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5 Acids, Bases, and Neutralization Reactions

5 Acids, Bases, and Neutralization Reactions

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142



CHAPTER 5



Classification and Balancing of Chemical Reactions



Another kind of neutralization reaction occurs between an acid and a carbonate

(or bicarbonate) to yield water, a salt, and carbon dioxide. Hydrochloric acid reacts

with potassium carbonate, for example, to give H2O, KCl, and CO 2:

2 HCl1aq2 + K2CO 31aq2 h H2O1l2 + 2 KCl1aq2 + CO 21g2



The reaction occurs because the carbonate ion 1CO 32-2 reacts initially with H+ to

yield H2CO 3, which is unstable and immediately decomposes to give CO 2 plus H2O.

We will defer a more complete discussion of carbonates as bases until Chapter 10,

but note for now that they yield OH- ions when dissolved in water just as KOH and

other bases do.

K2CO31s2 + H2O1l2



h 2K+1aq2 + HCO3-1aq2 + OH-1aq2



Dissolve in water



LOOKING AHEAD



Acids and bases are enormously important in biological

chemistry. We will see in Chapter 18, for instance, how acids and bases affect the structure

and properties of proteins.



Worked Example 5.7 Chemical Reactions: Acid–Base Neutralization

Write an equation for the neutralization reaction of aqueous HBr and aqueous

Ba1OH22.

SOLUTION



The reaction of HBr with Ba1OH22 involves the combination of a proton 1H+2 from

the acid with OH- from the base to yield water and a salt 1BaBr22.

2 HBr1aq2 + Ba1OH221aq2 h 2 H2O1l2 + BaBr21aq2



PROBLEM 5.11



Write and balance equations for the following acid–base neutralization reactions:

(a) CsOH1aq2 + H2SO 41aq2 h

(b) Ca1OH221aq2 + CH3CO 2H1aq2 h

(c) NaHCO 31aq2 + HBr1aq2 h



5.6 Redox Reactions

Oxidation–reduction (redox) reactions, the third and final category of reactions that

we will discuss here, are more complex than precipitation and neutralization reactions. Look, for instance, at the following examples and see if you can tell what they

have in common. Copper metal reacts with aqueous silver nitrate to form silver metal

and aqueous copper(II) nitrate; iron rusts in air to form iron(III) oxide; the zinc metal

container on the outside of a battery reacts with manganese dioxide and ammonium

chloride inside the battery to generate electricity and give aqueous zinc chloride plus

manganese(III) oxide. Although these and many thousands of other reactions appear

unrelated, all are examples of redox reactions.

Cu1s2 + 2 AgNO 31aq2 h 2 Ag1s2 + Cu1NO 3221aq2

2 Fe1s2 + 3 O 21g2 h Fe 2O 31s2

Zn1s2 + 2 MnO 21s2 + 2 NH4Cl1s2 h

ZnCl 21aq2 + Mn 2O 31s2 + 2 NH31aq2 + H2O1l2



Oxidation The loss of one or more

electrons by an atom.



Historically, the word oxidation referred to the combination of an element with

oxygen to yield an oxide, and the word reduction referred to the removal of oxygen from

an oxide to yield the element. Today, though, the words have taken on a much broader

meaning. An oxidation is now defined as the loss of one or more electrons by an atom,



SECTION 5.6



and a reduction is the gain of one or more electrons. Thus, an oxidation–reduction

reaction, or redox reaction, is one in which electrons are transferred from one atom to

another.



Redox Reactions



143



Reduction The gain of one or more

electrons by an atom.



Oxidation

A2−



A− + electron



A−



A



+ electron



A



A+



+ electron



A+



A2+



+ electron



Reactant A might be anything:

a neutral atom,

a monatomic ion,

a polyatomic ion,

or a molecule.



Reduction



Take the reaction of copper with aqueous Ag + as an example, as shown in Figure 5.1.

Copper metal gives an electron to each of 2 Ag + ions, forming Cu2+ and silver metal.

Copper is oxidized in the process, and Ag + is reduced. You can follow the transfer of the

electrons by noting that the charge on the copper increases from 0 to + 2 when it loses

2 electrons, whereas the charge on Ag + decreases from + 1 to 0 when it gains an electron.

+2 electrons = reduced!



Cu(s) + 2 Ag+(aq)



Cu2+(aq) + 2 Ag(s)



+1 charge +2 charge



0 charge



0 charge



−2 electrons = oxidized!



◀ Figure



5.1

The copper wire reacts with aqueous

Ag + ion and becomes coated with

metallic silver. At the same time,

copper(II) ions go into solution, producing the blue color.



Similarly, in the reaction of aqueous iodide ion with bromine, iodide ion gives an electron to bromine, forming iodine and bromide ion. Iodide ion is oxidized as its charge

increases from - 1 to 0, and bromine is reduced as its charge decreases from 0 to - 1.

+2 electrons = reduced!



2 I−(aq) + Br2(aq)

−1 charge



0 charge



I2(aq) + 2 Br−(aq)

0 charge



−1 charge



−2 electrons = oxidized!



As these examples show, oxidation and reduction always occur together. Whenever one substance loses an electron (is oxidized), another substance must gain

that electron (be reduced). The substance that gives up an electron and causes the

reduction—the copper atom in the reaction of Cu with Ag + and the iodide ion in

the reaction of I- with Br2 —is called a reducing agent. The substance that gains

an electron and causes the oxidation—the silver ion in the reaction of Cu with Ag +

and the bromine molecule in the reaction of I- with Br2 —is called an oxidizing agent.

The charge on the reducing agent increases during the reaction, and the charge on 

the oxidizing agent decreases.



Reducing agent A reactant that causes

a reduction in another reactant by

giving up electron to it.

Oxidizing agent A reactant that

causes an oxidation by taking electrons

from another reactant.



144



CHAPTER 5



Classification and Balancing of Chemical Reactions



Reducing agent



Loses one or more electrons

Causes reduction

Undergoes oxidation

Becomes more positive (less negative)

(May gain oxygen atoms)



Oxidizing agent



Gains one or more electrons

Causes oxidation

Undergoes reduction

Becomes more negative (less positive)

(May lose oxygen atoms)



Among the simplest of redox processes is the reaction of an element, usually a

metal, with an aqueous cation to yield a different element and a different ion. Iron

metal reacts with aqueous copper(II) ion, for example, to give iron(II) ion and copper

metal. Similarly, magnesium metal reacts with aqueous acid to yield magnesium ion

and hydrogen gas. In both cases, the reactant element (Fe or Mg) is oxidized, and the

reactant ion (Cu2+ or H+ ) is reduced.

Fe1s2 + Cu2+1aq2 h Fe2+1aq2 + Cu1s2

Mg1s2 + 2 H+1aq2 h Mg2+1aq2 + H21g2



The relationship between formation of ions and ionization energy/

electronegativity was discussed in

Chapter 3.



The reaction of a metal with water or aqueous acid 1H+2 to release H2 gas is a particularly important process. As you might expect based on the periodic properties discussed in Section 3.2, the alkali metals and alkaline earth metals (on the left side of the

periodic table) are the most powerful reducing agents (electron donors), so powerful

that they even react with pure water, in which the concentration of H+ is very low. This

is due in part to the fact that alkali metals and alkaline earth metals have low ionization energies. Ionization energy, which is a measure of how easily an element will lose

an electron, tends to decrease as we move to the left and down in the periodic table.

Thus, metals toward the middle of the periodic table, such as iron and chromium, have

higher ionization energies and do not lose electrons as readily; they react only with

aqueous acids but not with water. Those metals near the bottom right of the periodic

table, such as platinum and gold, react with neither aqueous acid nor water. At the

other extreme from the alkali metals, the reactive nonmetals at the top right of the

periodic table have the highest ionization energies and are extremely weak reducing

agents but powerful oxidizing agents (electron acceptors). This is, again, predictable

based on the periodic property of electron affinity (Section 3.2), which becomes more

energetically favored as we move up and to the right in the periodic table.

We can make a few generalizations about the redox behavior of metals and

nonmetals.

1. In reactions involving metals and nonmetals, metals tend to lose electrons

while nonmetals tend to gain electrons. The number of electrons lost or gained

can often be predicted based on the position of the element in the periodic

table. (Section 3.5)

2. In reactions involving nonmetals, the “more metallic” element (farther down

and/or to the left in the periodic table) tends to lose electrons, and the “less

metallic” element (up and/or to the right) tends to gain electrons.

Redox reactions involve almost every element in the periodic table, and they

occur in a vast number of processes throughout nature, biology, and industry.

Here are just a few examples:





Corrosion is the deterioration of a metal by oxidation, such as the rusting

of iron in moist air. The economic consequences of rusting are enormous:

it has been estimated that up to one-fourth of the iron produced in the

United States is used to replace bridges, buildings, and other structures

that have been destroyed by corrosion. (The raised dot in the formula

Fe 2O 3 # H2O for rust indicates that one water molecule is associated with

each Fe 2O 3 in an undefined way.)



SECTION 5.6







Combustion is the burning of a fuel by rapid oxidation with oxygen in air.

Gasoline, fuel oil, natural gas, wood, paper, and other organic substances of

carbon and hydrogen are the most common fuels that burn in air. Even some

metals, though, will burn in air. Magnesium and calcium are examples.

CH41g2 + 2 O 21g2 h CO 21g2 + 2 H2O1l2



Methane

1natural gas2



2 Mg1s2 + O 21g2 h 2 MgO1s2







Respiration is the process of breathing and using oxygen for the many biological redox reactions that provide the energy required by living organisms. We will see in Chapters 21–22 that in the respiration process, energy

is released from food molecules slowly and in complex, multistep pathways,

but that the overall result is similar to that of the simpler combustion reactions. For example, the simple sugar glucose 1C 6H12O 62 reacts with O 2 to

give CO 2 and H2O according to the following equation:

C 6H12O 6 + 6 O 2 h 6 CO 2 + 6 H2O + Energy

Glucose

1a carbohydrate2











Bleaching makes use of redox reactions to decolorize or lighten colored materials. Dark hair is bleached to turn it blond, clothes are bleached to remove

stains, wood pulp is bleached to make white paper, and so on. The oxidizing

agent used depends on the situation: hydrogen peroxide 1H2O 22 is used for

hair, sodium hypochlorite (NaOCl) for clothes, and elemental chlorine for

wood pulp, but the principle is always the same. In all cases, colored organic

materials are destroyed by reaction with strong oxidizing agents.

Metallurgy, the science of extracting and purifying metals from their

ores, makes use of numerous redox processes. Worldwide, approximately

800 million tons of iron are produced each year by reduction of the mineral hematite, Fe 2O 3, with carbon monoxide.

Fe 2O 31s2 + 3 CO1g2 h 2 Fe1s2 + 3 CO 21g2



Worked Example 5.8 Chemical Reactions: Redox Reactions

For the following reactions, indicate which atom is oxidized and which is reduced,

based on the definitions provided in this section. Identify the oxidizing and reducing agents.

(a) Cu1s2 + Pt2+1aq2 h Cu2+1aq2 + Pt1s2

(b) 2 Mg1s2 + CO 21g2 h 2 MgO1s2 + C1s2

ANALYSIS The definitions for oxidation include a loss of electrons, an increase in



charge, and a gain of oxygen atoms; reduction is defined as a gain of electrons, a

decrease in charge, and a loss of oxygen atoms.

SOLUTION



(a) In this reaction, the charge on the Cu atom increases from 0 to 2 + . This

corresponds to a loss of 2 electrons. The Cu is therefore oxidized and acts as the

reducing agent. Conversely, the Pt2+ ion undergoes a decrease in charge from 2 +

to 0, corresponding to a gain of 2 electrons for the Pt2+ ion. The Pt2+ is reduced,

and acts as the oxidizing agent.

(b) In this case, the gain or loss of oxygen atoms is the easiest way to identify which atoms

are oxidized and reduced. The Mg atom is gaining oxygen to form MgO; therefore, the

Mg is being oxidized and acts as the reducing agent. The C atom in CO 2 is losing oxygen. Therefore, the C atom in CO2 is being reduced, and so CO2 acts as the oxidizing

agent.



Redox Reactions



145



146



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Classification and Balancing of Chemical Reactions



Worked Example 5.9 Chemical Reactions: Identifying Oxidizing/Reducing Agents

For the respiration and metallurgy examples discussed previously, identify the

atoms being oxidized and reduced, and label the oxidizing and reducing agents.

ANALYSIS Again, using the definitions of oxidation and reduction provided in this

section, we can determine which atom(s) are gaining/losing electrons or gaining/

losing oxygen atoms.

SOLUTION



Respiration:



C 6H12O 6 + 6 O 2 h 6 CO 2 + 6 H2O



Because the charge associated with the individual atoms is not evident, we will use

the definition of oxidation/reduction as the gaining/losing of oxygen atoms. In this

reaction, there is only one reactant besides oxygen 1C6H12O62, so we must determine which atom in the compound is changing. The ratio of carbon to oxygen in

C6H12O4 is 1:1, while the ratio in CO2 is 1:2. Therefore, the C atoms are gaining

oxygen and are oxidized; the C6H12O16 is the reducing agent and O2 is the oxidizing

agent. Note that the ratio of hydrogen to oxygen in C6H12O6 and in H2O is 2:1.

The H atoms are neither oxidized nor reduced.

Fe 2O 31s2 + 3 CO1g2 h 2 Fe1s2 + 3 CO 21g2

Metallurgy:

The Fe2O3 is losing oxygen to form Fe(s); it is being reduced and acts as the

oxidizing agent. In contrast, the CO is gaining oxygen to form CO2; it is being

oxidized and acts as the reducing agent.



Worked Example 5.10 Chemical Reactions: Identifying Redox Reactions

For the following reactions, identify the atom(s) being oxidized and reduced:

(b) C1s2 + 2 Cl21g2 h CCl41l2

(a) 2 Al1s2 + 3 Cl21g2 h 2 AlCl31s2

ANALYSIS Again, there is no obvious increase or decrease in charge to indicate a



gain or loss of electrons. Also, the reactions do not involve a gain or loss of oxygen.

We can, however, evaluate the reactions in terms of the typical behavior of metals

and nonmetals in reactions.

SOLUTION



(a) In this case, we have the reaction of a metal (Al) with a nonmetal 1Cl22. Because

metals tend to lose electrons and nonmetals tend to gain electrons, we can assume

that the Al atom is oxidized (loses electrons) and the Cl2 is reduced (gains electrons).

(b) The carbon atom is the less electronegative element (farther to the left) and is

less likely to gain an electron. The more electronegative element (Cl) will tend to

gain electrons (be reduced).



PROBLEM 5.12



Identify the oxidized reactant, the reduced reactant, the oxidizing agent, and the

reducing agent in the following reactions:

(a) Fe1s2 + Cu2+1aq2 h Fe2+1aq2 + Cu1s2

(b) Mg1s2 + Cl21g2 h MgCl21s2

(c) 2 Al1s2 + Cr2O31s2 h 2 Cr1s2 + Al2O31s2

PROBLEM 5.13



Potassium, a silvery metal, reacts with bromine, a corrosive, reddish liquid, to yield

potassium bromide, a white solid. Write the balanced equation, and identify the

oxidizing and reducing agents.

PROBLEM 5.14



The redox reaction that provides energy for the lithium battery described in the

Chemistry in Action on p. 147 is 2 Li1s2 + I21s2 S 2 LiI1aq2. Identify which

reactant is being oxidized and which is being reduced in this reaction.



SECTION 5.6



Redox Reactions



147



CHEMISTRY

IN ACTION

Batteries

Imagine life without batteries: no cars (they do not start very

easily without their batteries!), no heart pacemakers, no flashlights, no hearing aids, no laptops, no radios, no cell phones,

nor thousands of other things. Modern society could not exist

without batteries.

Although they come in many types and sizes, all batteries

work using redox reactions. In a typical redox reaction carried

out in the laboratory—say, the reaction of zinc metal with Ag+

to yield Zn2+ and silver metal—the reactants are simply mixed in

a flask and electrons are transferred by direct contact between

the reactants. In a battery, however, the two reactants are kept

in separate compartments and the electrons are transferred

through a wire running between them.

The common household battery used for flashlights and

radios is the dry cell, developed in 1866. One reactant is a can

of zinc metal, and the other is a paste of solid manganese dioxide. A graphite rod sticks into the MnO2 paste to provide electrical contact, and a moist paste of ammonium chloride separates

the two reactants. If the zinc can and the graphite rod are connected by a wire, zinc sends electrons flowing through the wire

toward the MnO2 in a redox reaction. The resultant electrical

current can then be used to power a lightbulb or a radio. The

accompanying figure shows a cutaway view of a dry-cell battery.



▲ Think



of all the devices we use every day—laptop computers, cell phones, iPods—that depend on batteries.



Closely related to the dry-cell battery is the familiar alkaline

battery, in which the ammonium chloride paste is replaced by

an alkaline, or basic, paste of NaOH or KOH. The alkaline battery has a longer life than the standard dry-cell battery because

the zinc container corrodes less easily under basic conditions.

The redox reaction is:

Zn1s2 + 2 MnO21s2 h ZnO1aq2 + Mn2O31s2

The batteries used in implanted medical devices such as

pacemakers must be small, corrosion-resistant, reliable, and able

to last up to 10 years. Nearly all pacemakers being implanted

today—about 750,000 each year—use titanium-encased,

lithium–iodine batteries, whose redox reaction is:



Zn1s2 + 2 MnO21s2 + 2 NH4Cl1s2 h

ZnCl21aq2 + Mn2O31s2 + 2 NH31aq2 + H2O1l2



2 Li1s2 + I21s2 h 2 LiI1aq2

See Chemistry in Action Problems 5.65 and 5.66 at the

end of the chapter.

e−



+



Insulator

Graphite rod

MnO2 and

carbon black paste

NH4Cl and ZnCl2

paste (electrolyte)

Zinc metal can





▲A



e−



dry-cell battery. The cutaway view shows the two reactants that make up the redox

reaction.



148



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Classification and Balancing of Chemical Reactions



5.7 Recognizing Redox Reactions

How can you tell when a redox reaction is taking place? When ions are involved, it is

simply a matter of determining whether there is a change in the charges. For reactions

involving metals and nonmetals, we can predict the gain or loss of electrons as discussed previously. When molecular substances are involved, though, it is not as obvious. Is the combining of sulfur with oxygen a redox reaction? If so, which partner is the

oxidizing agent and which is the reducing agent?

S1s2 + O 21g2 h SO 21g2



Electronegativity, or the propensity of an atom in a covalent bond to

attract electrons, was introduced in

Section 4.9.



Oxidation number A number that

indicates whether an atom is neutral,

electron-rich, or electron-poor.



One way to evaluate this reaction is in terms of the oxygen gain by sulfur, indicating that S atoms are oxidized and O atoms are reduced. But can we also look at this

reaction in terms of the gain or loss of electrons by the S and O atoms? Because oxygen

is more electronegative than sulfur, the oxygen atoms in SO 2 attract the electrons in

the S—O bonds more strongly than sulfur does, giving the oxygen atoms a larger share

of the electrons than sulfur. By extending the ideas of oxidation and reduction to an

increase or decrease in electron sharing instead of complete electron transfer, we can

say that the sulfur atom is oxidized in its reaction with oxygen because it loses a share

in some electrons, whereas the oxygen atoms are reduced because they gain a share in

some electrons.

A formal system has been devised for keeping track of changes in electron sharing, and thus for determining whether atoms are oxidized or reduced in reactions. To

each atom in a substance, we assign a value called an oxidation number (or oxidation

state), which indicates whether the atom is neutral, electron-rich, or electron-poor. By

comparing the oxidation number of an atom before and after a reaction, we can tell

whether the atom has gained or lost shares in electrons. Note that oxidation numbers

do not necessarily imply ionic charges. They are simply a convenient device for keeping

track of electrons in redox reactions.

The rules for assigning oxidation numbers are straightforward:





An atom in its elemental state has an oxidation number of 0.

Oxidation number

0



Na







Na+



Review the Important Points

about Ion Formation and the Periodic

Table listed in Section 3.6.



Br2



A monatomic ion has an oxidation number equal to its charge.

Oxidation number

+1







H2



Oxidation number

+2



Oxidation number

−1



Cl−



Ca2+



Oxidation number

−2



O 2−



In a molecular compound, an atom usually has the same oxidation number

it would have if it were a monatomic ion. Recall from Chapters 3 and 4 that

the less electronegative elements (hydrogen and metals) on the left side of the

periodic table tend to form cations, and the more electronegative elements (oxygen, nitrogen, and the halogens) near the top right of the periodic table tend to

form anions. Hydrogen and metals therefore have positive oxidation numbers in

most compounds, whereas reactive nonmetals generally have negative oxidation

numbers. Hydrogen is usually + 1, oxygen is usually - 2, nitrogen is usually - 3,

and halogens are usually - 1:

+1



H



−1



Cl



+1



H



−2



O



+1



H



+1



H



−3



+1



N



H



H



+1



SEC TION 5.7



Recognizing Redox Reactions



For compounds with more than one nonmetal element, such as SO 2, NO, or CO 2,

the more electronegative element—oxygen in these examples—has a negative oxidation number and the less electronegative element has a positive oxidation number.

Thus, in answer to the question posed at the beginning of this section, combining

sulfur with oxygen to form SO 2 is a redox reaction because the oxidation number

of sulfur increases from 0 to + 4 and that of oxygen decreases from 0 to - 2.

−2



O







−2



+4



S



+2



O



−2



N



O



−2



O



−2



+4



C



O



The sum of the oxidation numbers in a neutral compound is 0. Using this rule,

the oxidation number of any atom in a compound can be found if the oxidation

numbers of the other atoms are known. In the SO 2 example just mentioned, each

of the 2 O atoms has an oxidation number of - 2, so the S atom must have an oxidation number of + 4. In HNO 3, the H atom has an oxidation number of + 1 and

the strongly electronegative O atom has an oxidation number of - 2, so the N atom

must have an oxidation number of + 5. In a polyatomic ion, the sum of the oxidation numbers equals the charge on the ion.

+1



H



−2



+5



O



N



−2



Total = 1 + 5 + 3(−2) = 0



O



−2



O



Worked Examples 5.11 and 5.12 show further instances of assigning and using

oxidation numbers.



Worked Example 5.11 Redox Reactions: Oxidation Numbers

What is the oxidation number of the titanium atom in TiCl 4? Name the compound

using a Roman numeral (Section 3.10).

SOLUTION



Chlorine, a reactive nonmetal, is more electronegative than titanium and has an

oxidation number of - 1. Because there are 4 chlorine atoms in TiCl 4, the oxidation

number of titanium must be + 4. The compound is named titanium(IV) chloride.

Note that the Roman numeral IV in the name of this molecular compound refers to

the oxidation number + 4 rather than to a true ionic charge.



Worked Example 5.12 Redox Reactions: Identifying Redox Reactions

Use oxidation numbers to show that the production of iron metal from its ore

1Fe 2O 32 by reaction with charcoal (C) is a redox reaction. Which reactant has been

oxidized, and which has been reduced? Which reactant is the oxidizing agent, and

which is the reducing agent?

2 Fe 2O 31s2 + 3 C1s2 h 4 Fe1s2 + 3 CO 21g2



SOLUTION



The idea is to assign oxidation numbers to both reactants and products and see

if there has been a change. In the production of iron from Fe 2O 3, the oxidation

number of Fe changes from + 3 to 0, and the oxidation number of C changes from

0 to + 4. Iron has thus been reduced (decrease in oxidation number), and carbon

has been oxidized (increase in oxidation number). Oxygen is neither oxidized nor

reduced because its oxidation number does not change. Carbon is the reducing

agent, and Fe 2O 3 is the oxidizing agent.

+3



−2



0



2 Fe2O3 + 3 C



0



+4 −2



4 Fe + 3 CO2



149



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Classification and Balancing of Chemical Reactions

PROBLEM 5.15



What are the oxidation numbers of the metal atoms in the following compounds?

Name each, using the oxidation number as a Roman numeral.

(a) VCl 3

(b) SnCl 4

(c) CrO 3

(d) Cu1NO 322

(e) NiSO 4

PROBLEM 5.16



Assign an oxidation number to each atom in the reactants and products shown here

to determine which of the following reactions are redox reactions:

(a) Na 2S1aq2 + NiCl 21aq2 h 2 NaCl1aq2 + NiS1s2

(b) 2 Na1s2 + 2 H2O1l2 h 2 NaOH1aq2 + H21g2

(c) C1s2 + O 21g2 h CO 21g2

(d) CuO1s2 + 2 HCl1aq2 h CuCl 21aq2 + H2O1l2

(e) 2 MnO 4-1aq2 + 5 SO 21g2 + 2 H2O1l2 h

2 Mn2+1aq2 + 5 SO 42-1aq2 + 4 H+1aq2

PROBLEM 5.17



For each of the reactions you identified as redox reactions in Problem 5.16, identify

the oxidizing agent and the reducing agent.



5.8 Net Ionic Equations

In the equations we have been writing up to this point, all the substances involved in

reactions have been written using their full formulas. In the precipitation reaction of

lead(II) nitrate with potassium iodide mentioned in Section 5.3, for example, only the

parenthetical aq indicated that the reaction actually takes place in aqueous solution,

and nowhere was it explicitly indicated that ions are involved:

Pb1NO 3221aq2 + 2 KI1aq2 h 2 KNO 31aq2 + PbI 21s2



Ionic equation An equation in which

ions are explicitly shown.



Spectator ion An ion that appears

unchanged on both sides of a reaction

arrow.

Net ionic equation An equation that

does not include spectator ions.



In fact, lead(II) nitrate, potassium iodide, and potassium nitrate dissolve in water

to yield solutions of ions. Thus, it is more accurate to write the reaction as an ionic

equation, in which all the ions are explicitly shown:

An ionic equation:



Pb2+1aq2 + 2 NO 3-1aq2 + 2 K+1aq2 + 2 I-1aq2 h

2 K+1aq2 + 2 NO 3-1aq2 + PbI 21s2



A look at this ionic equation shows that the NO 3- and K+ ions undergo no change

during the reaction. They appear on both sides of the reaction arrow and act merely

as spectator ions, that is, they are present but play no role. The actual reaction, when

stripped to its essentials, can be described more simply by writing a net ionic equation,

which includes only the ions that undergo change and ignores all spectator ions:

Ionic equation:

Net ionic equation:



Pb2+1aq2 + 2 NO 3-1aq2 + 2 K+1aq2 + 2 I-1aq2 h

2 K+1aq2 + 2 NO 3-1aq2 + PbI 21s2

Pb2+1aq2 + 2 I-1aq2 h PbI 21s2



Note that a net ionic equation, like all chemical equations, must be balanced both

for atoms and for charge, with all coefficients reduced to their lowest whole numbers.

Note also that all compounds that do not give ions in solution—all insoluble compounds and all molecular compounds—are represented by their full formulas.

We can apply the concept of ionic equations to acid–base neutralization reactions and

redox reactions as well. Consider the neutralization reaction between KOH and HNO3:

KOH1aq2 + HNO 31aq2 h H2O1l2 + KNO 31aq2



Since acids and bases are identified based on the ions they form when dissolved in

aqueous solutions, we can write an ionic equation for this reaction:

Ionic equation:



K+1aq2 + OH-1aq2 + H+1aq2 + NO3-1aq2 h

H2O1l2 + K+1aq2 + NO 3-1aq2



SECTION 5.8



Eliminating the spectator ions (K+ and NO 3- ), we obtain the net ionic equation for

the neutralization reaction:

OH-1aq2 + H+1aq2 h H2O1l2



Net ionic equation:



The net ionic equation confirms the basis of the acid–base neutralization; the OHfrom the base and the H+ from the acid neutralize each other to form water.

Similarly, many redox reactions can be viewed in terms of ionic equations. Consider the reaction between Cu1s2 and AgNO 3 from Section 5.6:

Cu1s2 + 2 AgNO 31aq2 h 2 Ag +1aq2 + Cu1NO 3221aq2



The aqueous products and reactants can be written as dissolved ions:

Ionic equation:



Cu1s2 + 2 Ag +1aq2 + 2 NO 3-1aq2 h

2 Ag1s2 + Cu2+1aq2 + 2 NO 3-1aq2



Again, eliminating the spectator ions 1NO 3-2, we obtain the net ionic equation for

this redox reaction:

Net ionic equation:



Cu1s2 + 2 Ag +1aq2 h 2 Ag1s2 + Cu2+1aq2



It is now clear that the Cu1s2 loses 2 electrons and is oxidized, whereas each Ag +

ion gains an electron and is reduced.



Worked Example 5.13 Chemical Reactions: Net Ionic Reactions

Write balanced net ionic equations for the following reactions:

(a) AgNO 31aq2 + ZnCl 21aq2 h

(b) HCl1aq2 + Ca1OH221aq2 h

(c) 6 HCl1aq2 + 2 Al1s2 h 2 AlCl 31aq2 + 3 H21g2

SOLUTION



(a) The solubility guidelines discussed in Section 5.4 predict that a precipitate of

insoluble AgCl forms when aqueous solutions of Ag + and Cl- are mixed. Writing all the ions separately gives an ionic equation, and eliminating spectator ions

Zn2+ and NO 3- gives the net ionic equation.

Ionic equation:

Net ionic equation:



2 Ag +1aq2 + 2NO 3-1aq2 + Zn2+1aq2 + 2 Cl-1aq2 h

2 AgCl1s2 + Zn2+1aq2 + 2 NO 31aq2

2 Ag +1aq2 + 2 Cl-1aq2 h 2 AgCl1s2



The coefficients can all be divided by 2 to give:

Net ionic equation:



Ag +1aq2 + Cl+1aq2 h AgCl1s2



A check shows that the equation is balanced for atoms and charge (zero on each

side).

(b) Allowing the acid HCl to react with the base Ca1OH22 leads to a neutralization

reaction. Writing the ions separately, and remembering to write a complete formula for water, gives an ionic equation. Then eliminating the spectator ions and

dividing the coefficients by 2 gives the net ionic equation.

Ionic equation:

Net ionic equation:



2 H+1aq2 + 2Cl-1aq2 + Ca2+1aq2 + 2 OH-1aq2 h

2 H2O1l2 + Ca2+1aq2 + 2 Cl-1aq2

H+1aq2 + OH-1aq2 h H2O1l2



A check shows that atoms and charges are the same on both sides of the

equation.



Net Ionic Equations



151



152



CHAPTER 5



Classification and Balancing of Chemical Reactions



(c) The reaction of Al metal with acid (HCl) is a redox reaction. The Al is oxidized,

since the oxidation number increases from 0 S + 3, whereas the H in HCl is

reduced from + 1 S 0. We write the ionic equation by showing the ions that are

formed for each aqueous ionic species. Eliminating the spectator ions yields the

net ionic equation.



Ionic equation: 6 H+1aq2 + 6 Cl-1aq2 + 2 Al1s2 h

2 Al 3+1aq2 + 6 Cl-1aq2 + 3 H21g2

Net ionic equation: 6 H+1aq2 + 2 Al1s2 h 2 Al 3+1aq2 + 3 H21g2



A check shows that atoms and charges are the same on both sides of the

equation.



PROBLEM 5.18



Write net ionic equations for the following reactions:

(a) Zn1s2 + Pb1NO 3221aq2 h Zn1NO 3221aq2 + Pb1s2

(b) 2 KOH1aq2 + H2SO 41aq2 h K2SO 41aq2 + 2 H2O1l2

(c) 2 FeCl 31aq2 + SnCl 21aq2 h 2 FeCl 21aq2 + SnCl 41aq2

PROBLEM 5.19



Identify each of the reactions in Problem 5.18 as an acid–base neutralization,

a precipitation, or a redox reaction.



SUMMARY: REVISITING THE CHAPTER GOALS

1. How are chemical reactions written? Chemical equations

must be balanced; that is, the numbers and kinds of atoms must

be the same in both the reactants and the products. To balance

an equation, coefficients are placed before formulas but the

formulas themselves cannot be changed (see Problems 21–23,

26–37, 59, 60, 64, 67, 68, 71, 72, 75, 76, 79, 80).

2. How are chemical reactions of ionic compounds

classified? There are three common types of reactions of ionic

compounds (see Problems 38–50, 65, 70, 79, 81).

Precipitation reactions are processes in which an insoluble

solid called a precipitate is formed. Most precipitations take place

when the anions and cations of two ionic compounds change

partners. Solubility guidelines for ionic compounds are used

to predict when precipitation will occur (see Problems 24, 25,

43–46, 49, 69, 76–78).

Acid–base neutralization reactions are processes in which

an acid reacts with a base to yield water plus an ionic compound

called a salt. Since acids produce H+ ions and bases produce OHions when dissolved in water, a neutralization reaction removes

H+ and OH- ions from solution and yields neutral H2O (see Problems 37, 39, 75, 81).

Oxidation–reduction (redox) reactions are processes in which

one or more electrons are transferred between reaction partners.



An oxidation is defined as the loss of one or more electrons by

an atom, and a reduction is the gain of one or more electrons.

An oxidizing agent causes the oxidation of another reactant by

accepting electrons, and a reducing agent causes the reduction

of another reactant by donating electrons (see Problems 51–54,

57–62, 65, 66, 68, 82).

3. What are oxidation numbers, and how are they used?

Oxidation numbers are assigned to atoms in reactants and products to provide a measure of whether an atom is neutral, electronrich, or electron-poor. By comparing the oxidation number of an

atom before and after reaction, we can tell whether the atom has

gained or lost shares in electrons and thus whether a redox reaction has occurred (see Problems 51–62, 65, 66, 70–74, 82).

4. What is a net ionic equation? The net ionic equation only

includes those ions that are directly involved in the ionic reaction. These ions can be identified because they are found in

different phases or compounds on the reactant and product

sides of the chemical equation. The net ionic equation does

not include spectator ions, which appear in the same state on

both sides of the chemical equation (see Problems 39, 47, 48,

50, 69, 76–78, 81).



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