Up to now, we have presented chemical reactions as a topic, but we have not discussed how the products of a chemical reaction can be predicted. Here, we will begin our study of certain types of chemical reactions that allow us to predict what the products of the reaction will be.

Single-Replacement Reactions

A single-replacement reaction is a chemical reaction in which one element is substituted for another element in a compound, generating a new element and a new compound as products. For example:

[latex]\ce{2HCl}{(aq)} + \ce{Zn}{(s)} → \ce{ZnCl}_{2}{(aq)} + \ce{H}_{2}{(g)}[/latex]

This is an example of a single-replacement reaction. The hydrogen atoms in [latex]\ce{HCl}[/latex] are replaced by [latex]\ce{Zn}[/latex] atoms, and in the process, a new element—hydrogen—is formed. Another example of a single-replacement reaction is:

[latex]\ce{2NaCl}{(aq)} + \ce{F}_{2}{(g)} → \ce{2NaF}{(s)} + \ce{Cl}_{2}{(g)}[/latex]

Here, the negatively charged ion changes from chloride to fluoride. A typical characteristic of a single-replacement reaction is that there is one element as a reactant and another element as a product.

Not all proposed single-replacement reactions will occur between two given reactants. This is most easily demonstrated with fluorine, chlorine, bromine, and iodine. Collectively, these elements are called halogens and are in the next-to-last column on the periodic table (see Figure 4.3.1 “Halogens on the Periodic Table”). The elements on top of the column will replace the elements below them on the periodic table, but not the other way around. Thus, the reaction represented by:

[latex]\ce{CaI}_{2}{(s)} + \ce{Cl}_{2}{(g)} → \ce{CaCl}_{2}{(s)} + \ce{I}_{2}{(s)}[/latex]

This reaction will occur, but the reaction,

[latex]\ce{CaF}_{2}{(s)} + \ce{Br}_{2}{(ℓ)} → \ce{CaBr}_{2}{(s)} + \ce{F}_{2}{(g)}[/latex]

will not because bromine is below fluorine on the periodic table. This is just one of many ways the periodic table helps us understand chemistry.

Chemical reactivity trends are easy to predict when replacing anions in simple ionic compounds—simply use their relative positions on the periodic table. However, when replacing the cations, the trends are not as straightforward. This is partly because there are so many elements that can form cations; an element in one column on the periodic table may replace another element nearby, or it may not. A list called the activity series does the same thing the periodic table does for halogens: it lists the elements that will replace other elements below them in single-replacement reactions. A simple activity series is shown below:

Using an activity series is similar to using the positions of the halogens on the periodic table. An element on top will replace an element below it in compounds undergoing a single-replacement reaction. Elements will not replace elements above them in compounds.

Double-Replacement Reactions

A double-replacement reaction occurs when parts of two ionic compounds are exchanged, making two new compounds. A characteristic of a double-replacement equation is that there are two compounds as reactants and two different compounds as products. An example is:

[latex]\ce{CuCl}_{2}{(aq)} + \ce{2AgNO}_{3}{(aq)} → \ce{Cu(NO_{3})}_{2}{(aq)} + \ce{2AgCl}{(s)}[/latex]

There are two equivalent ways of considering a double-replacement equation: either the cations are swapped, or the anions are swapped. You cannot swap both; you would end up with the same substances you started with. Either perspective should allow you to predict the proper products as long as you pair a cation with an anion and not a cation with a cation or an anion with an anion.

Precipitation Reactions and Solubility Rules

Predicting whether a double-replacement reaction will occur is somewhat more difficult than predicting a single-replacement reaction. However, there is one type of double-replacement reaction that we can predict: the precipitation reaction. A precipitation reaction occurs when two ionic compounds are dissolved in water and form a new ionic compound that does not dissolve; this new compound falls out of the solution as a solid precipitate. The formation of a solid precipitate is the driving force that makes the reaction proceed.

To judge whether double-replacement reactions will occur, we need to know what kinds of ionic compounds form precipitates. For this, we use solubility rules, which are general statements that predict which ionic compounds dissolve (are soluble) and which do not (are not soluble or insoluble). Tables 4.3.1 and 4.3.2 list some general solubility rules. These rules highlight that we need to consider each ionic compound (both the reactants and the possible products) in light of the solubility rules in Tables 4.3.1 and 4.3.2. If a compound is soluble, we use the (aq) label with it, indicating that it dissolves. If a compound is not soluble, we use the (s) label with it and assume that it will precipitate out of the solution. If all possible products are soluble in a reaction, then no precipitation reaction will be expected.

To illustrate the rules mentioned in the previous paragraph:

• consider the possible double-replacement reaction between [latex]\ce{Na}_{2}\ce{SO}_{4}[/latex] and [latex]\ce{SrCl}_{2}[/latex]. The solubility rules say that all ionic sodium compounds are soluble, and all ionic chloride compounds are soluble except for [latex]\ce{Ag}^{+}[/latex], [latex]\ce{Hg}^{2+}[/latex], and [latex]\ce{Pb}^{2+}[/latex], which are not being considered here. Therefore, [latex]\ce{Na}_{2}\ce{SO}_{4}[/latex] and [latex]\ce{SrCl}_{2}[/latex] are both soluble.
• The possible double-replacement reaction products are [latex]\ce{NaCl}[/latex] and [latex]\ce{SrSO}_{4}[/latex]. Are these soluble? [latex]\ce{NaCl}[/latex] is (by the same rule we just quoted), but what about [latex]\ce{SrSO}_{4}[/latex]? Compounds of the sulphate ion are generally soluble, but [latex]\ce{Sr}^{2+}[/latex] is an exception: we expect it to be insoluble—a precipitate. Therefore, we expect a reaction to occur, and the balanced chemical equation would be:

[latex]\ce{Na}_{2}\ce{SO}_{4}{(aq)} + \ce{SrCl}_{2}{(aq)} → \ce{2NaCl}{(aq)} + \ce{SrSO}_{4}{(s)}[/latex]

• You would expect to see a visual change corresponding to [latex]\ce{SrSO}_{4}[/latex]precipitating out of solution (Figure 4.3.2 “Double-Replacement Reactions”).