Chapter 6. Chemical Reactions and Equations

# 6.4 Oxidation-Reduction Reactions

### Learning Objectives

By the end of this section, you will be able to:

• Define oxidation and reduction.
• Assign oxidation numbers to atoms in simple compounds.
• Recognize a reaction as an oxidation-reduction reaction.
• Recognize composition, decomposition, combustion and single replacement reactions.
• Predict the products of a combustion reaction.

## Redox 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:

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

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:

$2\text{Na}(s) \longrightarrow 2\text{Na}^{+}(s) + 2\text{e}^{-}$
$\text{Cl}_2(g) + 2\text{e}^{-} \longrightarrow 2\text{Cl}^{-}(s)$

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:

$\begin{array}{r @ {{}={}} l} \pmb{\text{oxidation}} & \text{loss of electrons} \\[1em] \pmb{\text{reduction}} & \text{gain of electrons} \end{array}$

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.

$\begin{array}{r @ {{}={}} l} \pmb{\text{reducing agent}} & \text{species that is oxidized} \\[1em] \pmb{\text{oxidizing agent}} & \text{species that is reduced} \end{array}$

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

$\text{H}_2(g) + \text{Cl}_2(g) \longrightarrow 2 \text{HCl}(g)$
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 non-metals 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, O22−), very rarely $-\frac{1}{2}$ (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)
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.

### Example 1

Follow the guidelines in this section of the text to assign oxidation numbers to all the elements in the following species:

a) H2S

b) SO32−

c) Na2SO4

Solution
a) According to guideline 1, the oxidation number for H is +1.

Using this oxidation number and the compound’s formula, guideline 4 may then be used to calculate the oxidation number for sulfur:

$\text{charge on H}_2 \text{S} = 0 = (2 \times +1) + (1 \times x)$
$x = 0 = - (2 \times +1) = -2$

b) Guideline 3 suggests the oxidation number for oxygen is −2.

Using this oxidation number and the ion’s formula, guideline 4 may then be used to calculate the oxidation number for sulfur:

${\text{charge on SO}_3}^{2-} = -2 = (3 \times -2) + (1 \times x)$
$x = -2 - (3 \times -2) = +4$

c) For ionic compounds, it’s convenient to assign oxidation numbers for the cation and anion separately.

According to guideline 2, the oxidation number for sodium is +1.

Assuming the usual oxidation number for oxygen (-2 per guideline 3), the oxidation number for sulfur is calculated as directed by guideline 4:

${\text{charge on SO}_4}^{2-} = -2 = (4 \times -2) + (1 \times x)$
$x = -2 -(4 \times -2) = +6$

Test Yourself
Assign oxidation states to the elements whose atoms are underlined in each of the following compounds or ions:

a) KNO3     b) AlH3    c) NH4+    d) H2PO4

a) N, +5     b) Al, +3     c) N, −3    d) P, +5

### Example 2

Assign oxidation numbers to the atoms in each substance.

a) Br2      b) SiO2      c) Ba(NO3)2

Solution

a) Br2 is the elemental form of bromine. Therefore, by rule 1, each atom has an oxidation number of 0.

b) By rule 3, oxygen is normally assigned an oxidation number of −2. For the sum of the oxidation numbers to equal the charge on the species (which is zero), the silicon atom is assigned an oxidation number of +4.

c) The compound barium nitrate can be separated into two parts: the Ba2+ ion and the nitrate ion. Considering these separately, the Ba2+ ion has an oxidation number of +2 by rule 2. Now consider the NO3 ion. Oxygen is assigned an oxidation number of −2, and there are three oxygens. According to rule 4, the sum of the oxidation number on all atoms must equal the charge on the species, so we have the simple algebraic equation

x + 3(−2) = −1

where x is the oxidation number of the nitrogen atom and −1 represents the charge on the species. Evaluating,

x + (−6) = −1
x = +5

Thus, the oxidation number on the N atom in the nitrate ion is +5.

Test Yourself

Assign oxidation numbers to the atoms in H3PO4.

H = +1, O = −2, P = +5

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 5c). Definitions for the complementary processes of this reaction class are correspondingly revised as shown here:

$\pmb{\text{oxidation}} = \text{increase in oxidation number}$
$\pmb{\text{reduction}} = \text{decrease in oxidation number}$

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).

## Classification of Redox Reactions

Four classifications of chemical reactions will be reviewed in this section. Predicting the products in some of them may be difficult, but the reactions are still easy to recognize.

1 – A composition reaction (sometimes also called a combination reaction or a synthesis reaction) produces a single substance from multiple reactants. A single substance as a product is the key characteristic of the composition reaction. There may be a coefficient other than one for the substance, but if the reaction has only a single substance as a product, it can be called a composition reaction. In the reaction

2 H2(g) + O2(g) $\longrightarrow$ 2 H2O(ℓ)

water is produced from hydrogen and oxygen. Although there are two molecules of water being produced, there is only one substance—water—as a product. So this is a composition reaction.

2 – A decomposition reaction starts from a single substance and produces more than one substance; that is, it decomposes. One substance as a reactant and more than one substance as the products is the key characteristic of a decomposition reaction. For example, in the decomposition of sodium hydrogen carbonate (also known as sodium bicarbonate),

2 NaHCO3(s) $\longrightarrow$ Na2CO3(s) + CO2(g) + H2O(ℓ)

sodium carbonate, carbon dioxide, and water are produced from the single substance sodium hydrogen carbonate.

Composition and decomposition reactions are difficult to predict; however, they should be easy to recognize.

### Example 3

Identify each equation as a composition reaction, a decomposition reaction, or neither.

a) Fe2O3 + 3 SO3 $\longrightarrow$ Fe2(SO4)3

b) NaCl + AgNO3 $\longrightarrow$ AgCl + NaNO3

c) (NH4)2Cr2O7 $\longrightarrow$ Cr2O3 + 4 H2O + N2

Solution

a) In this equation, two substances combine to make a single substance. This is a composition reaction.

b) Two different substances react to make two new substances. This does not fit the definition of either a composition reaction or a decomposition reaction, so it is neither. In fact, you may recognize this as a double-replacement reaction.

c) A single substance reacts to make multiple substances. This is a decomposition reaction.

Test Yourself

Identify the equation as a composition reaction, a decomposition reaction, or neither.

C3H8 $\longrightarrow$ C3H4 + 2 H2

decomposition

3 – Combustion reactions in which the reductant, also called a fuel, and oxidant, molecular oxygen, react vigorously and produce significant amounts of heat, and often light, in the form of a flame.  Combustion reactions produce oxides of all other elements as products; any nitrogen in the reactant is converted to elemental nitrogen, N2. Many reactants, called fuels, contain mostly carbon and hydrogen atoms, reacting with oxygen to produce CO2 and H2O. For example, the balanced chemical equation for the combustion of methane, CH4, is as follows:

CH4 + 2 O2 $\longrightarrow$ CO2 + 2 H2O

Kerosene can be approximated with the formula C12H26, and its combustion equation is

2 C12H26 + 37 O2 $\longrightarrow$ 24 CO2 + 26 H2O

Sometimes fuels contain oxygen atoms, which must be counted when balancing the chemical equation. One common fuel is ethanol, C2H5OH, whose combustion equation is

C2H5OH + 3 O2 $\longrightarrow$ 2 CO2 + 3 H2O

If nitrogen is present in the original fuel, it is converted to N2, not to a nitrogen-oxygen compound. Thus, for the combustion of the fuel dinitroethylene, whose formula is C2H2N2O4, we have

2 C2H2N2O4 + O2 $\longrightarrow$ 4 CO2 + 2 H2O + 2 N2

### Example 4

Complete and balance each combustion equation.

a) the combustion of propane, C3H8

b) the combustion of ammonia, NH3

Solution

a) The products of the reaction are CO2 and H2O, so our unbalanced equation is

C3H8 + O2 $\longrightarrow$ CO2 + H2O

Balancing (and you may have to go back and forth a few times to balance this), we get

C3H8 + 5 O2 $\longrightarrow$ 3 CO2 + 4 H2O

b) The nitrogen atoms in ammonia will react to make N2, while the hydrogen atoms will react with O2 to make H2O:

NH3 + O2 $\longrightarrow$ N2 + H2O

To balance this equation without fractions (which is the convention), we get

4 NH3 + 3 O2 $\longrightarrow$ 2 N2 + 6 H2O

Test Yourself

Complete and balance the combustion equation for cyclopropanol, C3H6O.

C3H6O + 4 O2 $\longrightarrow$ 3 CO2 + 3 H2O

Watch a brief video showing the test firing of a small-scale, prototype, hybrid rocket engine planned for use in the new Space Launch System being developed by NASA. The first engines firing at 3 s (green flame) use a liquid fuel/oxidant mixture, and the second, more powerful engines firing at 4 s (yellow flame) use a solid mixture.

4 – 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:

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

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

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

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 2).

### Example 5

Identify which equations represent redox reactions, providing a name for the reaction if appropriate. For those reactions identified as redox, name the oxidant and reductant.

a) $\text{ZnCO}_3(s) \longrightarrow \text{ZnO}(s) + \text{CO}_2(g)$

b) $2\text{Ga}(l) + 3\text{Br}_2(l) \longrightarrow 2\text{GaBr}_3(s)$

c) $2\text{H}_2 \text{O}_2(aq) \longrightarrow 2\text{H}_2 \text{O}(l) + \text{O}_2(g)$

d) $\text{BaCl}_2(aq) + \text{K}_2 \text{SO}_4(aq) \longrightarrow \text{BaSO}_4(s) + 2\text{KCl}(aq)$

e) $\text{C}_2 \text{H}_4(g) + 3\text{O}_2(g) \longrightarrow 2\text{CO}_2(g) + 2\text{H}_2 \text{O}(l)$

Solution
Redox reactions are identified per definition if one or more elements undergo a change in oxidation number.

a) This is not a redox reaction, since oxidation numbers remain unchanged for all elements.

b) This is a redox reaction. Gallium is oxidized, its oxidation number increasing from 0 in Ga(l) to +3 in GaBr3(s). The reducing agent is Ga(l). Bromine is reduced, its oxidation number decreasing from 0 in Br2(l) to −1 in GaBr3(s). The oxidizing agent is Br2(l).

c) This is a redox reaction. It is a particularly interesting process, as it involves the same element, oxygen, undergoing both oxidation and reduction (a so-called disproportionation reaction). Oxygen is oxidized, its oxidation number increasing from −1 in H2O2(aq) to 0 in O2(g). Oxygen is also reduced, its oxidation number decreasing from −1 in H2O2(aq) to −2 in H2O(l). For disproportionation reactions, the same substance functions as an oxidant and a reductant.

d) This is not a redox reaction, since oxidation numbers remain unchanged for all elements.

e) This is a redox reaction (combustion). Carbon is oxidized, its oxidation number increasing from −2 in C2H4(g) to +4 in CO2(g). The reducing agent (fuel) is C2H4(g). Oxygen is reduced, its oxidation number decreasing from 0 in O2(g) to −2 in H2O(l). The oxidizing agent is O2(g).

Test Yourself
This equation describes the production of tin(II) chloride:

$\text{Sn}(s) + 2\text{HCl}(g) \longrightarrow \text{SnCl}_2(s) + \text{H}_2(g)$

Is this a redox reaction? If so, provide a more specific name for the reaction if appropriate, and identify the oxidant and reductant.

Yes, a single-replacement reaction. Sn(s) is the reductant, HCl(g) is the oxidant.

## Key Concepts and Summary

Chemical reactions are classified according to similar patterns of behavior. Redox reactions involve a change in oxidation number for one or more reactant elements. There are four classifications of chemical reactions: composition, decomposition, combustion and single displacement.

### Exercises

1. Is the reaction

2 K(s) + Br2(ℓ) $\longrightarrow$ 2 KBr(s)

2. In the reaction

2 Ca(s) + O2(g) $\longrightarrow$ 2 CaO

indicate what has lost electrons and what has gained electrons.

3. In the reaction

2 Li(s) + O2(g) $\longrightarrow$ Li2O2(s)

indicate what has been oxidized and what has been reduced.

4. Assign oxidation numbers to each atom in each substance.

a)  P4               b)  SO2

c)  SO22−       d)  Ca(NO3)2

5. .  Assign oxidation numbers to each atom in each substance.

a)  CO          b)  CO2

c)  NiCl2        d)  NiCl3

6.  Assign oxidation numbers to each atom in each substance.

a)  CH2O      b)  NH3

c)  Rb2SO4    d)  Zn(C2H3O2)2

7.  Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.

2 NO + Cl2 $\longrightarrow$ 2 NOCl

8.  Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.

2 KrF2 + 2 H2O $\longrightarrow$ 2 Kr + 4 HF + O2

9.  Identify what is being oxidized and reduced in this redox equation by assigning oxidation numbers to the atoms.

2 K + MgCl2 $\longrightarrow$ 2 KCl + Mg

10. Indicate what type, or types, of reaction each of the following represents:

a) $\text{Ca}(s) + \text{Br}_2(l) \longrightarrow \text{CaBr}_2(s)$

b) $\text{Ca(OH)}_2 (aq) + 2\text{HBr}(aq) \longrightarrow \text{CaBr}_2(aq) + 2\text{H}_2 \text{O}(l)$

c) $\text{C}_6 \text{H}_{12}(l) + 9\text{O}_2(g) \longrightarrow 6\text{CO}_2(g) + 6\text{H}_2 \text{O}(g)$

11. Indicate what type, or types, of reaction each of the following represents:

a) $\text{H}_2 \text{O}(g) + \text{C}(s) \longrightarrow \text{CO}(g) + \text{H}_2(g)$

b) $2\text{KClO}_3(s) \longrightarrow 2\text{KCl}(s) + 3\text{O}_2(g)$

c) $\text{Al(OH)}_3(aq) + 3\text{HCl}(aq) \longrightarrow \text{AlCl}_3(aq) + 3\text{H}_2 \text{O}(l)$

d) $\text{Pb(NO}_3)_2(aq) + \text{H}_2 \text{SO}_4(sq) \longrightarrow \text{PbSO}_4(s) + 2\text{HNO}_3(aq)$

12. Silver can be separated from gold because silver dissolves in nitric acid while gold does not. Is the dissolution of silver in nitric acid an acid-base reaction or an oxidation-reduction reaction? Explain your answer.

13. Determine the oxidation states of the elements in the compounds listed. None of the oxygen-containing compounds are peroxides or superoxides.

a) H3PO4      b) Al(OH)3      c) SeO2

d) KNO2       e) In2S3            f) P4O6

14. Classify the following as acid-base reactions or oxidation-reduction reactions:

a) $\text{Na}_2 \text{S}(aq) + 2 \text{HCl}(aq) \longrightarrow 2 \text{NaCl}(aq) + \text{H}_2 \text{S}(g)$

b) $2\text{Na}(s) + 2\text{HCl}(aq) \longrightarrow 2\text{NaCl}(aq) + \text{H}_2(g)$

c) $\text{Mg}(s) + \text{Cl}_2(g) \longrightarrow \text{MgCl}_2(aq)$

d) $\text{MgO}(s) + 2\text{HCl}(aq) \longrightarrow \text{MgCl}_2(s) + \text{H}_2 \text{O}(l)$

e) $\text{K}_3 \text{P}(s) + 2\text{O}_2(g) \longrightarrow \text{K}_3 \text{PO}_4(s)$

f) $3\text{KOH}(aq) + \text{H}_3 \text{PO}_4(aq) \longrightarrow \text{K}_3\text{PO}_4(aq) + 3 \text{H}_2 \text{O}(l)$

15. Complete and balance the following acid-base equations:

a) HCl gas reacts with solid Ca(OH)2(s).

b) A solution of Sr(OH)2 is added to a solution of HNO3.

16. Complete and balance the following oxidation-reduction reactions, which give the highest possible oxidation state for the oxidized atoms.

a) $\text{Al}(s) + \text{F}_2(g) \longrightarrow$

b) $\text{Al}(s) + \text{CuBr}_2(aq) \longrightarrow \;\text{(single displacement)}$

c) $\text{P}_4(s) + \text{O}_2(g) \longrightarrow$

d) $\text{Ca}(s) + \text{H}_2 \text{O}(l) \longrightarrow \;\text{(products are a strong base and a diatomic gas)}$

17. The military has experimented with lasers that produce very intense light when fluorine combines explosively with hydrogen. What is the balanced equation for this reaction?

18. Great Lakes Chemical Company produces bromine, Br2, from bromide salts such as NaBr, in Arkansas brine by treating the brine with chlorine gas. Write a balanced equation for the reaction of NaBr with Cl2.

19. Lithium hydroxide may be used to absorb carbon dioxide in enclosed environments, such as manned spacecraft and submarines. Write an equation for the reaction that involves 2 mol of LiOH per 1 mol of CO2. (Hint: Water is one of the products.)

20. Complete and balance the equations of the following reactions, each of which could be used to remove hydrogen sulfide from natural gas:

a) $\text{Ca(OH)}_2(s) + \text{H}_2 \text{S}(g) \longrightarrow$

b) $\text{Na}_2 \text{CO}_3(aq) + \text{H}_2 \text{S}(g) \longrightarrow$

21. Write balanced chemical equations for the reactions used to prepare each of the following compounds from the given starting material(s). In some cases, additional reactants may be required.

a) solid ammonium nitrate from gaseous molecular nitrogen via a two-step process (first reduce the nitrogen to ammonia, then neutralize the ammonia with an appropriate acid)

b) gaseous hydrogen bromide from liquid molecular bromine via a one-step redox reaction

c) gaseous H2S from solid Zn and S via a two-step process (first a redox reaction between the starting materials, then reaction of the product with a strong acid)

22. Which is a composition reaction and which is not?

a)  NaCl + AgNO3 $\longrightarrow$ AgCl + NaNO3

b)  CaO + CO2 $\longrightarrow$ CaCO3

23.  Which is a composition reaction and which is not?

a)  2 SO2 + O2 $\longrightarrow$ 2 SO3

b)  6 C + 3 H2 $\longrightarrow$ C6H6

24.  Which is a decomposition reaction and which is not?

a)  HCl + NaOH $\longrightarrow$ NaCl + H2O

b)  CaCO3 $\longrightarrow$ CaO + CO2

25.  Which is a decomposition reaction and which is not?

a)  Na2O + CO2 $\longrightarrow$ Na2CO3

b)  H2SO3 $\longrightarrow$ H2O + SO2

26.  Which is a combustion reaction and which is not?

a)  C6H12O6 + 6 O2 $\longrightarrow$ 6 CO2 + 6 H2O

b)  2 Fe2S3 + 9 O2 $\longrightarrow$ 2 Fe2O3 + 6 SO2

27.  Which is a combustion reaction and which is not?

a)  P4 + 5 O2 $\longrightarrow$ 2 P2O5

b)  2 Al2S3 + 9 O2 $\longrightarrow$ 2 Al2O3 + 6 SO2

28.  Is it possible for a composition reaction to also be a combustion reaction? Give an example to support your case.

29.  Complete and balance each combustion equation.

a)  C4H9OH + O2 $\longrightarrow$ ?

b)  CH3NO2 + O2 $\longrightarrow$ ?

1. Yes; both K and Br are changing oxidation numbers.

2. Ca has lost electrons, and O has gained electrons.

3. Li has been oxidized, and O has been reduced.

4. a)  P: 0

b)  S: +4; O: −2

c)  S: +2; O: −2

d)  Ca: 2+; N: +5; O: −2

5. a)  C: +2; O: −2

b)  C: +4; O: −2

c)  Ni: +2; Cl: −1

d)  Ni: +3; Cl: −1

6. a)  C: 0; H: +1; O: −2

b)  N: −3; H: +1

c)  Rb: +1; S: +6; O: −2

d)  Zn: +2; C: 0; H: +1; O: −2

7. N is being oxidized, and Cl is being reduced.

8. O is being oxidized, and Kr is being reduced.

9. K is being oxidized, and Mg is being reduced.

10. a) oxidation-reduction (addition); b) acid-base (neutralization); c) oxidation-reduction (combustion)

11. a) single replacement;   b) decomposition;   c) acid-base;   d) precipitation

12. It is an oxidation-reduction reaction because the oxidation state of the silver changes during the reaction.

13. a) H +1, P +5, O −2;    b) Al +3, H +1, O −2;    c) Se +4, O −2;

d) K +1, N +3, O −2;    e) In +3, S −2;    f) P +3, O −2

14. a) acid-base;    b) oxidation-reduction: Na is oxidized, H+ is reduced;

c) oxidation-reduction: Mg is oxidized, Cl2 is reduced;     d) acid-base;

e) oxidation-reduction: P3− is oxidized, O2 is reduced;     f) acid-base

15. a) $2\text{HCl}(g) + \text{Ca(OH)}_2(s) \longrightarrow \text{CaCl}_2(s) + 2\text{H}_2 \text{O}(l)$;
b) $\text{Sr(OH)}_2(aq) + 2\text{HNO}_3(aq) \longrightarrow \text{Sr(NO}_3)_2(aq) + 2\text{H}_2 \text{O}(l)$;

16. a) $2\text{Al}(s) + 3\text{F}_2 \longrightarrow 2\text{AlF}_3(s)$;
b) $2\text{Al}(s) + 3\text{CuBr}_2(aq) \longrightarrow 3\text{Cu}(s) + 2\text{AlBr}_3(aq)$;
c) $\text{P}_4(s) + 5\text{O}_2(g) \longrightarrow \text{P}_4 \text{O}_{10}(s)$;
d) $\text{Ca}(s) + 2\text{H}_2 \text{O}(l) \longrightarrow \text{Ca(OH)}_2(aq) + \text{H}_2(g)$;

17. $\text{H}_2(g) + \text{F}_2(g) \longrightarrow 2\text{HF}(g)$

18. $2\text{NaBr}(aq) + \text{Cl}_2(g) \longrightarrow 2\text{NaCl}(aq) + \text{Br}_2(l)$

19. $2\text{LiOH}(aq) + \text{CO}_2(g) \longrightarrow \text{Li}_2 \text{CO}_3(aq) + \text{H}_2 \text{O}(l)$

20. a) $\text{Ca(OH)}_2(s) + \text{H}_2 \text{S}(g) \longrightarrow \text{CaS}(s) + 2\text{H}_2\text{O}(l);$
b) $\text{Na}_2 \text{CO}_3(aq) + \text{H}_2 \text{S}(g) \longrightarrow \text{Na}_2 \text{S}(aq) + \text{CO}_2(g) + \text{H}_2 \text{O}(l)$

21. a) step 1: $\text{N}_2(g) + 3\text{H}_2(g) \longrightarrow 2\text{NH}_3(g)$,

step 2: $\text{NH}_3(g) + \text{HNO}_3(aq) \longrightarrow \text{NH}_4 \text{NO}_3(aq) \longrightarrow \text{NH}_4 \text{NO}_3\;\text{(s) (after drying)}$

b) $\text{H}_2(g) + \text{Br}_2(l) \longrightarrow 2\text{HBr}(g)$
c) $\text{Zn}(s) + \text{S}(s) \longrightarrow \text{ZnS}(s) \;\text{and} \;\text{ZnS}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2 \text{S}(g)$

22. a)  not composition    b)  composition

23. a)  composition    b)  composition

24. a)  not decomposition    b)  decomposition

25. a)  not decomposition    b)  decomposition

26. a)  combustion   b)  combustion

27. a)  combustion    b)  combustion

28. Yes; 2 H2 + O2 $\longrightarrow$ 2 H2O (answers will vary)

29. a)  C4H9OH + 6 O2 $\longrightarrow$ 4 CO2 + 5 H2O

b)  4 CH3NO2 + 3 O2 $\longrightarrow$ 4 CO2 + 6 H2O + 2 N2

## Glossary

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.

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

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

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