c 16hapter oxidation-reduction reactions

552

16CHAPTER

552

16CHAPTER

16.1 The Nature of Oxidation-Reduction ReactionsMiniLab 16.1 Corrosion of IronChemLab Copper Atoms and Ions:

Oxidation and Reduction

16.2 Applications of Oxidation-Reduction ReactionsMiniLab 16.2 Testing for Alcohol

by Redox

Chapter PreviewSections Why Do Things Rust?

When iron corrodes, iron metal reacts withoxygen from the air and water to formiron(III) oxide—rust. Rust is the result of

an oxidation-reduction reaction in which ironmetal loses electrons to oxygen. Given time, andoxygen from the air and water, all of the drumsin this photo will rust away completely.

Why Do Things Rust?

Oxidation-ReductionReactionsOxidation-ReductionReactions

553

Observing an Oxidation–Reduction ReactionRust is the result of a reaction of iron and oxygen. Ironnails can also react with substances other than oxygen,as you will find out in this experiment.

Safety Precautions

Always wear safety goggles and an apron in the laboratory.

Materials

• test tube• iron nail• steel wool or sandpaper• 1M copper(II) sulfate (CuSO4)

Procedure

1. Use a piece of steel wool to polish the end of an iron nail.

2. Add about 3 mL 1.0M CuSO4 to a test tube. Place thepolished end of the nail into the CuSO4 solution. Letstand and observe for about 10 minutes. Record yourobservations.

Analysis

What is the substance found clinging to the nail? Whathappened to the color of the copper(II) sulfate solution?Write the balanced chemical equation for the reactionyou observed.

Look through the Section Previewsfor this chapter, jotting down somekey ideas. As you read through thechapter, make an outline using thekey ideas you wrote down. For eachtopic, review any new vocabularywords.

Reading Chemistry

Review the following conceptsbefore studying this chapter.Chapter 3: patterns of valence electronsChapter 5: predicting oxidation number from the periodic tableChapter 6: types of reactions

What I Already Know

Preview this chapter’s content andactivities at chemistryca.com

Start-up ActivitiesStart-up Activities

http://chemistryca.com

16.1

Oxygen undergoes many reactions when it encounters other substances.One of these reactions is responsible for the browning of fruits. Anoth-er forms the rust that eats away at the metal parts of bikes and cars. In

both of these cases, a type of reaction called oxidation is taking place. Youcan probably guess how this reaction got its name; oxygen is a reactant. Butyou will learn that notall oxidation reactionsinvolve oxygen. Andoxidation reactionsare never lonelybecause they alwayshave partners—reduction reactions.You will see what thecharacteristics ofthese reactions areand why they alwaystake place together.

What is oxidation-reduction?Oxygen is the most abundant element in Earth’s crust. It is very reactive

and can combine with almost every other element. An element that bondsto oxygen to form a new compound, called an oxide, usually loses electronsbecause oxygen is more electronegative. You will recall that an electro-negative element has a strong attraction for electrons. Because of thisstrong attraction, oxygen is able to pull electrons away from other atoms.The reactions in which elements combine with oxygen to form oxides wereamong the first to be studied by early chemists, who grouped them togeth-er and called them oxidation reactions. Later, chemists realized that someother nonmetal elements can combine with substances in the same way asoxygen and that these reactions are similar to oxidation reactions. Modernchemists use the term oxidation to refer to any chemical reaction in whichan element or compound loses electrons to another substance.

A common oxidation reaction occurs when iron metal loses electronsto oxygen. Each year in the United States, corrosion of metals—especiallythe iron in steel—costs billions of dollars as automobiles, ships, andbridges and other structures are slowly eaten away. Figure 16.1 showssome of this damage and how it can be prevented.

The Nature of Oxidation-Reduction Reactions

SECTION

Objectives✓ Analyze the charac-teristics of an oxidation-reductionreaction.

✓ Distinguishbetween oxidationreactions and reduc-tion reactions by definition.

✓ Identify the sub-stances that are oxi-dized and those thatare reduced in a redoxreaction.

✓ Distinguish oxidiz-ing and reducingagents in redox reactions.

Review VocabularyBuffer: solution thatresists changes in pHwhen moderateamounts of acids orbases are added.

New Vocabularyoxidation-reduction

reactionoxidationreductionoxidizing agentreducing agent

SECTION PREVIEW

554 Chapter 16 Oxidation-Reduction Reactions

16.1 The Nature of Oxidation-Reduction Reactions 555

RedoxWhat happens to the zinc in galvanized steel? It reacts with oxygen to

form zinc oxide in the following reaction.

2Zn(s) � O2(g) ˇ 2Zn2�(s) � 2O2�(s)

Does this type of reaction look familiar? You learned in Chapter 6 that thisis classified as a synthesis reaction. You also know that early chemists calledit an oxidation reaction because oxygen is a reactant. The formation of zincoxide falls into another, broader class of reactions characterized by thetransfer of electrons from one atom or ion to another. This type of reactionis called an oxidation-reduction reaction, commonly known as a redoxreaction. Many important chemical reactions are redox reactions. Forma-tion of rust is one example; combustion of fuels is another. In each redoxreaction, one element loses electrons, and another element takes them.

How do atoms or ions lose electrons in a redox reaction? If you examinethe equation for the reaction between zinc and oxygen more closely, you cansee which atoms are gaining electrons and which are losing them. You alsocan determine where the electrons go during a redox reaction by comparingthe oxidation number of each type of atom or ion before and after the reac-tion takes place. Recall from Chapter 5 that the oxidation number of an ionis equal to its charge. All elements, when in their free form, have a charge ofzero and are assigned an oxidation number of zero. In the formation of zincoxide, the zinc atom and the diatomic oxygen molecule that react each hasan oxidation number of zero. In the ionic compound formed, each oxideion has a 2� charge and an oxidation number of 2�. Because the com-pound must be neutral, the total positive charge must be 4�; thus, each zincion must have a charge and an oxidation number of 2�.

� Steel can be protected from oxidation if it is coated with amore active metal such as zinc. Zinc loses electrons to oxygen more readily than iron does, so the zinc is oxi-dized preferentially, forming a tough protective layer ofzinc oxide. The coating of zinc and zinc oxide preventsthe formation of rust by keeping oxygen from reachingthe iron. Steel that has been coated with zinc is called gal-vanized steel. The bucket on the left has been galvanized.

Figure 16.1Corrosion of IronWhen iron corrodes, iron metal reacts with oxygen toform iron(III) oxide—rust. Corrosion of iron can be prevented by covering the surface of exposed steel withpaint or other coatings such as plastic. If the protectivecoating is damaged or cracked, rust forms quickly. �


OxidationYou have learned that a reaction in which an element loses electrons is

called an oxidation reaction. The element that loses the electrons becomesmore positively charged; that is, its oxidation number increases. That ele-ment is said to be oxidized during the reaction. Zinc is oxidized during theformation of zinc oxide because metallic zinc atoms each lose two electrons.The oxidation reaction can be written by itself to show how zinc changesduring the redox reaction. Here’s what happens to each atom of zinc.

ReductionWhat happens to the electrons that are lost by the zinc atom? Electrons

do not wander around by themselves; they must be transferred to anotheratom or ion. This is why oxidation reactions never occur alone. They arealways paired with reduction reactions. A reduction reaction is one inwhich an element gains one or more electrons. The element that picks upthe electrons and becomes more negatively charged during the reaction issaid to be reduced. Its oxidation number decreases, or is reduced. Becauseoxidation and reduction reactions occur together, each is referred to as ahalf-reaction.

In every redox reaction, at least one element undergoes reduction whileanother undergoes oxidation. Just as a successful pass in football requires aquarterback to throw the ball and a receiver to catch it, a redox reactionmust have one element that gives up electrons and one that accepts them.The electronic structure of both reactants changes during a redox reaction.

Figure 16.2 shows the movement of electrons in the formation of zincoxide. Oxygen accepts the electrons that zinc loses. Oxygen is reducedduring the reaction between zinc and oxygen because each oxygen atomgains two electrons. Like the oxidation reaction, the reduction reactioncan be written by itself. Here’s what happens to each atom of oxygen.

O O2–ˇ� 2e– (gain of electrons)

2+[Zn]Zn 2e–ˇ � (loss of electrons)

Figure 16.2Formation of Zinc OxideIn the formation of zincoxide, the zinc atom loses twoelectrons during the reaction,becoming a zinc ion. Its oxi-dation number increases fromzero to 2�. Theoxygen atom gainsthe two electronsfrom zinc, becom-ing an oxide ion.Its oxidation num-ber decreases fromzero to 2�.

Zincatom

Oxygenatom

2e�

2e�

8e� 2e� 6e�18e�

�

Zincion

Oxideion

2e� 8e� 2e� 8e�18e�

Zn � [Zn]2� � 2��

�

reduction:re (L) backducere (L) to lead

In a reductionreaction, the addi-tion of electronsresults in adecrease in oxida-tion number of anatom or ion.

Corrosion of IronCorrosion is the term generally used to describe the oxidation of a

metal during its interaction with the environment. In this MiniLab, youwill study the corrosion of a nail and determine the factors that affectthis process.

Procedure 1. Dissolve a package of clear,

unflavored gelatin in about 200 mL of warm water. Stir in 2 mL of phenolphthalein solu-tion and 2 mL of potassiumhexacyanoferrate(III) solution.Pour the prepared solution intoa widemouth glass jar or petridish to a depth of about 1 cm.

2. In the liquid gelatin, place aplain iron nail, an aluminumnail, a galvanized iron nail, anda painted iron nail of the typeused for paneling. Space thenails far apart.

3. Label the jar or petri dish withyour name, and leave it for sev-eral hours or overnight. Handlethe jar or dish carefully untilthe gelatin has set.

4. Record your observationsregarding any interactions ofthe nails with the substances inthe gelatin.

1


Combining the Half-Reactions The equation for the reduction half-reaction shows one atom of oxygen

reacting. However, oxygen is not found in nature as single atoms; twoatoms combine to form a diatomic molecule of O2. The reduction equa-tion must be multiplied by 2 to reflect this. Thus, the balanced equationfor the reduction reaction is written as follows.

Note that four electrons are gained by the oxygen molecule. To producethose four electrons, two atoms of zinc must take part in the reaction.Therefore, the balanced equation for the oxidation reaction must be writ-ten as follows.

2+2[Zn]2Zn 4e–ˇ �

O O2–

2 ˇ� 4e 2–

Analysis

1. Which of the nails havereacted with the substancesin the gelatin? What is theevidence of corrosion?

2. If any of the nails have notcorroded in the solution, canyou suggest a reason whythey haven’t? What methodsare commonly used to pre-vent or minimize corrosion?

3. Any blue color in the gelatinis due to the formation ofiron(II) ions and their inter-action with the hexacyano-ferrate(III) ion. Any pink orred color in the gelatin is dueto the gaining of electrons byoxygen and water molecules,forming basic hydroxide ionsthat turn the phenolph-thalein pink. Which of thesereactions is oxidation, andwhere does it occur on thereacting nail?


The balanced overall equation for the reaction now can be written asshown in Figure 16.3. This equation is the same as the equation for theformation of zinc oxide that you read at the beginning of the discussionon redox reactions. Now you know that it represents the net oxidation-reduction reaction and is the sum of an oxidation half-reaction and areduction half-reaction.

If an element is gaining electrons, why is this called a reduction reac-tion? After all, you don’t gain weight when you reduce. You have learnedthat the reason is because there is a reduction in the charge or oxidationnumber of an atom of the substance that is reduced. An older, historicreason for the use of the term reduction is that the name was first appliedto processes in furnaces in which metals are isolated from their ores athigh temperatures, Figure 16.4. During these processes, oxygen isremoved from the ores in which it is combined with the metal, so the oreis reduced to the free metal. There is a reduction in the amount of solidmaterial and a considerable decrease in volume.

Figure 16.4Furnaces: Old and ModernFor thousands of years, metals have been used by many different culturesfor making jewelry, cookware, and weapons. Because metals are normallyfound combined with other elements as ores, furnaces operating at hightemperatures are used to separate the free metal from other elements. Thepositively charged metal ions in the ore are reduced to the elemental state,while oxygen and other negatively charged elements in the ore are oxi-dized. Iron smelting in medieval England is shown here. �

Figure 16.3Summarizing the ReactionTwo zinc atoms combine withone diatomic oxygen mole-cule to form two formulaunits of zinc oxide. Becausezinc loses four electrons in theoxidation reaction and oxy-gen gains four electrons inthe reduction reaction, allelectrons are accounted for;the two half-reactions are balanced.

� A modern industrial iron blastfurnace is shown here.

2+22Zn 2Zn 2OO� �ˇ

�

�

2

from Zn

to O

2–0 0

0+4e

�

�

–4e

Each atom of zinc donates 2e to an oxygen atom and is oxidized.

Each atom of oxygen accepts 2e from a zinc atom and is reduced.


Identifying a Redox ReactionThe oxidation of zinc is a redox reaction in which oxygen is a reactant.

You have learned that elements other than oxygen can accept electrons andbecome reduced during redox reactions. You are already familiar with theexplosive reaction in which sodium and chlorine combine to form table salt.

2Na(s) � Cl2(g) ˇ 2NaCl(s)

Are electrons transferred during this reaction? Yes, because each sodiumatom loses one electron to become a sodium ion with a charge of 1�. Theoxidation number of sodium increases from 0 to 1�. Each chlorine atomgains one electron to form a chloride ion. The oxidation number of chlo-rine decreases from 0 to 1�. Therefore, this is another example of a redoxreaction.

Oxidizing and Reducing AgentsAnother redox reaction that doesn’t involve oxygen occurs when a strip

of zinc metal is placed in a solution of copper(II) sulfate. The progress ofthis reaction can be followed easily because a readily observable changetakes place. As shown in Figure 16.5, copper metal quickly begins to formon the zinc strip.

+200

Cl2Na 2Na 2Cl–ˇ� �

2e– ˇ

Figure 16.5The Reaction Between Zincand Copper(II) SulfateA blue copper(II) sulfate solu-tion gradually becomes color-less if a strip of zinc metal isplaced in it. The zinc gives upelectrons, becoming oxidizedto zinc ions. The colorlessZn2� ions that form go intosolution. The Cu2� ions pickup electrons from zinc andbecome reduced to coppermetal atoms, which aredeposited on the strip. �

2+2+ Zn ZnCuCu

2+Cu

� �

2+Zn

ˇ

ˇ

ˇ ˇ

ˇ

�

0

Zn0

0

Cu02e �2e

oxidizingagent

reduced to

reduced to

reducingagent

oxidized to

oxidized to

ˇ

ˇ

Cu2� is the oxidizing agent, andZn0 is the reducing agent. �


Copper Atoms and Ions: Oxidation and Reduction

Copper atoms and ions often take part in reac-tions by losing or gaining electrons, which areoxidation and reduction, respectively. If copperatoms lose electrons to form positive ions, copperis oxidized. Other atoms or ions must gain theelectrons that copper atoms lose. These atoms orions are reduced and are called oxidizing agents.In this ChemLab, you will observe two reactionsthat involve the oxidation or reduction of copper.

ProblemWhat are some typical reactions that involve

the oxidation or reduction of copper?

Objectives

•Observe reactions that involve the oxidation orreduction of copper.

•Classify the reactants as substance oxidized,reducing agent, substance reduced, and oxidiz-ing agent.

Materialscopper(II) oxidepowdered charcoal (carbon)weighing paperbalancelarge Pyrex or Kimax test tubes (2)1-hole rubber stopper fitted with glass tube with

bend as shown150-mL beakers (2)graduated cylinder, smallglass stirring rodBunsen or Tirrill burnerring standtest-tube clamp

thermal glovelimewater (calcium hydroxide solution)

Safety Precautions

Care should be taken in handling hot objectsand when working around open flames. Do notbreathe in the fumes that are produced duringthe teacher demonstration in step 1.

1. Teacher Demonstration Your teacher willperform this reaction as a demonstration eitherin the fume hood or outside the building.CAUTION: Do not perform this procedure byyourself. A 1-cm squareof copper foil will beplaced in a porcelainevaporating dish. First,5 mL of water, then 5 mL of concentratedHNO3 will be added.Note the color of theevolved gas and thecolor of the resulting solution. Record yourobservations in a table similar to the one underData and Observations.

2. On a piece of weighing paper, thoroughly mixapproximately 1 g of copper(II) oxide withtwice its volume of powdered charcoal. Placethe mixture in a clean, dry Pyrex or Kimax testtube. Add about 10 mL of limewater to a sec-ond test tube, and stand it in a 150-mL beaker.Assemble the apparatus as shown here, withthe copper-oxide test tube sloped slightlydownward and the delivery tube extendinginto the limewater.


3. Heat the mixture in the test tube, gently at firstand then strongly. As soon as you notice achange in the limewater, carefully remove thestopper and delivery tube from the reactiontest tube. CAUTION: Do not stop heating aslong as the tube is in the limewater. Record yourobservations of the limewater in your table.

4. Continue heating the reaction test tube until aglow spreads throughout the reactant mixture.Turn off the burner.

5. After the reaction test tube has cooled to near-ly room temperature, empty the contents intoa beaker that is about half full of water. In asink, slowly stir the mixture while runningwater into the beaker until all the unreactedcharcoal has washed away. Observe the prod-uct that remains in the beaker, and recordyour observations.

1. Interpreting Data What evidence of chemicalchange did you observe in each reaction?

2. Interpreting Data In the first reaction, a blue-colored solution indicates the presence ofCu2� ions, a brown gas is NO2, and a colorless

gas is NO. In the second reaction, if limewaterbecomes cloudy and white, carbon dioxide gashas reacted with the calcium hydroxide toform insoluble calcium carbonate. Using thisinformation, analyze your data and observa-tions. Determine which reactants (Cu andHNO3 for the first reaction, CuO and C forthe second reaction) were oxidized and whichwere reduced in each reaction.

3. Classifying Classify each of the four reactantsas an oxidizing agent or a reducing agent.

1. The mass of copper produced in the secondreaction is less than the mass of the reactingcopper(II) oxide. Why, then, is the gain ofelectrons known as reduction?

2. What mass of copper may be produced fromthe reduction of 1.000 metric ton ofcopper(II) oxide? Hint: Determine the formu-la mass of copper(II) oxide.

3. If the chlorine gas used at a water-treatmentplant reacts with organic materials in thewater to yield chloride ions, how would youclassify the chlorine gas in terms of oxidationand reduction?

Observations

Step 1: Gas

Solution

Step 3: Limewater

Step 5: Product


What role do the copper ions play in the redox reaction? Each copperion is reduced to uncharged copper metal when it accepts electrons fromthe zinc metal. Because the copper ion is the agent that oxidizes zincmetal to the zinc ion, Cu2� is called an oxidizing agent. An oxidizingagent is the substance that gains electrons in a redox reaction. The oxidiz-ing agent is the material that’s reduced. Because oxidation and reductiongo hand in hand, a reducing agent must be present. Zinc metal is theagent that supplies electrons and reduces the copper ion to copper metal;therefore, zinc is called the reducing agent. A reducing agent is the sub-stance that loses electrons in a redox reaction. The reducing agent is thematerial that’s oxidized. Figure 16.6 summarizes the roles of oxidizingand reducing agents in redox reactions.

Understanding Concepts1. Name and define the two half-reactions that

make up a redox reaction.

2. Identify which reactant is reduced and which isoxidized in each of the following reactions.

a) C5H12(l) � 8O2(g) ˇ 5CO2(g) � 6H2O(g)b) 2Al(s) � 3Cu2+(aq) ˇ 2Al3�(aq) � 3Cu(s)c) 2Cr3�(aq) � 3Zn(s) ˇ 2Cr(s) � 3Zn2�(aq)d) 2Au3�(aq) � 3Cd(s) ˇ 2Au(s) � 3Cd2�(aq)

3. What is the oxidizing agent when iron rusts?What is the reducing agent?

Thinking Critically4. Applying Concepts The following equation

represents the reaction between an acid and a

base to form a salt and water. Determine theoxidation number for each element. Is this aredox reaction? Explain.

2KOH(aq) � H2SO4(aq) ˇK2SO4(aq) � 2H2O(l)

Applying Chemistry5. Antioxidants Compounds that are easily oxi-

dized can act as antioxidants to prevent othercompounds from being oxidized. Vitamins C andE protect living cells from oxidative damage byacting as antioxidants. Why does adding lemonjuice to fruit salad prevent browning of the fruit?

SECTION REVIEW

Figure 16.6Oxidizing and Reducing AgentsWhen electrons are transferred from one element to another, a combinationof an electron-gaining—or reduction—reaction and an electron-losing—or oxidation—reaction takes place. This combination is called a redox reaction.The element that is reduced oxidizes another element by attracting electronsfrom it, so it is called an oxidizing agent. The element that is oxidized reducesthe first element by transferring electrons to it, so it is called a reducing agent.

� �ˇ

ˇ �e �e

�e �e

Oxidizingagent

Reducedoxidizing

agent

Oxidizedreducing

agentReducing

agent

Gainselectrons

Loseselectrons

For more practice with solvingproblems, see SupplementalPractice Problems,Appendix B.

chemistryca.com/self_check_quiz

http://chemistryca.com/self_check_quiz

16.2

Natural redox reactions are going on around you every day, everywhere.This is partly due to the abundance of oxygen, which acts as the oxidiz-ing agent as it is reduced in some redox reactions. Other oxidizing

agents take part in different redox reactions, especially in environmentswhere not much oxygen gas is found. Near the vents of volcanoes, where sul-fur compounds explode out from deep within Earth, enormous deposits ofsolid yellow sulfur are found. The element sulfur acts both as an oxidizingagent and as a reducing agent in the reaction that forms the sulfur deposits.Can you tell which sulfur compound serves each function in this reaction?

2H2S(g) � SO2(g) ˇ 3S(s) � 2H2(g) � O2(g)

Note that more than one element in a reaction can be oxidized or reduced.The sulfur in hydrogen sulfide and the oxygen in sulfur dioxide both are oxi-dized. Sulfur in sulfur dioxide and hydrogen in hydrogen sulfide both arereduced. Each reactant acts as both a reducing agent and an oxidizing agent.

Applications of Oxidation-Reduction Reactions

SECTION PREVIEW

SECTION

16.2 Applications of Oxidation-Reduction Reactions 563

Say Cheese: Redox in PhotographyUnderstanding natural redox reactions such as the one that occurs in

sulfur volcanoes has allowed chemists to develop many processes thatmake use of oxidation and reduction reactions. Without them, photo-graphs or steel wouldn’t exist, and stains would be much harder toremove from clothing.

Objectives✓ Analyze commonredox processes toidentify the oxidizingand reducing agents.

✓ Identify some redoxreactions that takeplace in living cells.

Review VocabularyOxidation: reactionin which an elementloses electrons.


Leonardo da Vinci described a primitive “camera” before 1519, in whichsomeone had to trace images focused on a glass plate inside a box. How-ever, it wasn’t until 1838 that the French inventor L.J.M. Daguerre success-fully fixed the images in a camera on highly polished, silver-plated copperto make the first photographs. These early photographs were calleddaguerreotypes in his honor, Figure 16.7.

Modern photographic film is made of a plastic backing covered with alayer of gelatin, in which millions of grains of silver bromide are embed-ded. When light strikes a grain, silver and bromide ions are converted intotheir elemental forms through a redox reaction. The equation for thisredox reaction is as follows.

2Ag� � 2Br� ˇ 2Ag � Br2

The reaction begins when the shutter on a camera is opened. Lightfrom the scene being photographed passes through the camera’s lens andshutter and strikes the light-sensitive silver bromide on the film. The lightenergy causes electrons to be ejected from a few of the bromide ions, oxi-dizing them to elemental bromine. The electrons are transferred to silverions, reducing them to metallic silver atoms. These grains are now activat-ed. The developing chemicals continue the redox reaction by causing theactivated grains to be converted to metallic silver. In areas where the lightis brightest, more grains are activated, and after developing, they becomethe darker areas. No silver atoms form in areas of the film that are notstruck by light, and that part of the film remains transparent. The exposedfilm is then developed into a negative, during which time the remainingAgBr and Br2 is washed away. Figure 16.8 describes the developing andprinting processes.

Figure 16.7Early PhotosIn a daguerreotype, a redoxreaction between silver andiodine fumes produced a layerof light-sensitive silver iodideon the surface of the polishedphotographic plate. Exposureto light caused decompositionof the silver iodide into ele-mental silver, which was thentreated with the fumes ofheated mercury to formbright amalgam areas. Theimage of Paris shown herewas made by Daguerre him-self.

photograph:photos (GK) lightgraphein (GK) towrite

Light is used torecord the imageof an object in aphotograph.

Figure 16.8Developing and Printing PicturesMaking photographic negatives by developing exposed film involves several steps. The process describes howblack-and-white pictures are made. For color photos, light-sensitive dyes are combined with the silver bromide inlayers on the film.

� 1. The exposed film is transferred to a canister, where it is devel-oped using a solution of a reducing agent, or developer. Theorganic compound hydroquinone is usually used for this pur-pose. The developer reduces all the silver ions to silver atoms inany grain of silver bromide that was hit by light, but it does notreact with silver ions in grains that were not exposed to light.Because metallic silver is dark and silver bromide is light, animage having light and dark areas is produced.

� 2. After the film has been developed, a solution of a fixer contain-ing thiosulfate ions is added. Thiosulfate ions react with unre-duced silver ions to form a soluble complex, which is washedaway. This prevents unreduced silver ions from becomingreduced and darkening slowly over time. The reaction follows.AgBr(s) � 2S2O3

2�(aq) ˇ [Ag(S2O3)2]3�(aq) � Br�(aq)

� 3. The fixed film is washed to remove any remaining developer orfixer solution. The photographic negative is the reverse of theimage photographed; that is, light areas in the scene are darkon the film, and vice versa.

� 4. When light is shonethrough the negative ontolight-sensitive photographicpaper, a photographic printis made. The print is posi-tive; light and dark areasare identical to those in the scene.



PHYSICS CONNECTION

Solid Rocket Booster EnginesIf you have ever built and launched a model rocket,

you probably noticed that the rocket engine was made ofa solid, highly combustible material packed into a card-board tube. After ignition, the expansion and expulsionof the gases produced enough downward force to launchthe lightweight rocket quickly into the air. Space shuttlesuse a similar type of technology, but on a much largerscale.

Engine systems The space shuttle has two differentengine systems. The three main engines attached directlyto the shuttle operate on liquid hydrogen and liquid oxygen reservoirs carried in the large, centrally locateddisposable fuel tank. The two smaller, reusable, strap-onbooster rockets on either side of the main fuel tank areloaded with a solid fuel, which undergoes a powerful,thrust-producing, oxidation-reduction reaction thathelps boost the shuttle into orbit.

Solid rocket fuel The solid rocket fuel is a mixture containing 12 percent aluminum powder, 74 percentammonium perchlorate, and 12 percent polymer binder. Once ignited, the engine cannotbe extinguished. The extremely reactive ammonium perchlorate supplies oxygen to theeasily oxidized aluminum powder, providing a greatly exothermic and fast reaction. Thepurpose of the polymer binder is to hold themixture together and to help it burn evenly.The overall redox reaction is shown here.

Shuttle forces Each solid rocket boosterweighs 591 000 kg at liftoff, produces 11.5million N of force, and operates for abouttwo minutes into the flight. For comparison,a 1000-kg car accelerating from 0 to 26.8 m/s(60 mph) in 7 seconds would require a forceof only 3830 N. The tremendous release ofchemical energy and expansion of hot gasesdue to the oxidation-reduction reactionthrough the engine of the solid rocket boost-er produces the tremendous thrust needed toget the 2 million-kg shuttle from 0 to almost700 m/s (1500 mph) in just 132 seconds.

2 3443NH ClO 43NH Cl8Al 4Al O� �

24e– ˇ

ˇ

Connecting to Chemistry

1. Applying Pow-dered aluminum isused in anothergreatly exothermicreaction, the ther-mite reaction,which is used forwelding metals. Thereaction is asshown.

2Al � Fe2O3 ˇAl2O3 � 2Fe

What role does thepowdered aluminumplay in this reaction?

2. Acquiring Informa-tion Investigate thelives and research ofRobert Goddardand Werner VonBraun, who bothexperimented withrockets in the 1930sand helped guidethe United Statesinto the space age.Write a short reportabout these men.


� Molten iron is drawn off at the bot-tom of the furnace. A combinationof by-products known as slag is alsoremoved at the bottom.

Molten iron

Limestone,coke, andiron ore

Compressedair

Exhaustgases

Slag

The Stone, Bronze, andIron Ages are historicalperiods named afterthe most commonmaterial that was usedfor making tools duringeach time. The BronzeAge came before theIron Age because cop-per and tin, the ele-ments that are meltedtogether to form thealloy bronze, were bothwidely available andeasily accessible metals.Bronze is stronger thaneither copper or tinalone. The Iron Agecame later because ironis harder to reduce toelemental form. Itrequires smelting at ahigher temperaturethan bronze.

Having a Blast: Redox in a Blast FurnaceIron is seldom found in the elemental form needed to make steel.

Metallic iron must be separated and purified from iron ore—usuallyhematite, Fe2O3. This process takes place in a blast furnace in a series ofredox reactions. The major reaction in which iron ore is reduced to ironmetal uses carbon monoxide gas as a reducing agent.

First, a blast of hot air causes coke, a form of carbon, to burn, produc-ing CO2 and heat. Limestone, CaCO3, which is mixed with the iron ore inthe furnace, decomposes to form lime (CaO) and more carbon dioxide.The carbon dioxide then oxidizes the coke in a redox reaction to formcarbon monoxide, which is used to reduce the iron ore to iron. Theprocess is outlined here and illustrated in Figure 16.9.

Redox in Bleaching ProcessesBleaches can be used to remove stains from clothing. Where do the

stains go? Bleach does not actually remove the chemicals in stains fromthe fabric; it reacts with them to form colorless compounds. In chlorinebleaches, an ionic chlorine compound in the bleach reacts with the com-pounds responsible for the stain. This ionic compound is sodiumhypochlorite (NaOCl). The hypochlorite ions oxidize the molecules thatcause dark stains.

OCl�(aq) � stain molecule(s) ˇ Cl�(aq) � oxidized stain molecule(s)(colored) (colorless)

2–2

3+

3 2

03O (s)2Fe (s)

2CO(g)

3CO(g) 3CO (g)2Fe (l)� �

�

�

�

6e–

CaCO (s) CaO(s) CO (g)

2CO (g) C(s)

ˇ

ˇ

ˇ

ˇ

Figure 16.9Blast FurnaceIron ore (Fe2O3), coke(C), and limestone(CaCO3) are added atthe top of the furnace.Hot air at about 900°C,blasted into the bottomof the furnace, burns thecoke in an exothermicreaction. This reactioncauses temperatures in ablast furnace to reachabout 2000°C. �

Testing for Alcohol by RedoxOrganic alcohols react with orange dichromate ions, producing blue-

green chromium(III) ions. This reaction is used in a Breathalyzer test totest for the presence of alcohol in a person’s breath. In this MiniLab,you will use this reaction to test for the presence of alcohol in a numberof household hygiene, cosmetic, and cleaning products.

Procedure

1. Label five small test tubes withthe names of the products tobe tested.

CAUTION: Do not allowdichromate reagent to come intocontact with skin. Wash withlarge volumes of water if it does.

4. Observe and record any colorchanges that occur within oneminute.

Analysis1. Which of the products that you

tested contain alcohol? Was thepresence of alcohol noted onthe label of the products?

2. If the orange Cr2O72� ion

reacts with alcohol to producethe blue-green Cr3� ion, whatsubstance is the reducing agentin the reaction?

2


Bleaches containing hypochlorite should be used carefully becausehypochlorite is a powerful oxidizing agent that can damage delicate fab-rics. These bleaches usually have a warning label telling the user to test aninconspicuous part of the fabric before using the product. In addition toacting as a bleaching agent, hypochlorite ions are also used as disinfec-tants, as Figure 16.10 shows.

Figure 16.10Hypochlorite as a DisinfectantHypochlorite is used in disinfectants to killbacteria in swimming pools and in drinkingwater. In both cases, the hypochlorite ionsact as oxidizing agents. Bacteria are killedwhen important compounds in them aredestroyed by oxidation. In this photo, theamount of chlorine in the water is beingmonitored. Chlorine reacts with the water toform hypochlorite ions.

2. Place approximate-ly 1 mL of eachproduct in theappropriate tube.

3. Wearing apronand goggles, addthree drops ofdichromatereagent to eachtube, and stir tomix the solutions.

See page 870 inAppendix F for

Testing the OxidationPower of Bleach

Lab


Breathalyzer TestThe alcohol in beverages, hair spray, and mouthwashes is

ethanol. Ethanol is a volatile liquid that evaporates rapidly atroom temperature. Because of this volatility, drinking analcoholic beverage results in a level of gaseous ethanol in thebreath that is proportional to the level of alcohol in thebloodstream. About 50 percent of all automobile accidentsthat result in a fatality are caused by intoxicated drivers. Lawofficers can determine quickly whether a person is legallyintoxicated by using an instrument called a breath analyzer,or Breathalyzer.

1. Suppose a personused mouthwashshortly before takinga Breathalyzer test.What might be theresult?

2. How would the colorproduced in aBreathalyzer testchange as theethanol content ofthe blood increases?

Thinking Critically

1. A simple Breathalyzer device has aninflatable plastic bag attached to a tubecontaining an orange solution of potassi-um dichromate and sulfuric acid.

2. During a Breathalyzer test,a person blows into themouthpiece of the bag.

3. If alcohol vapors are presentin the person’s breath,ethanol undergoes a redoxreaction with the dichromate.As ethanol is oxidized, theorange Cr6� ions are reducedto blue-green Cr3� ions.

4. The exact color produced dependson the amount of alcohol in thebreath. The color change that is pro-duced during the test is comparedto standard color mixtures of thetwo chromium ions to get an esti-mate of the blood alcohol level.


Corrosion of MetalsDid you know that the Statue of Liberty is made of copper sheets

attached to a steel skeleton? Why does it appear green rather than the reddish-brown color of copper? When copper is exposed to humid airthat contains sulfur compounds, it undergoes a slow oxidation process.Under these conditions, the copper metal atoms each lose two electrons toproduce Cu2� ions, which form the compounds CuSO4�3Cu(OH)2 andCu2(OH)2CO3. These compounds are responsible for the green coat orpatina found on the surface of copper objects that have been exposed toair for long periods of time, Figure 16.11.

You have learned that iron is oxidized by oxygen in the air to form rust.Aluminum is a more active metal than iron. As a result of its greater activ-ity, aluminum is oxidized more quickly than iron. If this is true, why doesan aluminum can degrade much more slowly than a tin can, which ismade of iron-containing steel that is coated with a thin layer of tin? Thereason is that, like copper, aluminum is oxidized to form a compoundthat coats the metal and protects it from further corrosion, as shown inFigure 16.12. Aluminum reacts with oxygen to form aluminum oxide in aredox reaction.

4Al(s) � 3O2(g) ˇ 2Al2O3(s)

A coating of aluminum oxide is tough and does not flake off easily, as ironoxide rust does. When rust flakes fall off a surface, additional metal isexposed to air and becomes corroded.

Figure 16.12Corrosion of Ironand AluminumBecause iron rust isporous and flaky, itdoes not form agood protectivecoating for itself. �

Figure 16.11The Green LadyThe green color of theStatue of Liberty in NewYork Harbor is due to alayer of patina, or pro-tective coating, that cov-ers the copper sheetsmaking up the statue.The presence of the pati-na helps keep the statuefrom corroding furtherbecause oxygen cannotget through the patinato reach the copper lay-ers underneath.

Steel can

Iron metal

Corrosive solution

Aluminum canTin-coated steel can

Iron metal

Tin

Corrosive solution

Aluminum metal

Aluminum oxide

Corrosive solution

A tin coating offers some protectionto the iron. However, if a hole orcrack develops in the thin tin coat-ing, the underlying iron corrodesrapidly. A tin-coated steel can willdegrade completely in about 100years. The aluminum oxide coatingon an aluminum can is tough andclosely packed. It protects theunderlying aluminum from furthercorrosion so that the can will takeabout 400 years to degrade. �


Lightning-Produced FertilizerDid you know that one of the main nutrients

plants need is nitrogen? Although the air sur-rounding Earth is almost 80 percent nitrogen, thenitrogen is in the form of N2 molecules, a form thatmost plants and animals cannot use. Nitrogen fromthe air is converted to a form thatplants can use by a process callednitrogen fixation. Plants can best usenitrogen when it is in the form of theammonium ion, NH4

�, where thenitrogen has an oxidation number of3�, but they can also use the nitrateion, NO3

�, with nitrogen having anoxidation number of 5�.

Nitrogen fixation Nitrogen can befixed for plants in three ways: by light-ning, by nitrogen-fixing bacteria livingin the roots of plants or in the soil,and by commercial synthesis reactionssuch as the Haber ammonia process.

Nitrogen is a fairly inert gasbecause the triple bond of N2 isstrong and resists breaking. However,the exceptionally high energy and temperatures oflightning can easily break bonds and allow forrecombination of gases in the atmosphere.

Lightning-driven reactions In the process oflightning-driven nitrogen fixation, nitrogen andoxygen combine to form nitrogen monoxide.Nitrogen monoxide then combines with more oxy-gen to form nitrogen dioxide. This nitrogen diox-ide mixes with water in the air to form nitric acidand more nitrogen monoxide, which is available tocontinue the cycle.

N2 � O2 ˇ 2NO2NO � O2 ˇ 2NO2

3NO2 � H2O ˇ 2HNO3 � NO

Fertilizer production The pH of rainwater is nat-urally slightly acidic, and some of this acidity is

due to the dissolved nitric acid, HNO3, from nitro-gen fixation. As the rain soaks into the soil, bacte-ria convert the nitrate ions into ammonium ions.

How does nature’s manufacturing of fixednitrogen compare with commercial production offixed nitrogen? Lightning may seem uncommon,but it is estimated that there are approximately

10 000 lightning stormsevery day over the sur-face of Earth. Statedanother way, lightningstrikes 100 times a sec-ond on the planet as awhole. Approximately10 billion kg of nitrogenare fixed yearly in theatmosphere. Biologicalagents such as bacteriafix about 100 billion kgof nitrogen yearly, andan amount equal to thatis fixed through themanufacture of fertilizerand other industrialprocesses.

Chemistry

Exploring Further

1. Classifying Nitrogen fixation in the soil isaccomplished by bacteria living in the roots ofcertain plants. Name some of these plants.

2. Applying In each of the three equationsshown, what is oxidized and what is reduced?

3. Acquiring Information The process by whichnitrogen is put back into the air is called deni-trification. Find out what conditions are neces-sary for this process and what reaction occurs.

To learn more about the nitrogen cycle, visit theChemistry Web site at chemistryca.com

http://chemistryca.com


Silver Tarnish: A Redox ReactionImagine if, along with your usual chores of taking out the trash, wash-

ing dishes, feeding your pets, and taking care of your younger siblings,you also had to polish the silver—as people did back in your great-grandparents’ days. How would you find time for any fun? Fortunately,other materials such as stainless steel have replaced most “silverware.”Why do silver utensils have to be polished, but those made of stainlesssteel or aluminum don’t? Silver becomes tarnished through a redox reac-tion that is a form of corrosion, as rusting is. Tarnish is formed on thesurface of a silver object when silver reacts with H2S in air. The product,black silver sulfide, forms the coating of tarnish on the silver.

O2(g) � 4Ag(s) � 2H2S(g) ˇ 2Ag2S(s) � 2H2O(l)

Many commercial silver polishes contain abrasives that help to removetarnish. Unfortunately, they also remove some of the silver. A more gentleway to remove tarnish from the surface of a silver object involves anotherredox reaction. In this reaction, aluminum foil scraps act as a reducing agent.

Figure 16.13Removing Silver TarnishEven though corrosion is an unwanted redox reaction, removing thetarnish makes use of another redox reaction. A nest of crumpledaluminum foil scraps is made at the bottom of a large pot. The tarnishedsilver object is added, making sure the silver is in contact with the foilscraps. Baking soda is added, and the silver is covered with water. Whenthe pot is heated on a stove, the silver sulfide tarnish is reduced to silveratoms, and the silver object becomes shiny and bright.

2– 3+2

�2

+6Ag (s)2Al (s) 2Al (aq)3S (s) 6OH (aq)6H O(l) 6Ag (s)� � � ��

6e–

3H S(g)

ˇ

ˇ0 0

This reaction is essentially the reverse of the reaction that forms tarnish.Here, silver ions in the Ag2S tarnish are reduced to silver atoms, while alu-minum atoms in the foil are oxidized to aluminum ions. The tarnish-removing solution usually includes baking soda (sodium hydrogen car-bonate) to help remove any aluminum oxide coating that forms and tomake the cleaning solution more conductive. Figure 16.13 shows how thismethod of silver cleaning is done.

tarnish:terne (OF) dull,wan

The shiny surfacesof many metalobjects lose lusterand become dulland tarnished asthe metal atomsundergo oxidation.

Forensic Blood DetectionThe gas station at the corner was robbed, and the

cashier was shot. On television, police announce thatSuspect A has been taken into custody. They haveconfiscated a jacket, allegedly worn by the suspect.After preliminary examination by the police depart-ment, the jacket is sent to a forensic laboratory forscientific investigation. One of the first tests a techni-cian at the laboratory will carry out determineswhether or not there are blood stains on the jacket.

The Luminol TestThe technician may choose from several chem-

ical tests for blood, all based on the fact that thehemoglobin in blood catalyzes the oxidation of anumber of organic indicators to produce a col-ored product that emits light, or luminesces.

The technician on this case chooses the lumi-nol test. Luminol has an organic double-ringstructure, shown below. In 1928, Germanchemists first observed the blue-green lumines-cence when the compound was oxidized in alka-line solution. It was soon found that a number ofoxidizing agents, such as hydrogen peroxide,bring about the luminescence. Later, workersnoted that the luminescence was greatlyenhanced by the presence of blood, which led toits current use in forensic investigations.

The technician carefully mixes an alkalinesolution of luminol with aqueous sodium perox-ide and, in a darkened workplace, sprays the solu-tion onto suspected spots on the jacket. Bingo!An intense, blue-green chemiluminescence isemitted from several spots. Because the glow will

last for a few minutes, the technician photo-graphs the spots and their telltale light.

Ruling Out with LuminolYou may wonder if this relatively simple proce-

dure will serve to convict Suspect A. Certainlynot. However, if the test had been negative, Sus-pect A might have been cleared from suspicion. Anegative result ensures that a stain is not blood.But, because this is not the case with the stainson the jacket, the luminol test is preliminary andwill be used with other tests.

The luminol test is especially useful because itworks well with both fresh and dried blood.Luminol has one particularly useful feature. Thesame stains can bemade luminescentover and over again ifthe spray is allowedto dry and the stainsare resprayed.

A positive testshould not be takenas absolute proof ofblood because lumi-nol reacts with copper and cobalt ions, as well aswith the iron in hemoglobin. However, it reactsmuch more strongly with hemoglobin. A largenumber of forensic authorities believe that theluminol test has value as a preliminary sortingtechnique.ONH2

NH

O

NHLuminol

1. Applying If a lumi-nol test yields a pos-itive reaction, whatis the next logicalstep?

2. HypothesizingWhy can it almost

never be assumedthat stains areuncontaminated,although stain evi-dence is importantin a criminal inves-tigation?

DISCUSSING THE TECHNOLOGY


&TECHNOLOGYC H E M I S T R Y


Chemiluminescence: It’s CoolSome redox reactions can release light energy at room temperature.

The production of this kind of cool light by a chemical reaction is calledchemiluminescence. The light from chemiluminescent reactions can beused in emergency light sticks that work without an external energysource. You may recall learning in Chapter 6 how these light sticks work.Now you know that the reaction that takes place when the two solutionsin the light sticks are mixed involves an oxidation and a reduction.

Some chemiluminescent redox reactions occur naturally in the atmo-sphere as a result of lightning, Figure 16.14. Other chemiluminescentreactions involve luminol, an organic compound that emits cool lightwhen it is oxidized. Luminol reactions are utilized by forensic chemists toanalyze evidence in crime investigations. They spray luminol onto a loca-tion where the presence of blood is suspected. If blood is present, theiron(II) ions in the blood oxidize the luminol to form a chemilumines-cent compound that glows in the dark. The iron is reduced by the lumi-nol. Figure 16.14 shows the glow from the oxidized form of luminol.

Biochemical Redox ProcessesHow are bears able to stay warm enough to keep from freezing during

their winter hibernation? How do marathon runners get the energy to fin-ish a race without stopping to eat? In both cases, fats stored in the bodyare oxidized. Oxygen molecules from the air are reduced as they gain elec-trons to form water. In a series of redox reactions called respiration,

When luminol is oxidizedand is observed in the dark,an eerie blue-green glow isproduced through chemi-luminescence. �

Figure 16.14Chemiluminescence

� When lightning is produced by an electrical dis-charge in the atmosphere, electrons in moleculesof O2 and N2 gases are excited to higher energylevels. Energy from the electricity breaks the mole-cules into atoms. When the atoms recombine toform molecules and the electrons return to lowerenergy levels, light energy is released throughchemiluminescence.

Nitrous oxide is pro-duced by a redox reac-tion between oxygenand nitrogen duringlightning storms. It wasdiscovered and studiedin the late 1700s byJoseph Priestley, whofound that inhaling itresulted in unusual sideeffects including laugh-ing, singing, and fight-ing. For this reason, itwas called laughinggas. Its anestheticproperties were discov-ered by accident inConnecticut in 1844 ata public demonstrationgiven for amusementwhen a man whoinhaled nitrous oxidecut his leg badly in ascuffle but felt no painuntil the gas wore off.


energy is released. Figure 16.15 shows one effect of this heat in plants.Respiration will be discussed in Chapter 19. Many other redox reactionstake place in living things. Electrons are transferred between molecules inredox reactions during photosynthesis and in the reactions that firefliesuse to flash light signals to potential mates. You will study photosynthesisin Chapter 20.

Some organisms can use the energy released during redox reactions toconvert chemical energy into light energy, a process called biolumines-cence. You are probably familiar with the flashing lights given off by fire-flies during courtship, but did you know that many different organisms—including some fish, at least one type of mushroom, and a caterpillarknown as a glowworm—also are bioluminescent? Figure 16.16 shows bio-luminescence in fireflies.

Now that you have learned what redox reactions are and have readabout some of the processes of which they are part, you can reexaminethe redox reaction that makes cut fruit turn brown. The color is due tobrown pigments that are formed by the oxidation of colorless compoundsnormally present in the cells of the fruit. Oxygen in the air is the oxidizingagent that reacts with the colorless compounds to produce the brown pig-ments. The oxygen is reduced when it accepts electrons from the pig-ments, so the pigments function as reducing agents. This combination ofoxidation and reductiongoes hand in hand in aredox reaction becauseelectrons that are lost byone element must begained by another.

Figure 16.15Keeping WarmAlthough it is common to think that only mammals keepwarm, in truth, all plants and animals maintain a tempera-ture at which their enzymes function best. Plants keep fromfreezing because heat is produced as a by-product of respira-tion and photosynthesis. One of the first plants to pokethrough the snow in early spring is the heat-producing skunkcabbage. The heat it releases allows it to get a head start onother plants and also contributes to the unpleasant odorthat gives it its name.

Figure 16.16Firefly SignalsFireflies use flashes of light to attract mates. Lightenergy is released during an enzyme-catalyzedredox reaction. Luciferase is the name given to theenzyme that speeds up the reaction in which theorganic molecule luciferin is oxidized.

bioluminescent:bios (GK) lifelumen (L) lightescentis (L) begin-ning to be, have,or do

A bioluminescentsubstance under-goes a chemicalreaction in livingthings in whichpotential energy inchemical bonds isconverted intolight energy.


The skin of a fruit keeps oxygen out, which is why unbroken fruit doesnot turn brown. Coating cut fruit with an antioxidant can prevent brown-ing and keep a fruit salad looking fresh longer. The vitamin C in lemons is agood antioxidant. If lemon juice is squirted onto cut banana or apple slices,they will not brown as quickly because the vitamin C reacts with oxygenmore readily than do the fruit-browning compounds, Figure 16.17.

Connecting IdeasMost reactions involve electron transfer and thus are redox reactions. You

have learned to identify which element is reduced and which is oxidizedwhen you are given the equation for a redox reaction. You might wonderwhy one element accepts electrons from another and whether you can predict which element will be oxidized and which will be reduced. Learningto make those predictions is the next step in your study of electron-transferprocesses in compounds and will help you understand how redox reactionsin batteries produce electricity.

Understanding Concepts1. What role does the reducing agent hydro-

quinone play in the production of a photo-graphic negative?

2. How is most of the iron that is used for makingsteel purified from iron ores?

3. Why do aluminum cans degrade more slowlythan cans made of iron?

Thinking Critically 4. Applying Concepts Oxygen is required for the

production of light by fireflies. What role doesthe oxygen play in the reaction?

Applying Chemistry5. Bleaching Why can’t rust stains be removed

with bleach?

SECTION REVIEW

Figure 16.17AntioxidantsVitamin C owes its antioxidant propertiesto the fact that it reacts so readily withoxygen. When added to a food product,oxygen reacts preferentially with vita-min C, thereby sparing the foodproduct from oxidation. Otheranti-oxidant food additivesinclude the synthetic com-pounds BHA and BHT andthe natural antioxidant,vitamin E.

Myoglobin, found inmuscle tissue, is aniron-containing proteinthat stores oxygen.Myoglobin in livingmuscle tissue is boundto oxygen and is a redcolor. It becomes palepurple after deathwhen the oxygen islost. Heating meatresults in oxidation ofthe iron in myoglobin,which then has thebrown color that tellsyou the meat is cooked.

For more practice with solvingproblems, see SupplementalPractice Problems,Appendix B.

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Chapter 16 Assessment 577

16.1 The Nature of Oxidation-Reduction Reactions

■ Oxidation occurs when an atom or ion losesone or more electrons and attains a more pos-itive oxidation number. Reduction takes placewhen an atom or ion gains electrons andattains a more negative oxidation number.

■ Oxidation and reduction reactions alwaysoccur together in a net process called a redoxreaction.

■ An oxidizing agent is the substance that gainselectrons and is reduced during a redox reac-tion. A reducing agent is the substance thatloses electrons and is oxidized during a redoxreaction.

16.2 Applications of Oxidation-Reduction Reactions■ In photography, light triggers the reduction of

silver ions to silver metal on photographic film.

■ Bleach removes stains from clothing by oxi-dizing colored molecules to form colorlessmolecules.

■ Metals such as copper and aluminum areresistant to corrosion even though they are

easily oxidized because the products of theirreactions with oxygen form protective coat-ings on the surface of the metal.

■ Chemiluminescent reactions in emergencylight sticks, lightning, and the luminol reac-tion convert the energy of chemical bondsinto light energy.

■ Some organisms use redox reactions to pro-duce light, which they use in communication.This light production is called biolumines-cence.

■ Cut fruits turn brown because compounds inthe fruit cells react with oxygen in a redoxreaction to produce brown pigments. Coatingthe fruits with antioxidants can prevent thisbrowning.

VocabularyFor each of the following terms, write a sentence that showsyour understanding of its meaning.

oxidation oxidizing agentoxidation-reduction reducing agent

reaction reduction

REVIEWING MAIN IDEAS

UNDERSTANDING CONCEPTS1. What is the difference between an oxidizing

agent and a reducing agent?

2. Which of the changes indicated are oxidationsand which are reductions?

a) Cu becomes Cu2�

b) Sn4� becomes Sn2�

c) Cr3� becomes Cr6�

d) Ag becomes Ag�

3. Identify the oxidizing agent in each of the fol-lowing reactions.

a) Cu2�(aq) � Mg(s) ˇ Cu(s) � Mg2+(aq)b) Fe2O3(s) � 3CO(g) ˇ 2Fe(l) � 3CO2(g)

4. What is the oxidizing agent in householdbleach?

5. Why does a photographic negative need to befixed?

6. In which direction do electrons move during aredox reaction: from oxidizing agent to reduc-ing agent or vice versa?

7. Why is aluminum metal used to remove tar-nish from silver?

8. What chemical process do hibernating animalsuse to stay warm?

9. Write the equation for the redox reaction thatoccurs when a piece of iron metal is dipped ina solution of copper(II) sulfate.

CHAPTER 16 ASSESSMENT

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10. Identify the following as an oxidation reactionor a reduction reaction.

Fe2� ˇ Fe3� � e�

APPLYING CONCEPTS11. If galvanized nails, which have been coated

with zinc, are placed in a brown solution con-taining I2, the solution slowly turns colorless.Adding a few drops of bleach to the colorlesssolution results in a return of the brown color.Explain what makes these changes occur.

12. List several ways in which a steel chain-linkfence could be treated to prevent corrosion.

13. When hydrogen peroxide is added to a color-less solution of potassium iodide, a red-browncolor appears. What substance is responsiblefor the color?

14. Write the equation for a reaction that is not aredox reaction. Are electrons transferred in thisreaction?

15. Indigo is one of the oldest known dyes. It hasbeen detected in cloth used to wrap mummiesthat are more than 5000 years old. When cottonjeans are dyed with indigo, they are dipped intoa yellow solution of indigo and sodium hydro-sulfite, which is a good reducing agent. Withinminutes after being taken out of the solution,the jeans turn blue. How can you explain this?

16. A shiny copper mirror can be formed on theinside of a test tube in which the followingreaction takes place.

H2C=O(aq) � Cu2�(aq) � 2OH�(aq) ˇformaldehyde

Cu(s) � HCOOH(aq) � H2O(l)formic acid

a) Identify the substance that is reduced dur-ing this reaction.

b) Identify the substance that is oxidized dur-ing this reaction.

c) What is the oxidizing agent in this reaction?d) What is the reducing agent in this reaction?

17. Is oxygen a necessary reactant for an oxidationreaction? Explain.

18. Sodium nitrite is often added to meat to inhib-it the growth of microorganisms and to keepthe meat from spoiling. Under the acidic con-ditions in our stomachs, nitrites can be con-verted into potentially cancer-causing sub-stances. Vitamin C can convert nitrite ions intonitrogen monoxide gas and may help protectus from the effects of these ions.

NO2�(aq) ˇ NO(g)

a) Is the nitrite ion oxidized or reduced in thisreaction?

b) Does vitamin C act as an oxidizing agent ora reducing agent?

19. Why is gold rather than copper used to coatelectrical connections in expensive electronicequipment?

Everyday Chemistry20. Just before World War I, a German chemist

named Fritz Haber developed a process for fix-ing atmospheric nitrogen into ammonia. Theammonia produced this way can be convertedinto ammonium nitrate, an important fertiliz-er and explosive.

3H2 � N2 ˇ 2NH3

a) What element is oxidized during this reac-tion? What is reduced?

b) What is the oxidizing agent? What is thereducing agent?

Physics Connection21. By passing an electric current through water,

the water can be separated into its componentelements in the reverse of the reaction used topower the main stage of the space shuttle.

2H2O(l) � energy ˇ 2H2(g) � O2(g)

a) Is this a redox reaction? If so, what elementis oxidized?

Fe3� � e�Fe2�


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Chapter 16 Assessment 579

b) Where does the energy for this endothermicreaction come from?

How It Works22. If ethanol were less volatile, how might the

usefulness of a Breathalyzer test be affected?Explain.

Chemistry and Technology23. Why should a positive result from the luminol

test not be taken as proof of the presence ofblood?

THINKING CRITICALLYUsing a Table24. The table below lists some of the most com-

mon compounds that are used as oxidizingagents.

a) Name each of the compounds in the table.b) List at least one practical application, men-

tioned in this chapter or from a referencebook, of each of these oxidizing agents.

c) Make a similar table for common reducingagents. Should any compounds be listed inboth tables?

Making Predictions25. Hydrogen peroxide (H2O2) can be used to

restore white areas of paintings that have dark-ened from the reaction of lead paint pigmentswith polluted air containing hydrogen sulfidegas.

PbS � 4H2O2 ˇ PbSO4 � 4H2O(black) (white)

Could hydrogen peroxide be used to removetarnish from silver objects? Would the reactionhave any undesirable effects?

Interpreting Data26. ChemLab Write the balanced equation for the

reaction that caused the limewater to becomecloudy. Is this a redox reaction? Explain.

27. MiniLab 1 Why do you think corrosion seemsto occur mostly at the head and point of a nail?

Making Inferences28. MiniLab 2 When a pile of orange ammonium

dichromate is ignited, it decomposes in anexothermic reaction in which the green prod-uct and flames shoot upward like an eruptingvolcano. (CAUTION: Do NOT perform thisreaction.)

(NH4)2Cr2O7(s) ˇCr2O3(s) � N2(g) � 4H2O(g)

a) What is the reducing agent in this reaction?The oxidizing agent?

b) How is this reaction similar to the Breath-alyzer reaction?

CUMULATIVE REVIEW29. Identify each of the following as a pure sub-

stance or a mixture. (Chapter 1)

a) petroleum d) diamondb) fruit juice e) milkc) smog f) iron ore

30. List some characteristic properties of metals.(Chapter 3)

31. Name each of the following ionic compounds.(Chapter 5)

a) NaF d) Na2Cr2O7

b) CaS e) KCNc) Al(OH)3 f) NH4Cl

32. How many grams of nitrogen are needed toreact completely with 346 g of hydrogen toform ammonia by the Haber process? (Chapter 12)

N2 � 3H2 ˇ 2NH3


O2 K2Cr2O7

H2O2 HNO3

KMnO4 NaClO

Cl2 KClO3

Common Oxidizing Agents


33. Draw Lewis electron dot diagrams for each ofthe following covalent molecules. (Chapter 9)

a) CHCl3 c) CH3CH3

b) CH3CH2OH

SKILL REVIEW34. Designing an Experiment Do you think silver

will tarnish more quickly in clean air or in pol-luted air? Design an experiment to test yourhypothesis.

WRITING IN CHEMISTRY35. Research the evidence that suggests that the

antioxidant properties of vitamin C may helpprevent cancer in people who take large dosesof this vitamin. Write a summary of your find-ings in which you propose how you would domore tests to determine whether or not vita-min C has anticarcinogenic properties.

PROBLEM SOLVING36. A flask filled with acid-washed steel wool is fit-

ted with a long, thin glass tube in a rubberstopper. When the flask is inverted so the tubeopening is in a beaker of colored water, thewater slowly begins to rise in the tube. Write asummary of this experiment, as if you had per-formed it. Explain what makes the water rise.Predict what portion of the flask will be filledwith water at the end of the experiment.

37. The patina coating on the Statue of Liberty haspreserved most of the copper metal in the stat-ue. Some damage does occur wherever steelrivets are in contact with copper and exposedto water. Do library research to determine whythose sites are more susceptible to corrosionthan the rest of the statue. Write up your find-ings in a short report. Include a diagram ormake a poster showing the movement of elec-trons in the process.

38. Metallic lithium reacts vigorously with fluorinegas to form lithium fluoride.

a) Write an equation for this process.

b) Is this an oxidation-reduction reaction?

c) If it is an oxidation-reduction, which ele-ment is oxidized? Which is reduced?

d) If 2.0 g of lithium are reacted with 0.1 L flu-orine at STP, which reactant is limiting?

e) If 0.04 g of lithium fluoride is formed inreaction in part d., what is the percent yield?

39. Identify the oxidizing reagent in each of thefollowing reactions.

a) C2H5OH(l) � 3O2(g) → 2CO2(g) � 3H2O(l)

b) CuO(s) � H2(g) → Cu(s) � H2O(l)

c) 2FeO(s) � C(s) → 2Fe(s) � CO2(g)

d) 2Fe2+(aq) � Br2(l) → 2Fe3+(aq) � 2 Br�(aq)

40. When coal and other fossil fuels containingsulfur are burned, sulfur is converted to sulfurdioxide: S(s) � O2(g) → SO2(g)

a) Is this an oxidation-reduction reaction?

b) If it is an oxidation-reduction, which ele-ment is oxidized? Which is reduced?

c) If 7.0 � 103 kg of fuels containing 3.5 per-cent sulfur are burned in a city on a givenday, how much SO2 will be emitted? Assumethat the sulfur reacts completely.

41. Sodium nitrite is formed when sodium nitratereacts with lead:

NaNO3(s) � Pb(s) → NaNO2(s) � PbO(s)

a) What is the oxidizing reagent in this reac-tion? What is the reducing reagent?

b) If 5.00 g of sodium nitrate is reacted withan excess of lead, what mass of sodiumnitrite will form if the yield is 100 percent?


Standardized Test Practice

1. The term used to describe a chemical reactionin which a substance losses electrons to anoth-er substance is

a) oxidation. c) reduction.b) redox. d) corrosion.

2. The oxidation numbers of the elements inCuSO4 are

a) Cu � �2, S � �6, O � �2b) Cu � �3, S � �5, O � �2c) Cu � �2, S � �2, O � �1d) Cu � �2, S � 0, O � �2

3. For the reaction X + Y → XY, the elementthat will be reduced is the one that is

a) more reactive. c) more electronegative.b) more massive. d) more radioactive.

4. The reaction between sodium iodine andchlorine is:

2NaI(aq) � Cl2(aq) → 2NaCl(aq) � I2(aq)

The oxidation state of Na remains unchangedbecause

a) Na� is a spectator ion.b) Na� cannot be reduced.c) Na� is an uncombined element.d) Na� is a monatomic ion.

5. The reaction between nickel and copper(II)chloride is:

Ni(s) � CuCl2(aq) → Cu(s) + NiCl2(aq)

The half reactions for this redox reaction are

a) Ni → Ni 2� � 2e�; Cl2 → 2Cl� � 2e�

b) Ni → Ni � � e�; Cu+ + e� → Cuc) Ni → Ni 2� � 2e�; Cu2� � 2e� → Cud) Ni → Ni 2� � 2e�; 2C� � 2e� → Cu

Interpreting Tables: Use the table in the next col-umn to answer questions 6–8.

6. Which of the following elements forms amonatomic ion that is a spectator in the redoxreaction?

a) Zn c) Nb) O d) H

7. The oxidation number of N in Zn(NO3)2 is

a) �3. a) �1.b) �5. d) �6.

8. The element that is oxidized in this reaction is

a) Zn. c) N.b) O. d) H.

9. Why are redox reactions so common in every-day situations?

a) Nitrogen is an abundant reducing agent.b) Nitrogen is an abundant oxidizing agent.c) Oxygen is an abundant reducing agent.d) Oxygen is an abundant oxidizing agent.

Standardized Test Practice 581

Write It Down! Most tests ask you a large num-ber of questions in a small amount of time. Writedown your work whenever possible. Write out thehalf-reactions for a redox problem, and make surethey add up. Do math on paper, not in your head.Underline and reread important facts in passagesand diagrams—don’t try to memorize them.

Test Taking Tip

chemistryca.com/standardized_test

Data for Elements in the Redox Reaction Zn + HNO3 0 Zn(NO3)2 + NO2 + H2O

Element Oxidation Complex ion of whichNumber element is a part

Zn 0 none

Zn in Zn(NO3)2 �2 none

H in HNO3 �1 none

H in H2O ? none

N in HNO3 ? NO3�

N in NO2 �4 none

N in Zn(NO3)2 ? NO3�

O in HNO3 �2 NO3�

O in NO2 ? none

O in Zn(NO3)2 ? NO3�

O in H2O �2 none

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c 16hapter oxidation-reduction reactions

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