Binary Ionic Compounds

An ionic compound that contains only two elements, one present as a cation and one as an anion, is called a binary ionic compoundAn ionic compound that contains only two elements, one present as a cation and one as an anion.. One example is MgCl_2, a coagulant used in the preparation of tofu from soybeans. For binary ionic compounds, the subscripts in the empirical formula can also be obtained by crossing charges: use the absolute value of the charge on one ion as the subscript for the other ion. This method is shown schematically as follows:

Crossing charges. One method for obtaining subscripts in the empirical formula is by crossing charges.

When crossing charges, you will sometimes find it necessary to reduce the subscripts to their simplest ratio to write the empirical formula. Consider, for example, the compound formed by Mg²⁺ and O^2-. Using the absolute values of the charges on the ions as subscripts gives the formula Mg₂O₂:

This simplifies to its correct empirical formula MgO. The empirical formula has one Mg²⁺ ion and one O^2- ion.

Polyatomic Ions

Polyatomic ionsA group of two or more atoms that has a net electrical charge. are groups of atoms that bear a net electrical charge, although the atoms in a polyatomic ion are held together by the same covalent bonds that hold atoms together in molecules. Just as there are many more kinds of molecules than simple elements, there are many more kinds of polyatomic ions than monatomic ions. Two examples of polyatomic cations are the ammonium (NH₄⁺) and the methylammonium (CH₃NH₃⁺) ions. Polyatomic anions are much more numerous than polyatomic cations; some common examples are in Table 2.4 "Common Polyatomic Ions and Their Names".

As you acquaint yourself with the discipline of chemistry, these names, formulas and charges in table 2.4 will become second nature. Until then, you need to get a toe hold. A suggestion is to start with the oxoacids which are coming up in a future section. Consider the Table 2.4b below:

Table 2.4b

See if you can sit down with a blank piece of paper and from memory write down the first two columns which represent the names and formulas of the important oxoacids. Once that part is well in hand, then try adding the remaining columns of Table 2.4b. Once that's done, catch the five stray species highlighted in yellow in Table 2.4 above. At this point it's brute force memorization, but just watch how often these ions come up in the remainder of the textbook and in your laboratory activities. Soon you'll be a pro!

The method we used to predict the empirical formulas for ionic compounds that contain monatomic ions can also be used for compounds that contain polyatomic ions. The overall charge on the cations must balance the overall charge on the anions in the formula unit. Thus K⁺ and NO₃^- ions combine in a 1:1 ratio to form KNO₃ (potassium nitrate or saltpeter), a major ingredient in black gunpowder. Similarly, Ca²⁺ and SO₄^2- form CaSO₄ (calcium sulfate), which combines with varying amounts of water to form gypsum and plaster of Paris. The polyatomic ions NH₄⁺ and NO₃^- form NH₄NO₃ (ammonium nitrate), which is a widely used fertilizer and, in the wrong hands, an explosive. One example of a compound in which the ions have charges of different magnitudes is calcium phosphate, which is composed of Ca²⁺ and PO₄^3- ions; it is a major component of bones. The compound is electrically neutral because the ions combine in a ratio of three Ca²⁺ ions [3(+2) = +6] for every two ions [2(-3) = -6], giving an empirical formula of Ca₃(PO₄)₂; the parentheses around PO₄ in the empirical formula indicate that it is a polyatomic ion. Writing the formula for calcium phosphate as Ca₃P₂O₈ gives the correct number of each atom in the formula unit, but it obscures the fact that the compound contains readily identifiable PO₄^3- ions.

Hydrates

Many ionic compounds occur as hydratesA compound that contains specific ratios of loosely bound water molecules, called waters of hydration., compounds that contain specific ratios of loosely bound water molecules, called waters of hydrationThe loosely bound water molecules in hydrate compounds. These waters of hydration can often be removed by simply heating the compound.. Waters of hydration can often be removed simply by heating. For example, calcium dihydrogen phosphate can form a solid that contains one molecule of water per Ca(H₂PO₄)₂ unit and is used as a leavening agent in the food industry to cause baked goods to rise. The empirical formula for the solid is Ca(H₂PO₄)₂·H₂O. In contrast, copper sulfate usually forms a blue solid that contains five waters of hydration per formula unit, with the empirical formula CuSO₄·5H₂O. When heated, all five water molecules are lost, giving a white solid with the empirical formula CuSO₄ (Figure 2.9 "Hydration of Anhydrous Copper(II) Sulfate").

Figure 2.9 Hydration of Anhydrous Copper(II) Sulfate

When a drop of water is placed on white CuSO₄ blue CuSO₄·5H₂O is formed . Image Credit: By Benjah-bmm27 (Own work) [Public domain], via Wikimedia Commons

Compounds that differ only in the numbers of waters of hydration can have very different properties. For example, CaSO₄·H₂O is plaster of Paris, which was often used to make sturdy casts for broken arms or legs, whereas CaSO₄·2H₂O is the less dense, flakier gypsum, a mineral used in drywall panels for home construction. When a cast would set, a mixture of plaster of Paris and water crystallized to give solid CaSO₄·2H₂O. Similar processes are used in the setting of cement and concrete.