3.4: Polyatomic Ions and Formulae for Ionic Compounds

Learning Objectives

• Recognize polyatomic ions in chemical formulas.
• Write the chemical formula for an ionic compound and name them.

We have already encountered some chemical formulas for ionic compounds from monoatomic ions. A chemical formula is a concise list of the elements in a compound and the ratios of these elements. To better understand what a chemical 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 chemical 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.

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.

Polyatomic Ions

Some ions consist of groups of atoms covalently bonded together and have an overall electric charge. Because these ions contain more than one atom, they are called polyatomic ions. The Lewis structures, names and formulas of some polyatomic ions are found in Figure 3.4.1.

Figure \(\PageIndex{1}\): Names, formula, and charges of polyatomic ions

Polyatomic ions have defined formulas, names, and charges that cannot be modified in any way. Table \(\PageIndex{1}\) 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.

Table \(\PageIndex{1}\): Ion Names and Ion Formulas of Common Polyatomic Ions

Ion Name	Ion Formula
ammonium ion	NH4^+1
hydronium ion	H3O^+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 there are two polyatomic ions in this table, the ammonium ion and hydronium ion that are cations. The ammonium ion contains one nitrogen and four hydrogen's that collectively bear a +1 charge. The hydronium ion contains one oxygen and three hydrogen's 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 oxygen's within the polyatomic unit can vary. No two chemical formulas should share a common chemical 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.

Most of the polyatomic ions in table 3.4.1 are found in significant amounts in all living things and required for survival. For example, the hydrogen carbonate ion (HCO₃^-) is involved in the transport of carbon dioxide (CO₂) from the tissues to the lungs.

Naming Compounds

Now that we know how to name ions, we are ready to name ionic compounds. We do so by placing the name of the cation first, followed by the name of the anion, and dropping the word ion from both parts. For example, what is the name of the compound whose formula is \(\ce{Ba(NO3)2}\)?

The compound’s name does not indicate that there are two nitrate ions for every barium ion. You must determine the relative numbers of ions by balancing the positive and negative charges. If you are given a formula for an ionic compound whose cation can have more than one possible charge, you must first determine the charge on the cation before identifying its correct name. For example, consider \(\ce{FeCl2}\) and \(\ce{FeCl3}\). In the first compound, the iron ion has a 2+ charge because there are two \(\ce{Cl^{−}}\) ions in the formula (1− charge on each chloride ion). In the second compound, the iron ion has a 3+ charge, as indicated by the three \(\ce{Cl^{−}}\) ions in the formula. These are two different compounds that need two different names. By the Stock system, the names are iron(II) chloride and iron(III) chloride. If we were to use the stems and suffixes of the common system, the names would be ferrous chloride and ferric chloride, respectively.

Formulae for Ionic Compounds from Polyatomic Ions

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:

  

The image asks a two-part question: Write the formula for the compound formed by (a) barium and nitrate, and (b) ammonium and phosphate. For (a), Ba has a charge of 2+ and NO₃ has a charge of -1. By crossing charges, the formula is Ba(NO₃)₂. For (b), NH₄ has a charge of +1 and PO₄ has a charge of 3-, so by crossing charges the formula is (NH₄)₃PO₄.

Example \(\PageIndex{1}\)

Write the chemical formula for an ionic compound composed of each pair of ions.

1. the potassium ion and the sulfate ion
2. the calcium ion and the nitrate ion

Solution

1. Potassium ions have a charge of 1+, while sulfate ions have a charge of 2−. We will need two potassium ions to balance the charge on the sulfate ion, so the proper chemical formula is \(\ce{K_2SO_4}\).
2. Calcium ions have a charge of 2+, while nitrate ions have a charge of 1−. We will need two nitrate ions to balance the charge on each calcium ion. The formula for nitrate must be enclosed in parentheses. Thus, we write \(\ce{Ca(NO3)2}\) as the formula for this ionic compound.

Exercise \(\PageIndex{1}\)

Write the chemical formula for an ionic compound composed of each pair of ions.

1. the magnesium ion and the carbonate ion
2. the aluminum ion and the acetate ion

Answer a:

Mg²⁺ and CO₃^2- = MgCO₃

Answer b:

Al³⁺ and C₂H₃O₂^- = Al(C₂H₃O₂)₃

Recognizing Ionic Compounds

There are two ways to recognize ionic compounds. First, compounds between metal and nonmetal elements are usually ionic. For example, CaBr₂ contains a metallic element (calcium, a group 2A metal) and a nonmetallic element (bromine, a group 7A nonmetal). Therefore, it is an ionic compound. In contrast, the compound NO₂ contains two elements that are both nonmetals (nitrogen, from group 5A, and oxygen, from group 6A). It is not an ionic compound; it belongs to the category of covalent compounds also know as molecular compounds. Also note that this combination of nitrogen and oxygen has no electric charge specified, so it is not the nitrite ion.

Second, if you recognize the formula of a polyatomic ion in a compound, the compound is ionic. For example, if you see the formula \(\ce{Ba(NO3)2}\), you may recognize the “NO₃” part as the nitrate ion, \(\rm{NO_3^−}\). (Remember that the convention for writing formulas for ionic compounds is not to include the ionic charge.) This is a clue that the other part of the formula, \(\ce{Ba}\), is actually the \(\ce{Ba^{2+}}\) ion, with the 2+ charge balancing the overall 2− charge from the two nitrate ions. Thus, this compound is also ionic.

If the compound is not ionic then it is a molecular compound.

Example \(\PageIndex{2}\)

Identify each compound as ionic or molecular (not ionic).

1. \(\ce{Na2O}\)
2. \(\ce{PCl3}\)
3. \(\ce{NH4Cl}\)
4. \(\ce{OF2}\)

Solution

1. Sodium is a metal, and oxygen is a nonmetal; therefore, \(\ce{Na2O}\) is expected to be ionic.
2. Both phosphorus and chlorine are nonmetals. Therefore, \(\ce{PCl3}\) is a molecular compound.
3. The \(\ce{NH4}\) in the formula represents the ammonium ion, \(\ce{NH4^{+}}\), which indicates that this compound is ionic.
4. Both oxygen and fluorine are nonmetals. Therefore, \(\ce{OF2}\) is not ionic.

Exercise \(\PageIndex{2}\)

Identify each compound as ionic or molecular (not ionic).

1. \(\ce{N2O}\)
2. \(\ce{FeCl3}\)
3. \(\ce{(NH4)3PO4}\)
4. \(\ce{SOCl2}\)

Answer a:

molecular

Answer b:

ionic

Answer c:

ionic

Answer d:

molecular

Looking Closer: Blood and Seawater

Science has long recognized that blood and seawater have similar compositions. After all, both liquids have ionic compounds dissolved in them. The similarity may be more than mere coincidence; many scientists think that the first forms of life on Earth arose in the oceans. A closer look, however, shows that blood and seawater are quite different. A 0.9% solution of sodium chloride approximates the salt concentration found in blood. In contrast, seawater is principally a 3% sodium chloride solution, over three times the concentration in blood. Here is a comparison of the amounts of ions in blood and seawater:

Table \(\PageIndex{2}\):  A comparison of the amounts of ions in blood and seawater.

Ion	Percent in Seawater	Percent in Blood
Na^+	2.36	0.322
Cl^−	1.94	0.366
Mg^2^+	0.13	0.002
SO4^2^−	0.09	—
K^+	0.04	0.016
Ca^2^+	0.04	0.0096
HCO3^−	0.002	0.165
HPO4^2^−, H2PO4^−	—	0.01

Most ions are more abundant in seawater than they are in blood, with some important exceptions. There are far more hydrogen carbonate ions (\(\ce{HCO3^{−}}\)) in blood than in seawater. This difference is significant because the hydrogen carbonate ion and some related ions have a crucial role in controlling the acid-base properties of blood. The amount of hydrogen phosphate ions—\(\ce{HPO4^{2−}}\) and \(\ce{H2PO4^{−}}\)—in seawater is very low, but they are present in higher amounts in blood, where they also affect acid-base properties. Another notable difference is that blood does not have significant amounts of the sulfate ion (\(\ce{SO4^{2−}}\)), but this ion is present in seawater.

A synopsis of how to name simple ionic compounds is shown in figure 3.4.2.

Figure \(\PageIndex{2}\): A Guide to Naming Simple Ionic Compounds. Follow these steps to name a simple ionic compound.

Identify the cation name and the anion name. If the cation can have more than one possible charge, either use the Stock system name of the cation and name of the anion, or use the stem of the cation name and -ic/-ous and the name of the anion. Examples of this would be FeCl₂, which is iron(II) chloride of ferrous chloride, CuSO₄, which is copper(II) sulfate or cupric sulfate, and Cr₂O₃, which is chromium(III) oxide or chromic oxide. If the cation can not have more than one possible charge, use the name of the cation and the name of the anion. Examples of this would be KBr, which is potassium bromide, NaNO₃, which is sodium nitrate, and (NH₄)₂S, which is ammonium sulfide.

Key Takeaways

• Proper chemical formulas for ionic compounds balance the total positive charge with the total negative charge.
• Groups of atoms with an overall charge, called polyatomic ions, also exist.

EXERCISES

1. What information is contained in the formula of an ionic compound?

• Why do the chemical formulas for some ionic compounds contain subscripts, while others do not?

3. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Mg²⁺ and I⁻
2. Na⁺ and O²⁻

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

1. Na⁺ and Br⁻
2. Mg²⁺ and Br⁻
3. Mg²⁺ and S²⁻

5. Write the chemical formula for the ionic compound formed by each pair of ions.

1. K⁺ and Cl⁻
2. Mg²⁺ and Cl⁻
3. Mg²⁺ and Se²⁻

6. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Na⁺ and N³⁻
2. Mg²⁺ and N³⁻
3. Al³⁺ and S²⁻

7. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Li⁺ and N³⁻
2. Mg²⁺ and P³⁻
3. Li⁺ and P³⁻

8. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Fe³⁺ and Br⁻
2. Fe²⁺ and Br⁻
3. Au³⁺ and S²⁻
4. Au⁺ and S²⁻

9. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Cr³⁺ and O²⁻
2. Cr²⁺ and O²⁻
3. Pb²⁺ and Cl⁻
4. Pb⁴⁺ and Cl⁻

10. Write the chemical formula for the ionic compound formed by each pair of ions.

1. Cr³⁺ and NO₃⁻
2. Fe²⁺ and PO₄³⁻
3. Ca²⁺ and CrO₄²⁻
4. Al³⁺ and OH⁻

11. Write the chemical formula for the ionic compound formed by each pair of ions.

1. NH₄⁺ and NO₃⁻
2. H⁺ and Cr₂O₇²⁻
3. Cu⁺ and CO₃²⁻
4. Na⁺ and HCO₃⁻

12. For each pair of elements, determine the charge for their ions and write the proper formula for the resulting ionic compound between them.

1. Ba and S
2. Cs and I

13. For each pair of elements, determine the charge for their ions and write the proper formula for the resulting ionic compound between them.

1. K and S
2. Sc and Br

14. Which compounds would you predict to be ionic?

1. Li₂O
2. (NH₄)₂O
3. CO₂
4. FeSO₃
5. C₆H₆
6. C₂H₆O

15. Which compounds would you predict to be ionic?

1. Ba(OH)₂
2. CH₂O
3. NH₂CONH₂
4. (NH₄)₂CrO₄
5. C₈H₁₈
6. NH₃

12. Name the ionic compound formed by each pair of ions. Use both the Stock and common systems, where appropriate.

1. Cr³⁺ and NO₃⁻
2. Fe²⁺ and PO₄³⁻
3. Ca²⁺ and CrO₄²⁻
4. Al³⁺ and OH⁻

13. Name the ionic compound formed by each pair of ions. Use both the Stock and common systems, where appropriate.

1. NH₄⁺ and NO₃⁻
2. K⁺ and Cr₂O₇²⁻
3. Cu⁺ and CO₃²⁻
4. Na⁺ and HCO₃⁻

14. Give two names for each compound.

1. Al(HSO₄)₃
2. Mg(HSO₄)₂

15. Give two names for each compound.

1. Co(HCO₃)₂
2. LiHCO₃

11.

1. chromium(III) oxide or chromic oxide
2. chromium(II) oxide or chromous oxide
3. lead(II) chloride or plumbous chloride
4. lead(IV) chloride or plumbic chloride

12.

1. chromium(III) nitrate or chromic nitrate
2. iron(II) phosphate or ferrous phosphate
3. calcium chromate
4. aluminum hydroxide

13.

1. ammonium nitrate
2. potassium dichromate
3. copper(I) carbonate or cuprous carbonate
4. sodium hydrogen carbonate or sodium bicarbonate

14.

1. aluminum hydrogen sulfate or aluminum bisulfate
2. magnesium hydrogen sulfate or magnesium bisulfate

15.

1. cobalt hydrogen carbonate or cobalt bicarbonate
2. lithium hydrogen carbonate or lithium bicarbonate

 

 

 

Answers

1. the ratio of each kind of ion in the compound

2. Sometimes more than one ion is needed to balance the charge on the other ion in an ionic compound.

3.

1. MgI₂
2. Na₂O

4.

1. NaBr
2. MgBr₂
3. MgS

5.

1. KCl
2. MgCl₂
3. MgSe

6.

1. Na₃N
2. Mg₃N₂
3. Al₂S₃

7.

1. Li₃N
2. Mg₃P₂
3. Li₃P

8.

1. FeBr₃
2. FeBr₂
3. Au₂S₃
4. Au₂S

9.

1. Cr₂O₃
2. CrO
3. PbCl₂
4. PbCl₄

10.

1. Cr(NO₃)₃
2. Fe₃(PO₄)₂
3. CaCrO₄
4. Al(OH)₃

11.

1. NH₄NO₃
2. H₂Cr₂O₇
3. Cu₂CO₃
4. NaHCO₃

12.

1. Ba²⁺, S²⁻, BaS
2. Cs⁺, I⁻, CsI

13.

1. K⁺, S²⁻, K₂S
2. Sc³⁺, Br⁻, ScBr₃

14.

1. ionic
2. ionic
3. not ionic
4. ionic
5. not ionic
6. not ionic

15.

1. ionic
2. not ionic
3. not ionic
4. ionic
5. not ionic
6. not ionic