8.3 Other Types of Chemical Reactions

Acid-Base Reactions

An acid-base reaction is one in which a hydrogen ion, [latex]\ce{H+}[/latex], is transferred from one chemical species to another. Such reactions are of central importance to numerous natural and technological processes, ranging from the chemical transformations that take place within cells and the lakes and oceans, to the industrial-scale production of fertilizers, pharmaceuticals, and other substances essential to society. The subject of acid-base chemistry, therefore, is worthy of thorough discussion, and a full chapter is devoted to this topic later in the text.

For purposes of this brief introduction, we will consider only the more common types of acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that will dissolve in water to yield hydronium ions, [latex]\ce{H3O+}[/latex]. As an example, consider the equation shown here:

[latex]\text{HCl(}aq\text{)}+{\text{H}}_{2}\text{O(}aq\text{)}\rightarrow{\text{Cl}}^{-}\text{(}aq\text{)}+{\text{H}}_{3}{\text{O}}^{\text{+}}\text{(}aq\text{)}[/latex]

The process represented by this equation confirms that hydrogen chloride is an acid. When dissolved in water, [latex]\ce{H3O+}[/latex] ions are produced by a chemical reaction in which [latex]\ce{H+}[/latex] ions are transferred from [latex]\ce{HCl}[/latex] molecules to [latex]\ce{H2O}[/latex] molecules (Figure 8.3.1).

The nature of [latex]\ce{HCl}[/latex] is such that its reaction with water as just described is essentially 100% efficient: Virtually every [latex]\ce{HCl}[/latex] molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and [latex]\ce{HCl}[/latex] is one among just a handful of common acid compounds that are classified as strong (Table 8.3.1).

A far greater number of compounds behave as acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor. A familiar example of a weak acid is acetic acid, the main ingredient in food vinegars:

[latex]{\text{CH}}_{3}{\text{CO}}_{2}\text{H(}aq\text{)}+{\text{H}}_{2}\text{O(}l\text{)}\rightleftharpoons{\text{CH}}_{3}{\text{CO}}_{2}{}^{-}\text{(}aq\text{)}+{\text{H}}_{3}{\text{O}}^{\text{+}}\text{(}aq\text{)}[/latex]

When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, [latex]\ce{CH3CO2-}[/latex] (Figure 8.3.2). (The use of a double-arrow in the equation above denotes the partial reaction aspect of this process, a concept addressed fully in the chapter on chemical equilibrium.)

A base is a substance that will dissolve in water to yield hydroxide ions, [latex]\ce{OH-}[/latex]. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, [latex]\ce{NaOH}[/latex] and [latex]\ce{Ca(OH)2}[/latex]. When these compounds dissolve in water, hydroxide ions are released directly into the solution. For example, [latex]\ce{KOH}[/latex] and [latex]\ce{Ba(OH)2}[/latex] dissolve in water and dissociate completely to produce cations ([latex]\ce{K+}[/latex] and [latex]\ce{Ba^{2+}}[/latex], respectively) and hydroxide ions, [latex]\ce{OH-}[/latex]. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

Consider as an example the dissolution of lye (sodium hydroxide) in water:

[latex]\text{NaOH(}s\text{)}\rightarrow{\text{Na}}^{\text{+}}\text{(}aq\text{)}+{\text{OH}}^{-}\text{(}aq\text{)}[/latex]

This equation confirms that sodium hydroxide is a base. When dissolved in water, [latex]\ce{NaOH}[/latex] dissociates to yield [latex]\ce{Na+}[/latex] and [latex]\ce{OH-}[/latex] ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, [latex]\ce{NaOH}[/latex] and other ionic hydroxides are all classified as strong bases.

Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners (Figure 8.3.3). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:

[latex]{\text{NH}}_{3}\text{(}aq\text{)}+{\text{H}}_{2}\text{O(}l\text{)}\rightleftharpoons{\text{NH}}_{4}{}^{\text{+}}\text{(}aq\text{)}+{\text{OH}}^{-}\text{(}aq\text{)}[/latex]

This is, by definition, an acid-base reaction, in this case involving the transfer of [latex]\ce{H+}[/latex] ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as [latex]{\text{NH}}_{4}{}^{+}[/latex] ions.

A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base, and the products are often a salt and water:

[latex]\text{acid}+\text{base}\rightarrow\text{salt}+\text{water}[/latex]

To illustrate a neutralization reaction, consider what happens when a typical antacid such as milk of magnesia (an aqueous suspension of solid [latex]\ce{Mg(OH)2}[/latex]) is ingested to ease symptoms associated with excess stomach acid ([latex]\ce{HCl}[/latex]):

[latex]\text{Mg}{\text{(OH)}}_{2}\text{(}s\text{)}+2\text{HCl(}aq\text{)}\rightarrow{\text{MgCl}}_{2}\text{(}aq\text{)}+2{\text{H}}_{2}\text{O(}l\text{)}[/latex]

Note that in addition to water, this reaction produces a salt, magnesium chloride.

CHEMISTRY IN EVERYDAY LIFE

Culinary Aspects of Chemistry

Examples of acid-base chemistry are abundant in the culinary world. One example is the use of baking soda, or sodium bicarbonate in baking. [latex]\ce{NaHCO3}[/latex] is a base. When it reacts with an acid such as lemon juice, buttermilk, or sour cream in a batter, bubbles of carbon dioxide gas are formed from decomposition of the resulting carbonic acid, and the batter “rises.” Baking powder is a combination of sodium bicarbonate, and one or more acid salts that react when the two chemicals come in contact with water in the batter.

Many people like to put lemon juice or vinegar, both of which are acids, on cooked fish (Figure 8.3.4). It turns out that fish have volatile amines (bases) in their systems, which are neutralized by the acids to yield involatile ammonium salts. This reduces the odor of the fish, and also adds a “sour” taste that we seem to enjoy.

Pickling is a method used to preserve vegetables using a naturally produced acidic environment. The vegetable, such as a cucumber, is placed in a sealed jar submerged in a brine solution. The brine solution favors the growth of beneficial bacteria and suppresses the growth of harmful bacteria. The beneficial bacteria feed on starches in the cucumber and produce lactic acid as a waste product in a process called fermentation. The lactic acid eventually increases the acidity of the brine to a level that kills any harmful bacteria, which require a basic environment. Without the harmful bacteria consuming the cucumbers they are able to last much longer than if they were unprotected. A byproduct of the pickling process changes the flavor of the vegetables with the acid making them taste sour.

Oxidation-Reduction Reactions

Earth’s atmosphere contains about 20% molecular oxygen, [latex]\ce{O2}[/latex], a chemically reactive gas that plays an essential role in the metabolism of aerobic organisms and in many environmental processes that shape the world. The term oxidation was originally used to describe chemical reactions involving [latex]\ce{O2}[/latex], but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction (redox) reactions. A few examples of such reactions will be used to develop a clear picture of this classification.

Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride:

[latex]2\text{Na}(s)+\text{Cl}_2(g)\rightarrow{2}\text{NaCl}(s)[/latex]

It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction:

[latex]\begin{array}{l}2\text{Na(}s\text{)}\rightarrow 2{\text{Na}}^{\text{+}}\text{(}s\text{)}+2{\text{e}}^{-}\\ {\text{Cl}}_{2}\text{(}g\text{)}+2{\text{e}}^{-}\rightarrow 2{\text{Cl}}^{-}\text{(}s\text{)}\end{array}[/latex]

These equations show that Na atoms lose electrons while [latex]\ce{Cl}[/latex] atoms (in the [latex]\ce{Cl2}[/latex] molecule) gain electrons, the “s” subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur:

oxidation = loss of electrons

reduction = gain of electrons

In this reaction, then, sodium is oxidized and chlorine is undergoes reduction. Viewed from a more active perspective, sodium functions as a reducing agent (reductant), since it provides electrons to (or reduces) chlorine. Likewise, chlorine functions as an oxidizing agent (oxidant), as it effectively removes electrons from (oxidizes) sodium.

reducing agent = species that is oxidized

oxidizing agent = species that is reduced

Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding [latex]\ce{NaCl}[/latex]:

[latex]{\text{H}}_{2}\text{(}g\text{)}+{\text{Cl}}_{2}\text{(}g\text{)}\rightarrow 2\text{HCl(}g\text{)}[/latex]

The product of this reaction is a covalent compound, so transfer of electrons in the explicit sense is not involved. To clarify the similarity of this reaction to the previous one and permit an unambiguous definition of redox reactions, a property called oxidation number has been defined. The oxidation number (or oxidation state) of an element in a compound is the charge its atoms would possess if the compound was ionic. The following guidelines are used to assign oxidation numbers to each element in a molecule or ion.

1. The oxidation number of an atom in an elemental substance is zero.
2. The oxidation number of a monatomic ion is equal to the ion’s charge.
3. Oxidation numbers for common nonmetals are usually assigned as follows:
   • Hydrogen: +1 when combined with nonmetals, –1 when combined with metals
   • Oxygen: –2 in most compounds, sometimes –1 (so-called peroxides, [latex]{\text{O}}_{2}{}^{\text{2}-}\text{),}[/latex] very rarely [latex]-\frac{1}{2}[/latex] (so-called superoxides, [latex]{\text{O}}_{2}{}^{-}\text{),}[/latex] positive values when combined with F (values vary)
   • Halogens: –1 for F always, –1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values)
4. The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion.

Note: The proper convention for reporting charge is to write the number first, followed by the sign (e.g., 2+), while oxidation number is written with the reversed sequence, sign followed by number (e.g., +2). This convention aims to emphasize the distinction between these two related properties.

Using the oxidation number concept, an all-inclusive definition of redox reaction has been established. Oxidation-reduction (redox) reactions are those in which one or more elements involved undergo a change in oxidation number. (While the vast majority of redox reactions involve changes in oxidation number for two or more elements, a few interesting exceptions to this rule do exist, as in Example 4.) Definitions for the complementary processes of this reaction class are correspondingly revised as shown here:

[latex]\begin{array}{lll}\hfill \mathbf{\text{oxidation}}& =& \text{increase in oxidation number}\hfill \\ \hfill \mathbf{\text{reduction}}& =& \text{decrease in oxidation number}\hfill \end{array}[/latex]

Returning to the reactions used to introduce this topic, they may now both be identified as redox processes. In the reaction between sodium and chlorine to yield sodium chloride, sodium is oxidized (its oxidation number increases from 0 in Na to +1 in [latex]\ce{NaCl}[/latex]) and chlorine is reduced (its oxidation number decreases from 0 in [latex]\ce{Cl2}[/latex] to –1 in [latex]\ce{NaCl}[/latex]). In the reaction between molecular hydrogen and chlorine, hydrogen is oxidized (its oxidation number increases from 0 in [latex]\ce{H2}[/latex] to +1 in [latex]\ce{HCl}[/latex]) and chlorine is reduced (its oxidation number decreases from 0 in [latex]\ce{Cl2}[/latex] to –1 in [latex]\ce{HCl}[/latex]).

Several subclasses of redox reactions are recognized, including combustion reactions in which the reductant (also called a fuel) and oxidant (often, but not necessarily, molecular oxygen) react vigorously and produce significant amounts of heat, and often light, in the form of a flame. Solid rocket-fuel reactions are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation:

[latex]10\text{Al(}s\text{)}+6{\text{NH}}_{4}{\text{ClO}}_{4}\text{(}s\text{)}\rightarrow 4{\text{Al}}_{2}{\text{O}}_{3}\text{(}s\text{)}+2{\text{AlCl}}_{3}\text{(}s\text{)}+12{\text{H}}_{2}\text{O(}g\text{)}+3{\text{N}}_{2}\text{(}g\text{)}[/latex]

Single-displacement (replacement) reactions are redox reactions in which an ion in solution is displaced (or replaced) via the oxidation of a metallic element. One common example of this type of reaction is the acid oxidation of certain metals:

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

Metallic elements may also be oxidized by solutions of other metal salts; for example:

[latex]\text{Cu(}s\text{)}+2{\text{AgNO}}_{3}\text{(}aq\text{)}\rightarrow\text{Cu}{\text{(}{\text{NO}}_{3}\text{)}}_{2}\text{(}aq\text{)}+2\text{Ag(}s\text{)}[/latex]

This reaction may be observed by placing copper wire in a solution containing a dissolved silver salt. Silver ions in solution are reduced to elemental silver at the surface of the copper wire, and the resulting [latex]\ce{Cu^2+}[/latex] ions dissolve in the solution to yield a characteristic blue color (Figure 8.3.5).

Balancing Redox Reactions via the Half-Reaction Method

Redox reactions that take place in aqueous media often involve water, hydronium ions, and hydroxide ions as reactants or products. Although these species are not oxidized or reduced, they do participate in chemical change in other ways (e.g., by providing the elements required to form oxyanions). Equations representing these reactions are sometimes very difficult to balance by inspection, so systematic approaches have been developed to assist in the process. One very useful approach is to use the method of half-reactions, which involves the following steps:

1. Write the two half-reactions representing the redox process.
2. Balance all elements except oxygen and hydrogen.
3. Balance oxygen atoms by adding [latex]\ce{H2O}[/latex] molecules.
4. Balance hydrogen atoms by adding [latex]\ce{H+}[/latex] ions.
5. Balance charge by adding electrons.
6. If necessary, multiply each half-reaction’s coefficients by the smallest possible integers to yield equal numbers of electrons in each.
7. Add the balanced half-reactions together and simplify by removing species that appear on both sides of the equation.
8. For reactions occurring in basic media (excess hydroxide ions), carry out these additional steps:
   • Add [latex]\ce{OH-}[/latex] ions to both sides of the equation in numbers equal to the number of [latex]\ce{H+}[/latex] ions.
   • On the side of the equation containing both [latex]\ce{H+}[/latex] and [latex]\ce{OH-}[/latex] ions, combine these ions to yield water molecules.
   • Simplify the equation by removing any redundant water molecules.
9. Finally, check to see that both the number of atoms and the total charges are balanced.

Glossary

acid: substance that produces [latex]\ce{H3O+}[/latex] when dissolved in water

acid-base reaction: reaction involving the transfer of a hydrogen ion between reactant species

base: substance that produces [latex]\ce{​OH-}[/latex] when dissolved in water

complete ionic equation: chemical equation in which all dissolved ionic reactants and products, including spectator ions, are explicitly represented by formulas for their dissociated ions

combustion reaction: vigorous redox reaction producing significant amounts of energy in the form of heat and, sometimes, light

half-reaction: an equation that shows whether each reactant loses or gains electrons in a reaction.

insoluble: of relatively low solubility; dissolving only to a slight extent

molecular equation: chemical equation in which all reactants and products are represented as neutral substances

neutralization reaction: reaction between an acid and a base to produce salt and water

net ionic equation: chemical equation in which only those dissolved ionic reactants and products that undergo a chemical or physical change are represented (excludes spectator ions)

oxidation: process in which an element’s oxidation number is increased by loss of electrons

oxidation-reduction reaction: (also, redox reaction) reaction involving a change in oxidation number for one or more reactant elements

oxidation number: (also, oxidation state) the charge each atom of an element would have in a compound if the compound were ionic

oxidizing agent: (also, oxidant) substance that brings about the oxidation of another substance, and in the process becomes reduced

precipitate: insoluble product that forms from reaction of soluble reactants

precipitation reaction: reaction that produces one or more insoluble products; when reactants are ionic compounds, sometimes called double-displacement or metathesis

reactant: substance undergoing a chemical or physical change; shown on the left side of the arrow in a chemical equation

reduction: process in which an element’s oxidation number is decreased by gain of electrons

reducing agent: (also, reductant) substance that brings about the reduction of another substance, and in the process becomes oxidized

salt: ionic compound that can be formed by the reaction of an acid with a base that contains a cation and an anion other than hydroxide or oxide

single-displacement reaction: (also, replacement) redox reaction involving the oxidation of an elemental substance by an ionic species

soluble: of relatively high solubility; dissolving to a relatively large extent

solubility: the extent to which a substance may be dissolved in water, or any solvent

spectator ion: ion that does not undergo a chemical or physical change during a reaction, but its presence is required to maintain charge neutrality

strong acid: acid that reacts completely when dissolved in water to yield hydronium ions

strong base: base that reacts completely when dissolved in water to yield hydroxide ions

weak acid: acid that reacts only to a slight extent when dissolved in water to yield hydronium ions

weak base: base that reacts only to a slight extent when dissolved in water to yield hydroxide ions