Oxidation-Reduction Reactions

8.3 Oxidation-Reduction Reactions

Learning Objectives

To identify fundamental types of chemical reactions.
To predict the types of reactions substances will undergo.

In the previous section we learned how to use six rules to assign an oxidation state to an element, given a formula or symbol. In this section we are going to look more at the chemical changes associated with oxidation-reduction reactions.

Oxidation–Reduction Reactions

The term oxidationThe loss of one or more electrons in a chemical reaction. The substance that loses electrons is said to be oxidized. was first used to describe reactions in which metals react with oxygen in air to produce metal oxides. When iron is exposed to air in the presence of water, for example, the iron turns to rust—an iron oxide. When exposed to air, aluminum metal develops a continuous, coherent, transparent layer of aluminum oxide on its surface. In both cases, the metal acquires a positive charge by transferring electrons to the neutral oxygen atoms of an oxygen molecule. As a result, the oxygen atoms acquire a negative charge and form oxide ions (O²⁻). Because the metals have lost electrons to oxygen, they have been oxidized; oxidation is therefore the loss of electrons. Conversely, because the oxygen atoms have gained electrons, they have been reduced, so reduction is the gain of electrons. For every oxidation, there must be an associated reduction.

Note the Pattern

Any oxidation must be accompanied by a reduction and vice versa.

Originally, the term reductionThe gain of one or more electrons in a chemical reaction. The substance that gains electrons is said to be reduced. referred to the decrease in mass observed when a metal oxide was heated with carbon monoxide, a reaction that was widely used to extract metals from their ores. As you have seen or soon will see in chemistry lab, when solid copper(II) oxide is heated with methane, CH₄, its mass decreases because the formation of pure copper is accompanied by the loss of oxygen atoms as a volatile product (water). The reaction is as follows:

Equation 8.3(eq1)

4CuO(s) + CH₄(g) → 4Cu(s) + 2H₂O(g)+ CO₂(g)

Oxidation and reduction reactions are now defined as reactions that exhibit a change in the oxidation states of one or more elements in the reactants, which follows the mnemonic oxidation is loss reduction is gain, or oil rig. The oxidation stateThe charge that each atom in a compound would have if all its bonding electrons were transferred to the atom with the greater attraction for electrons. of each atom in a compound is the charge an atom would have if all its bonding electrons were transferred to the atom with the greater attraction for electrons. Atoms in their elemental form, such as O₂ or H₂, are assigned an oxidation state of zero. For example, the reaction of aluminum with oxygen to produce aluminum oxide is

Equation 8.3(eq2)

4Al(s) + 3O₂(g) → 2Al₂O₃(s)

Each neutral oxygen atom gains two electrons and becomes negatively charged, forming an oxide ion; thus, oxygen has an oxidation state of −2 in the product and has been reduced. Each neutral aluminum atom loses three electrons to produce an aluminum ion with an oxidation state of +3 in the product, so aluminum has been oxidized. In the formation of Al₂O₃, electrons are transferred as follows (the superscript 0 emphasizes the oxidation state of the elements):

Equation 8.3(eq3)

4Al⁰ + 3O₂⁰ → 4Al³⁺ + 6O²⁻

Equation 8.3(eq1) and Equation 8.3(eq2) are examples of oxidation–reduction (redox) reactions. In redox reactions, there is a net transfer of electrons from one reactant to another. In any redox reaction, the total number of electrons lost must equal the total of electrons gained to preserve electrical neutrality. In Equation 8.3(eq3), for example, the total number of electrons lost by aluminum is equal to the total number gained by oxygen:

Equation 8.3(eq4)

\begin{array}{l} electrons lost = 4 Al atoms \times \frac{{3 e}^{-} lost}{Al atom} = {12 e}^{-} lost \\ electrons gained = 6 O atoms \times \frac{{2 e}^{-} gained}{O atom} = {12 e}^{-} gained \end{array}

The same pattern is seen in all oxidation–reduction reactions: the number of electrons lost must equal the number of electrons gained.

Note the Pattern

In all oxidation–reduction (redox) reactions, the number of electrons lost equals the number of electrons gained.

Oxidants and Reductants

Compounds that are capable of accepting electrons, such as O₂ or F₂, are called oxidants (or oxidizing agents)A compound that is capable of accepting electrons; thus it is reduced. because they can oxidize other compounds. In the process of accepting electrons, an oxidant is reduced. Compounds that are capable of donating electrons, such as sodium metal or cyclohexane (C₆H₁₂), are called reductants (or reducing agents)A compound that is capable of donating electrons; thus it is oxidized. because they can cause the reduction of another compound. In the process of donating electrons, a reductant is oxidized. These relationships are summarized in Equation 8.3(eq5):

Equation 8.3(eq5)

oxidant + reductant → oxidation−reduction

\underset{\begin{matrix} {gains e}^{–} \\ (is reduced) \end{matrix}}{O_{2} (g)} + \underset{\begin{matrix} {loses e}^{–} \\ (is oxidized) \end{matrix}}{4Na(s)} \to \underset{redox reaction}{{2Na}_{2} O(s)}

Some oxidants have a greater ability than others to remove electrons from other compounds. Oxidants can range from very powerful, capable of oxidizing most compounds with which they come in contact, to rather weak. Both F₂ and Cl₂ are powerful oxidants: for example, F₂ will oxidize H₂O in a vigorous, potentially explosive reaction. In contrast, S₈ is a rather weak oxidant, and O₂ falls somewhere in between. Conversely, reductants vary in their tendency to donate electrons to other compounds. Reductants can also range from very powerful, capable of giving up electrons to almost anything, to weak. The alkali metals are powerful reductants, so they must be kept away from atmospheric oxygen to avoid a potentially hazardous redox reaction.

A combustion reactionAn oxidation–reduction reaction in which the oxidant is $O_{2}$ . is an oxidation–reduction reaction in which the oxidant is O₂. Consider, for example, the combustion of cyclohexane, a typical hydrocarbon, in excess oxygen. The balanced chemical equation for the reaction, with the oxidation state shown for each atom, is as follows:

Equation 8.3(eq6)

\underset{–2}{C_{6}} \overset{+1}{H_{12}} + \overset{0}{{9O}_{2}} \to 6 \overset{+4}{C} \underset{–2}{O_{2}} + 6 \overset{+1}{H_{2}} \underset{–2}{O}

If we compare the oxidation state of each element in the products and the reactants, we see that hydrogen is the only element whose oxidation state does not change; it remains +1. Carbon, however, has an oxidation state of −2 in cyclohexane and +4 in CO₂; that is, each carbon atom changes its oxidation state by six electrons during the reaction. Oxygen has an oxidation state of 0 in the reactants, but it gains electrons to have an oxidation state of −2 in CO₂ and H₂O. Because carbon has been oxidized, cyclohexane is the reductant; because oxygen has been reduced, it is the oxidant. All combustion reactions are therefore oxidation–reduction reactions.

Other Examples of Redox Reactions

The reaction of bromine with ethylene to give 1,2-dibromoethane, which is used in agriculture to kill nematodes in soil, is as follows:

Equation 8.3(eq7)

C₂H₄(g) + Br₂(g) → BrCH₂CH₂Br(g)

Some reaction classification schemes call this a condensation reaction because it has the general form A + B → AB. This reaction, however, can also be viewed as an oxidation–reduction reaction, in which electrons are transferred from carbon (−2 → −1) to bromine (0 → −1). Another example of a condensation reaction is the one used for the industrial synthesis of ammonia:

Equation 8.3(eq8)

3H₂(g) + N₂(g) → 2NH₃(g)

Although this reaction also has the general form of a condensation reaction, hydrogen has been oxidized (0 → +1) and nitrogen has been reduced (0 → −3), so it can also be classified as an oxidation–reduction reaction.

Example 8.3-1

The following reactions have important industrial applications. For each redox reaction, identify the oxidant and reductant and specify which reactant is oxidized or reduced.

C₂H₄(g) + Cl₂(g) → ClCH₂CH₂Cl(g)
Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

Given: balanced chemical equation

Asked for: oxidant, reductant, oxidized, reduced

Strategy:

Assign oxidation states to each atom present in the reactants and the products. If the oxidation state of one or more atoms changes, then the reaction is a redox reaction.

Solution:

A This reaction is used to prepare 1,2-dichloroethane, one of the top 25 industrial chemicals in the United States. We need to look at the oxidation states of the atoms:
$\underset{–2}{C_{2}} \overset{+1}{H_{4}} + \overset{0}{{Cl}_{2}} \to \underset{–1}{Cl} \underset{–1}{C} \overset{+1}{H_{2}} \underset{–1}{C} \overset{+1}{H_{2}} \underset{–1}{Cl}$
The oxidation states show that chlorine is reduced from 0 to −1 and carbon is oxidized from −2 to −1. Ethylene is the reductant, and chlorine is the oxidant. Ethylene is oxidized and chlorine is reduced.
This reaction occurs in a conventional car battery every time the engine is started. An acid (H₂SO₄) is present and transfers protons to oxygen in PbO₂ to form water during the reaction. The reaction can therefore be described as an acid–base reaction.

B The oxidation states are as follows:
$\overset{0}{Pb} + \overset{+4}{Pb} \underset{–2}{O_{2}} + 2 \overset{+1}{H_{2}} \overset{+6}{S} \underset{–2}{O_{4}} \to 2 \overset{+2}{Pb} \overset{+6}{S} \underset{–2}{O_{4}} + 2 \overset{+1}{H_{2}} \underset{–2}{O}$
The oxidation state of lead changes from 0 in Pb and +4 in PbO₂ (both reactants) to +2 in PbSO₄. This is also a redox reaction, in which elemental lead is the reductant, and PbO₂ is the oxidant. Elemental lead is oxidized, lead(IV) oxide is reduced.

Schematic drawing of a 12-volt car battery. The locations of the reactants (lead metal in a spongy form with large surface area) and PbO₂ are shown. The product (PbSO₄) forms as a white solid between the plates.

Exercise

For each redox reaction, identify the oxidant and the reductant and specify which reactants are oxidized or reduced.

Al(s) + OH⁻(aq) + 3H₂O(l) → 3/2H₂(g) + [Al(OH)₄]⁻(aq)
TiCl₄(l) + 2Mg(l) → Ti(s) + 2MgCl₂(l)
CO(g) + Cl₂(g) → Cl₂CO(l)

Answer:

Redox reaction; reductant is Al, oxidant is H₂O; Al is oxidized, water is reduced. This is the reaction that occurs when Drano is used to clear a clogged drain.
Redox reaction; reductant is Mg, oxidant is TiCl₄; Mg is oxidized, titanium(IV) chloride is reduced.
Both a condensation reaction and a redox reaction; reductant is CO, oxidant is Cl₂; carbon monoxide is oxidized, chlorine is reduced. The product of this reaction is phosgene, a highly toxic gas used as a chemical weapon in World War I. Phosgene is now used to prepare polyurethanes, which are used in foams for bedding and furniture and in a variety of coatings.

Key Takeaways

Oxidation-reduction (redox) reactions involve the transfer of electrons from one atom to another.
Oxidation is an increase of oxidation number (a loss of electrons); reduction is a decrease in oxidation number (a gain of electrons).

Summary

To keep track of electrons in chemical reactions, oxidation states are assigned to atoms in compounds. In an oxidation–reduction reaction, one atom must lose electrons and another must gain electrons. Oxidation is the loss of electrons, and an element whose oxidation state increases is said to be oxidized. Reduction is the gain of electrons, and an element whose oxidation state decreases is said to be reduced. Oxidants are compounds that are capable of accepting electrons from other compounds, so they are reduced during an oxidation–reduction reaction. In contrast, reductants are compounds that are capable of donating electrons to other compounds, so they are oxidized during an oxidation–reduction reaction. A combustion reaction is a redox reaction in which the oxidant is O₂(g).

Conceptual Problems

What is a combustion reaction? How can it be distinguished from an exchange reaction?
What two products are formed in the combustion of an organic compound containing only carbon, hydrogen, and oxygen? Is it possible to form only these two products from a reaction that is not a combustion reaction? Explain your answer.
What factors determine whether a reaction can be classified as a redox reaction?
Name three characteristics of a balanced redox reaction.
Does an oxidant accept electrons or donate them?
Does the oxidation state of a reductant become more positive or more negative during a redox reaction?
Nitrogen, hydrogen, and ammonia are known to have existed on primordial earth, yet mixtures of nitrogen and hydrogen do not usually react to give ammonia. What natural phenomenon would have enough energy to initiate a reaction between these two primordial gases?

Numerical Problems

Is the reaction
2K(s) + Br₂(l) → 2KBr(s)
an oxidation-reduction reaction? Explain your answer.
Is the reaction
NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl(s)
an oxidation-reduction reaction? Explain your answer.
In the reaction
2Ca(s) + O₂(g) → 2CaO
indicate what has lost electrons and what has gained electrons.
In the reaction
16Fe(s) + 3S₈(s) → 8Fe₂S₃(s)
indicate what has lost electrons and what has gained electrons.
In the reaction
2Li(s) + O₂(g) → Li₂O₂(s)
indicate what has been oxidized and what has been reduced.
In the reaction
2Ni(s) + 3I₂(s) → 2NiI₃(s)
indicate what has been oxidized and what has been reduced.
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
2NO + Cl₂ → 2NOCl
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
Fe + SO₃ → FeSO₃
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
2KrF₂ + 2H₂O → 2Kr + 4HF + O₂
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
SO₃ + SCl₂ → SOCl₂ + SO₂
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
2K + MgCl₂ → 2KCl + Mg
Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.
C₇H₁₆ + 11O₂ → 7CO₂ + 8H₂O
For each redox reaction, determine the identities of the oxidant, the reductant, the species oxidized, and the species reduced.
1. H₂(g) + I₂(s) → 2HI(g)
2. 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)
3. 2F₂(g) + 2NaOH(aq) → OF₂(g) + 2NaF(aq) + H₂O(l)
For each redox reaction, determine the identities of the oxidant, the reductant, the species oxidized, and the species reduced.
1. 2Na(s) + Cl₂(g) → 2NaCl(s)
2. SiCl₄(l) + 2Mg(s) → 2MgCl₂(s) + Si(s)
3. 2H₂O₂(aq) → 2H₂O(l) + O₂(g)
Balance each chemical equation. Then identify the oxidant, the reductant, the species oxidized, and the species reduced. (Δ indicates that the reaction requires heating.)
1. H₂O(g) + CO(g) → CO₂(g) + H₂(g)
2. the reaction of aluminum oxide, carbon, and chlorine gas at 900°C to produce aluminum chloride and carbon monoxide
3. $HgO(s) \overset{Δ}{\to} Hg(l) + O_{2} (g)$
  
  Video: Red, mercury(II) oxide is heated. Mercury metal is deposited on the side of the heated tube. The oxygen evolved is collected over water. Evidence of the presence of oxygen is demonstrated by the reignition of a glowing wood splint. Video Credit: Part of NCSSM CORE collection: http://www.dlt.ncssm. Please attribute this work as being created by the North Carolina School of Science and Mathematics. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via YouTube
Balance each chemical equation. Then identify the oxidant, the reductant, the species oxidized, and the species reduced. (Δ indicates that the reaction requires heating.)
1. the reaction of water and carbon at 800°C to produce hydrogen and carbon monoxide
2. Mn(s) + S₈(s) + CaO(s) → CaS(s) + MnO(s)
3. the reaction of ethylene and oxygen at elevated temperature in the presence of a silver catalyst to produce ethylene oxide
4. ZnS(s) + H₂SO₄(aq) + O₂(g) → ZnSO₄(aq) + S₈(s) + H₂O(l)
Silver is tarnished by hydrogen sulfide, an atmospheric contaminant, to form a thin layer of dark silver sulfide (Ag₂S) along with hydrogen gas.
1. Write a balanced chemical equation for this reaction.
2. Which species has been oxidized and which has been reduced?
The following reaction is used in the paper and pulp industry:
Na₂SO₄(aq) + C(s) + NaOH(aq) → Na₂CO₃(aq) + Na₂S(aq) + H₂O(l)
1. Balance the chemical equation.
2. Identify the oxidant and the reductant.

Answers

Yes; both K and Br are changing oxidation numbers.
Ca has lost electrons, and O has gained electrons.
Li has been oxidized, and O has been reduced.
N is being oxidized, and Cl is being reduced.
O is being oxidized, and Kr is being reduced.
K is being oxidized, and Mg is being reduced.
1. Na is the reductant and is oxidized. Cl₂ is the oxidant and is reduced.
2. Mg is the reductant and is oxidized. Si is the oxidant and is reduced.
3. H₂O₂ is both the oxidant and reductant. One molecule is oxidized, and one molecule is reduced.
1. H₂O(g) + C(s) $\overset{Δ}{\to}$ H₂(g) + CO(g)
  
  C is the reductant and is oxidized. H₂O is the oxidant and is reduced.
2. 8Mn(s) + S₈(s) + 8CaO(s) → 8CaS(s) + 8MnO(s)
  
  Mn is the reductant and is oxidized. The S₈ is the oxidant and is reduced.
3. 2C₂H₄(g) + O₂(g) $\overset{Δ}{\to}$ 2C₂H₄O(g)
  
  Ethylene is the reductant and is oxidized. O₂ is the oxidant and is reduced.
4. 8ZnS(s) + 8H₂SO₄(aq) + 4O₂(g) → 8ZnSO₄(aq) + S₈(s) + 8H₂O(l)
  
  Sulfide in ZnS is the reductant and is oxidized. O₂ is the oxidant and is reduced.
1. Na₂SO₄ + 2C + 4NaOH → 2Na₂CO₃ + Na₂S + 2H₂O
2. The sulfate ion is the oxidant, and the reductant is carbon.