therefore, are classified as anions. However, only two of these, the hydroxide ion and the cyanide ion, are named using the "-ide" suffix that is typically indicative of negatively-charged particles. The remaining polyatomic anions, which all contain oxygen, in combination with another non-metal, exist as part of a series in which the number of oxygens within the polyatomic unit can vary. As has been repeatedly emphasized in several sections of this text, no two mmsanotherstage2019.comical formulas should share a common mmsanotherstage2019.comical name. A single suffix, "-ide," is insufficient for distinguishing the names of the anions in a related polyatomic series. Therefore, "-ate" and "-ite" suffixes are employed, in order to denote that the corresponding polyatomic ions are part of a series. Additionally, these suffixes also indicate the relative number of oxygens that are contained within the polyatomic ions. Note that all of the polyatomic ions whose names end in "-ate" contain one more oxygen than those polyatomic anions whose names end in "-ite." Unfortunately, much like the common system for naming transition metals, these suffixes only indicate the relative number of oxygens that are contained within the polyatomic ions. For example, the nitrate ion, which is symbolized as NO3−1, has one more oxygen than the nitrite ion, which is symbolized as NO2−1. However, the sulfate ion is symbolized as SO4−2. While both the nitrate ion and the sulfate ion share an "-ate" suffix, the former contains three oxygens, but the latter contains four. Additionally, both the nitrate ion and the sulfite ion contain three oxygens, but these polyatomic ions do not share a common suffix. Unfortunately, the relative nature of these suffixes mandates that the ion formula/ion name combinations of the polyatomic ions must simply be memorized.

The rule for constructing formulas for ionic compounds containing polyatomic ions is the same as for formulas containing monatomic (single-atom) ions: the positive and negative charges must balance. If more than one of a particular polyatomic ion is needed to balance the charge, the entire formula for the polyatomic ion must be enclosed in parentheses, and the numerical subscript is placed outside the parentheses. This is to show that the subscript applies to the entire polyatomic ion. Two examples are shown below:

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Mg2+ and I− Na+ and O2−

4. Write the mmsanotherstage2019.comical formula for the ionic compound formed by each pair of ions.