Classifying Chemical Reactions (7.2)

Humans interact with one another in various and complex ways, and we classify these interactions according to common patterns of behavior. When two humans exchange information, we say they are communicating. When they exchange blows with their fists or feet, we say they are fighting. Faced with a wide range of varied interactions between chemical substances, scientists have likewise found it convenient (or even necessary) to classify chemical interactions by identifying common patterns of reactivity. This module will provide an introduction to three of the most prevalent types of chemical reactions: precipitation, acid-base, and oxidation-reduction.

Precipitation Reactions and Solubility Rules

A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solution and are sometimes referred to as double displacement, double replacement, or metathesis reactions. These reactions are common in nature and are responsible for the formation of coral reefs in ocean waters and kidney stones in animals. They are used widely in industry for the production of a number of commodity and specialty chemicals. Precipitation reactions also play a central role in many chemical analysis techniques, including spot tests used to identify metal ions and gravimetric methods for determining the composition of matter (see the last module of this chapter).

The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility, defined as the maximum concentration of a substance that can be achieved under specified conditions. Substances with relatively large solubilities are said to be soluble. A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble, and these are the substances that readily precipitate from solution. More information on these important concepts is provided in a later chapter on solutions. For purposes of predicting the identities of solids formed by precipitation reactions, one may simply refer to patterns of solubility that have been observed for many ionic compounds (Table 7.8).

A vivid example of precipitation is observed when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide:

$$2\text{KI}(aq)+\text{Pb}{({\text{NO}}_3)}_2(aq)⟶{\text{PbI}}_2(s)+2{\text{KNO}}_3(aq)$$

This observation is consistent with the solubility guidelines: The only insoluble compound among all those involved is lead iodide, one of the exceptions to the general solubility of iodide salts.

The net ionic equation representing this reaction is:

$${\text{Pb}}^{\mathrm{2+}}(aq)+2{\text{I}}^{\text{−}}(aq)⟶{\text{PbI}}_2(s)$$

Lead iodide is a bright-yellow solid that was formerly used as an artist’s pigment known as iodine yellow (Figure 7.4). The properties of pure PbI₂ crystals make them useful for the fabrication of X-ray and gamma-ray detectors.

The solubility guidelines in Table 7.8 may be used to predict whether a precipitation reaction will occur when solutions of soluble ionic compounds are mixed together. One merely needs to identify all the ions present in the solution and then consider if possible cation/anion pairing could result in an insoluble compound. For example, mixing solutions of silver nitrate and sodium fluoride will yield a solution containing Ag⁺, NO₃ ⁻, Na⁺, and F⁻ ions. Aside from the two ionic compounds originally present in the solutions, AgNO₃ and NaF, two additional ionic compounds may be derived from this collection of ions: NaNO₃ and AgF. The solubility guidelines indicate all nitrate salts are soluble but that AgF is one of the exceptions to the general solubility of fluoride salts. A precipitation reaction, therefore, is predicted to occur, as described by the following equations:

$$\begin{array}{c}\text{NaCl}(aq)+{\text{AgNO}}_3(aq)⟶\text{AgCl}(s)+\text{NaN}{\text{O}}_3(aq)\text{(molecular)}\\ {\text{Ag}}^{\text{+}}(aq)+{\text{Cl}}^{\text{−}}(aq)⟶\text{AgCl}(s)\text{(net ionic)}\end{array}$$

Acid-Base Reactions

An acid-base reaction is one in which a hydrogen ion, H+, 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, H₃O+. As an example, consider the equation shown here:

$$\text{HCl}(aq)+{\text{H}}_2\text{O}(aq)⟶{\text{Cl}}^{\text{−}}(aq)+{\text{H}}_3{\text{O}}^{\text{+}}(aq)$$

The process represented by this equation confirms that hydrogen chloride is an acid. When dissolved in water, H₃O+ ions are produced by a chemical reaction in which H+ ions are transferred from HCl molecules to H₂O molecules (Figure 7.5).

The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table 7.2). A far greater number of compounds behave as weak 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:

$${\text{CH}}_3{\text{CO}}_2\text{H}(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{CH}}_3{\text{CO}}_2{}^{\text{−}}(aq)+{\text{H}}_3{\text{O}}^{\text{+}}(aq)$$

When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form, CH₃CO₂ − (Figure 7.6). (The use of a double-arrow in the equation above denotes the partial reaction aspect of this process, a concept addressed fully in the chapters on chemical equilibrium.)

A base is a substance that will dissolve in water to yield hydroxide ions, OH−. 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, NaOH and Ca(OH)₂. Unlike the acid compounds discussed previously, these compounds do not react chemically with water; instead, they dissolve and dissociate, releasing hydroxide ions directly into the solution. For example, KOH and Ba(OH)₂ dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH−. 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:

$$\text{NaOH}(s)⟶{\text{Na}}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq)$$

This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na+ and OH− 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, NaOH 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 the chemical synthesis of other compounds and an active ingredient in household cleaners (Figure 7.7). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:

$${\text{NH}}_3(aq)+{\text{H}}_2\text{O}(l)\rightleftharpoons {\text{NH}}_4{}^{\text{+}}(aq)+{\text{OH}}^{\text{−}}(aq)$$

This is, by definition, an acid-base reaction, in this case involving the transfer of H+ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as NH₄+ ions.

A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base (but not water), and the products are often a salt and water:

$$\text{acid}+\text{base}⟶\text{salt}+\text{water}$$

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

$$\text{Mg}{\text{(OH)}}_2(s)+2\text{HCl}(aq)⟶{\text{MgCl}}_2(aq)+2{\text{H}}_2\text{O}(l\text{).}$$

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. NaHCO₃ 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 the 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 7.8). 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, O₂, 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 O₂, 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:

$$\text{2Na}(s)+{\text{Cl}}_2(g)⟶2\text{NaCl}(s)$$

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:

$$\begin{array}{l}2\text{Na}(s)⟶2{\text{Na}}^{\text{+}}(s)+2{\text{e}}^{-}\\ {\text{Cl}}_2(g)+2{\text{e}}^{-}⟶2{\text{Cl}}^{\text{−}}(s)\end{array}$$

These equations show that Na atoms lose electrons while Cl atoms (in the Cl₂ 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:

$$\begin{array}{lll}\hfill \mathbf{\text{oxidation}}& =& \text{loss of electrons}\hfill \\ \hfill \mathbf{\text{reduction}}& =& \text{gain of electrons}\hfill \end{array}$$

In this reaction, then, sodium is oxidized, and chlorine 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.

$$\begin{array}{lll}\hfill \mathbf{\text{reducing agent}}& =& \text{species that is oxidized}\hfill \\ \hfill \mathbf{\text{oxidizing agent}}& =& \text{species that is reduced}\hfill \end{array}$$

Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding NaCl:

$${\text{H}}_2(g)+{\text{Cl}}_2(g)⟶2\text{HCl}(g)$$

The product of this reaction is a covalent compound, so the 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.

• The oxidation number of an atom in an elemental substance is zero.
• The oxidation number of a monatomic ion is equal to the ion’s charge.
• 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, O₂ ²⁻), very rarely − [latex]\frac{1}{2}[/latex] (so-called superoxides, O₂ ^–) 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)
• 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 the oxidation number is written with the reversed sequence, the sign followed by the 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 Example 7.6.) Definitions for the complementary processes of this reaction class are correspondingly revised as shown here:

$$\begin{array}{lll}\hfill \mathbf{\text{oxidation}}& =& \text{increase in oxidation number}\hfill \\ \hfill \mathbf{\text{reduction}}& =& \text{decrease in oxidation number}\hfill \end{array}$$

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 NaCl), and chlorine is reduced (its oxidation number decreases from 0 in Cl₂ to −1 in NaCl). In the reaction between molecular hydrogen and chlorine, hydrogen is oxidized (its oxidation number increases from 0 in H₂ to +1 in HCl), and chlorine is reduced (its oxidation number decreases from 0 in Cl₂ to −1 in HCl).

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 such as the one depicted in Figure 7.1 are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation:

$$10\text{Al}(s)+6{\text{NH}}_4{\text{ClO}}_4(s)⟶4{\text{Al}}_2{\text{O}}_3(s)+2{\text{AlCl}}_3(s)+12{\text{H}}_2\text{O}(g)+3{\text{N}}_2(g)$$

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:

$$\text{Zn}(s)+2\text{HCl}(aq)⟶{\text{ZnCl}}_2(aq)+{\text{H}}_2(g)$$

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

$$\text{Cu}(s)+2{\text{AgNO}}_3(aq)⟶\text{Cu}{({\text{NO}}_3)}_2(aq)+2\text{Ag}(s)$$

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 Cu²⁺ ions dissolve in the solution to yield a characteristic blue color (Figure 7.9).

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 H₂O molecules.
4. Balance hydrogen atoms by adding H⁺ 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 OH⁻ ions to both sides of the equation in numbers equal to the number of H⁺ ions.
   • On the side of the equation containing both H⁺ and OH⁻ 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^[1] are balanced.