Unit 7 Electrochemistry - · PDF fileElectrochemical reactions may be the most important ... enzymes for biosensor applications, ... Unit 7 Electrochemistry ARE YOU READY?

Unit 7 Electrochemistry

77unit

ElectrochemistryElectrochemistry

You have previously classified chemical reactions usingempirical generalizations such as single replacement. Thisclassification of reactions is useful because it helps you tomake predictions that generally can be verified in a labo-ratory. However, these predictions do not require any the-oretical knowledge about how the reaction actually occurs.Chemists classify most common reactions based on a the-oretical understanding of the reaction process. One veryimportant example is the electron–transfer reaction, knownas an electrochemical reaction. The study of chemical reac-tions associated with the transfer of electrons is known aselectrochemistry.

Electrochemical reactions may be the most importantreactions on Earth. Living things use electrochemical reactions for photosynthesis andmetabolism. Nonliving things undergo electrochemical reactions such as corrosion, (asshown by the metal door in the photograph), metallurgy, and combustion.

Electrochemistry can be applied to solve various technological problems. For example,the research conducted by Dr. Birss (Figure 1) and her research team at the Universityof Calgary is rooted in electrochemistry. According to Dr. Birss, “One of the main goalsof our research over the last 15 years has been to produce useful thin film materials thatcan serve as electrocatalysts in fuel cells, to protect metals from corrosion, as a matrix forenzymes for biosensor applications, and to create useful nanoarchitectures.”

Research into fuel cells involves developing new catalysts of nanometre-sized metal ormetal oxide particles, researching new low-cost cathode materials for the reduction ofoxygen in fuel cells, and investigating factors that affect the performance of these cells.Dr. Birss’ team is also researching the development of novel electrochemical methods ofcoating metals with protective oxide films, which is of particular interest to the aero-space and oil and gas industries, and the development of glucose biosensors for peoplewith diabetes.

As you progress through the unit, think about these focusing questions:

• What is an electrochemical change?

• How have scientific knowledge and technological innovation been integratedin the field of electrochemistry?

552 Unit 7 NEL

Figure 1Dr. Viola Birss

Unit 7 - Ch 13 CHEM30 11/1/06 12:55 PM Page 552

Unit 7

Electrochemistry 553NEL

GENERAL OUTCOMESIn this unit, you will• explain the nature of oxidation–reduction

reactions

• apply the principles of oxidation–reduction to electrochemical cells


Unit 7Electrochemistry

ARE YOU READY?

554 Unit 7 NEL

These questions will help you find out what you already know, and what you need toreview, before you continue with this unit.

Knowledge1. When a metal atom forms an ion, the atom _____ electrons to form a ________

charged ion.

2. When a nonmetal atom forms an ion, the atom _____ electrons to form a________ charged ion.

3. According to trends in the periodic table and your general chemistry knowledge,copy and complete Table 1.

4. Complete the following chemical equations using Lewis symbols or formulas forthe products.

(a)potassium � chlorine →

(b)phosphorus � chlorine →

(c) Compare electron rearrangement in (a) to electron rearrangement in (b).(d) According to the bonding theory you have studied, what is believed to

determine whether two atoms transfer or share electrons?

5. (a) Write the generalized chemical equation for a single replacement reaction.(b) How do you know what class of element—metal or nonmetal—forms in a

single replacement reaction?

6. Complete and balance the chemical equation for each of the following reactions:(a) __ Zn(s) � __ AgNO3(aq) →(b) __ Cl2(aq) � __ KBr(aq) →(c) __ Al(s) � __ HCl(aq) →(d) __ C3H8(g) � __ O2(g) → __ CO2(g) � __

STS Connections7. When lead ion solutions are used in a laboratory, disposal of reactant solutions

and products requires special care. The recommended method for lead disposalis to precipitate all lead(II) ions in the form of the insoluble lead(II) silicate.This solid can then be disposed of as regular solid waste as long as it is buriedand not incinerated.(a) Why should lead compounds not be incinerated or disposed in a soluble

form down the drain?

Concepts

• atoms and ions

• ionic and covalent bonding

• balanced chemical reactionequations

• dissociation and ionization

• mole concept andstoichiometry

Skills

• safe handling and disposalbased on WHMIS

• diagnostic tests

• solution preparation

• reaction equipment andprocedures

• chemical amount andstoichiometry calculations

• designing and evaluatingexperiments

• SI and IUPACcommunication conventions

You can review prerequisiteconcepts and skills in theChemistry Review unit, onthe Nelson Web site, and inthe Appendices.

A Unit Pre-Test is alsoavailable online.

Prerequisites

Table 1 Reactivity of Elements

Category Groups or examples

most reactive metals

least reactive metals

most reactive nonmetals

least reactive nonmetals

www.science.nelson.com GO

ClK + →

Cl+P →


Unit 7


(b) Briefly describe the issue related to lead in the environment. What was themost common source of lead pollution until legislation forced a change inthe substance sold?

(c) Suggest a reason why it is acceptable to bury a lead compound that is notsoluble in water.

(d) When excess sodium silicate solution was added to 150 mL of lead(II)nitrate solution, 2.41 g of dried precipitate was obtained. Calculate theamount concentration of the lead(II) nitrate solution.

8. Compare the nature of science and the nature of technology, noting the keydifferences in the characteristics of these two forms of human endeavours.

Skills9. What should you check before plugging in electrical equipment?

10. List some safety precautions for operating electrical equipment.

11. When studying chemical reactions, diagnostic tests are important to identifyproducts (see Appendix C.4). Interpret the evidence in the photos below toidentify the product or type of product formed in a chemical reaction.

12. List the general steps of a procedure to prepare a standard solution starting witha pure solid.

13. Titration is a common technological process used in different types of chemicalanalysis. List the main laboratory equipment needed to conduct a titration.

halogen test (a) (b)

acidity test (using (c) (d)bromothymol blue)

ion colour test (e) (f)


1313 Redox Reactions

chapter

Redox Reactions

556 Chapter 13 NEL

In this chapter

Career Connections: Materials/Metallurgical Engineer; Conservator

Exploration: CleaningSilver

Case Study: EarlyMetallurgy

Investigation 13.1: SingleReplacement Reactions

Web Activity: HenryTaube and RudolphMarcus

Web Activity: RedoxReaction

Biology Connection:Metabolic Redox

Investigation 13.2:Spontaneity of RedoxReactions

Web Activity: Piercings: A Rash Decision

Lab Exercise 13.A:Building a Redox Table

Investigation 13.3:Predicting the Reaction ofSodium Metal(Demonstration)

Lab Exercise 13.B:Oxidation States ofVanadium

Web Activity: CatalyticConverters

Case Study: BleachingWood Pulp

Lab Exercise 13.C:Analyzing for Tin

Web Activity: ImantsLauks

Lab Exercise 13.D:Analyzing for Chromiumin Steel

Investigation 13.4:Analyzing a HydrogenPeroxide Solution

Answer these questions as best you can with your current knowledge. Then, usingthe concepts and skills you have learned, you will revise your answers at the end ofthe chapter.

1. What are the key theoretical concepts that distinguish electrochemical reactions fromother kinds of chemical reactions?

2. How can you predict whether or not a mixture of chemicals will react?

3. What are the similarities and differences between redox stoichiometry and otherstoichiometry that you have learned?

STARTING Points

Of all chemical changes, electrochemical (electron transfer) reactions are the mostcommon in both living and nonliving systems. Photosynthesis, cellular respiration, andmetabolism are all electrochemical processes. Technologies involving electrochemistry,such as combustion and the production of metals from their ores, have been used for thousands of years. In the development of these applications, the technology was suc-cessfully developed long before there was any scientific understanding of the processes.In other words, technology led science. More recently, however, a sophisticated under-standing of modern electrochemistry has led to the invention of modern technologiessuch as aureate (gold-like) plating of coins (Figure 1), and the development of fuel cells and biosensors. Today, science and technology nurture each other in a symbiotic relationship.

Knowledge of electrochemistry will help you connect and clarify many seeminglyunrelated reactions, and understand the interactions of science and technology. Forexample, paper for this book is produced from trees that used photosynthesis reactionsto grow. Harvesting trees requires machinery made from steel, which is produced bythe electrochemical reduction of iron ore. The energy used to run the machines in apulp mill comes from the combustion of fossil fuels. Electrochemical reactions play a rolein the production of paper from wood pulp. Photographs used in books may involvethe reduction of silver ions to silver metal to form a negative image, which is printedusing metal plates made by electrochemical reactions. All of the people involved, fromthose who harvest the trees to those who read the book, metabolize food to live andwork. An understanding of electrochemistry will give you a broader comprehension ofmany chemical reactions and their importance in both living and nonliving systems.


Redox Reactions 557NEL

Exploration Cleaning Silver

Have you ever noticed how silver becomes darkened and cloudyover time? This dark tarnish is mainly silver sulfide. The sulfurmay come from several sources: air, which contains smallamounts of hydrogen sulfide from industrial operations or fromthe decomposition of organic matter; natural gas, whichcontains a sulfur compound to give the odourless methane anoticeable odour; and some food items, including egg whites,mustard, and mayonnaise.

There are several technological solutions to restore thesilvery appearance of tarnished silverware. Many silver-cleaningmixtures contain abrasives that are used to scour off the tarnish.However, these cleaners can remove the silver as well. A bettertechnological process is to convert the silver sulfide back tosilver metal. A common household remedy based on thisapproach is to react the tarnish with aluminium in a hot solutionof baking soda and table salt.

Materials: tarnished silverware such as a small spoon, 400 or600 mL beaker, piece of aluminium foil about 5 cm by 5 cm, hotplate, stirring rod, measuring spoon, tongs, baking soda, tablesalt, tap water

Do not touch the surface of the hot plate. Switchoff the hot plate immediately after use.

• Fill the beaker about three-quarters full with tap water.• Add about one teaspoon each of baking soda and table salt

to the water. Stir to dissolve.• Heat the solution on the hot plate to near boiling and then

turn off the hot plate.(a) Describe the appearance of the tarnished silver object and

the aluminium foil.• Use the tarnished silver object to push the aluminium foil to

the bottom of the beaker. Make sure the silver objectremains in contact with the foil.

• Observe any changes immediately and after severalminutes.

• Use the tongs to remove the silver object and the aluminiumfoil.

(b) Describe the final appearance of the silver object and thealuminium foil.

(c) Write a balanced chemical equation for the singlereplacement reaction of solid silver sulfide and aluminium.How well does this equation explain your observations?

(d) Evaluate this technological process for cleaning tarnishedsilver.

• When cool, dispose of the solution down the sink and putthe aluminium foil into the regular garbage.

Figure 1The aureate plating on the $1 and $2 coins gives themtheir golden colour. This plating process was invented inFort Saskatchewan, Alberta, at the Sherritt Gordon plant.


558 Chapter 13 NEL

13.113.1 Oxidation and ReductionIn prehistoric times, people learned to extract metals from rocks and minerals (Figure 1).This discovery initiated both the technology of metallurgy and humanity’s progressionfrom the Stone Age, through the Bronze Age and the Iron Age, to our increasingly tech-nological modern age. Only a few metals, such as gold and silver, exist naturally in theform of a pure element. Most metals exist in a variety of compounds mixed with othersubstances in rocks called ores. Metallurgy is the science and technology of extractingmetals from their naturally occurring compounds and adapting these metals for usefulpurposes. For some metals, the basic procedures are quite simple and were developed earlyin human history; for others, more complex procedures have been developed morerecently. Making steel requires higher temperatures than those provided by a simplewood fire; therefore, this technology developed much later than the making of copperand bronze. A relatively recent example is the production of aluminium. This requiresthe technology to produce current electricity, which was not discovered until the 1800s.In all of these examples, from copper to aluminium, the technology was developed to solvepractical problems before there was any scientific understanding of the processes involved.

5000 4000 3000 2000 1000 B.C.E. C.E. 1000 2000

copperbronze

(copper and tin) ironbrass

(copper and zinc) aluminium

Figure 1The technology of metallurgy has along history, preceding bythousands of years the scientificunderstanding of the processes.

CAREER CONNECTION

Materials/MetallurgicalEngineerBeing a Materials/MetallurgicalEngineer is a challenging andrewarding career. The job involveslearning the properties ofmaterials, as well as developingnew materials to meet specificrequirements. Materials/Metallurgical Engineers also studyand develop solutions for materialfractures and breakages.

Read about the experiences ofNalaine Morin of the Tahltan FirstNation, a metallurgical engineer.


From metallurgy, the term reduction came to be associated with producing metalsfrom their compounds. For example, iron(III) oxide is “reduced” by carbon monoxidegas to iron metal. Tin and copper metals are other examples where a metal compoundis reduced to the metal.

Fe2O3(s) � 3 CO(g) → 2 Fe(s) � 3 CO2(g)

SnO2(s) � C(s) → Sn(s) � CO2(g)

CuS(s) � H2(g) → Cu(s) � H2S(g)

As you can see from these chemical equations, a substance called a reducing agent causesor promotes the reduction of a metal compound to an elemental metal. In the precedingexamples, carbon monoxide is the reducing agent for the production of iron fromiron(III) oxide, carbon (charcoal) is the reducing agent for the production of tin fromtin(IV) oxide, and hydrogen is the reducing agent for the production of copper fromcopper(II) sulfide. These are three of the most common reducing agents used in metal-lurgical processes.



Although the discovery of fire occurred much earlier than that of metal refining, bothdiscoveries advanced the development of civilization significantly. There are also impor-tant similarities in the chemistry behind these technological developments. Of course, itdoes not require a detailed scientific understanding of the processes to use fire to refinemetals (Figure 2). Only in the 18th century did we realize the role of oxygen in burning.Understanding the connection between corrosion and burning is an even more recentdevelopment. Scientists now understand that corrosion, including the rusting of metals,is similar to combustion, although corrosion reactions occur more slowly. Historically,corrosion was considered to be the spontaneous reaction of metals with air (oxygen) to formmetal compounds (such as rust). In effect, corrosion was returning the metal to its nat-ural state as a compound and, therefore, can be considered to be the opposite of metal-lurgy. Chemists eventually called the reactions of substances with oxygen, whether theywere the explosive combustion of gunpowder, the burning of wood, or the slow corrosionof iron, oxidation. As the study of chemistry developed, chemists realized that oxygenwas not the only substance that could cause reactions similar to oxidation reactions. Forexample, metals can be converted to compounds by most nonmetals and by some othersubstances as well. The rapid reaction process we call burning may even take place withgases other than oxygen, such as chlorine or bromine (Figure 3). The term “oxidation” hasbeen extended beyond reactions with oxygen to include a wide range of combustion andcorrosion reactions, such as the following:

2 Mg(s) � O2(g) → 2 MgO(s)2 Al(s) � 3 Cl2(g) → 2 AlCl3(s)

Cu(s) � Br2(g) → CuBr2(s)

A substance that causes or promotes the oxidation of a metal to produce a metal com-pound is called an oxidizing agent. In the reactions shown above, the oxidizing agents areoxygen, chlorine, and bromine. As you will see, an understanding of reduction and oxidation is necessary to explain many seemingly unrelated chemical reactions.

Section 13.1

Figure 2Making samurai swords requires aspecial type of fire (hardwood coalsand bellows) to produce the highertemperatures required. This was atechnological development longbefore any scientific understanding.

Practice1. Using the historical context of metallurgy and corrosion, write an empirical definition

for each of the following terms: (a) reduction (d) reducing agent (b) oxidation (e) metallurgy(c) oxidizing agent (f) corrosion

2. For each of the following, classify the reaction of the metal or metal compound in thehistorical context of reduction or oxidation, and identify the oxidizing agent or thereducing agent. (a) 4 Fe(s) � 3 O2(g) → 2 Fe2O3(s) (b) 2 PbO(s) � C(s) → 2 Pb(s) � CO2(g)(c) NiO(s) � H2(g) → Ni(s) � H2O(l) (d) Sn(s) � Br2(l) → SnBr2(s) (e) Fe2O3(s) � 3 CO(g) → 2 Fe(s) � 3 CO2(g)(f) Cu(s) � 4 HNO3(aq) → Cu(NO3)2(aq) � 2 H2O(l) � 2 NO2(g)

3. List three reducing agents used in metallurgy.

4. What class of elements behaves as oxidizing agents for metals?

5. In the history of metallurgy, which came first, technological applications or scientificunderstanding? Elaborate on your answer.

6. Technologies that are intended to provide useful products and processes usuallyhave unintended consequences. Using previous knowledge, list some unintendedconsequences that result from the production and refining of metals.

Figure 3Copper metal is oxidized by reactivenonmetals such as bromine.

Putting Out Class D FiresA Class D fire includescombustible metals such asmagnesium and the alkali metals.These metals burn at very hightemperatures and water can makethe fire much worse because ofviolent reactions. Carbon dioxideextinguishers don’t help eithersince magnesium burns very wellin carbon dioxide. So how can youput out such a fire? Sand is asimple option, but a special fireextinguisher such as MET-L-X,which contains sodium chloride, isthe preferred method.

DID YOU KNOW ??


560 Chapter 13 NEL

Early Metallurgy Metallurgy, which is the process of extracting metals from oresand forming them into useful objects, is an example of an earlytechnology that was developed to solve practical problems.There is considerable archeological evidence that earlyhumans discovered that gold, silver, and copper occurred intheir natural state as nuggets or veins in rocks thousands ofyears ago. Stone tools were used to extract these metals fromrocks and to hammer them into various objects. In particular,North American Aboriginal peoples developed the technologyto extract copper from large copper pits in the Lake Superiorregion. The copper was used to make various ornaments, tools,and weapons. The problem was that the use of naturallyoccurring metals was limited to areas in which these metalswere easily unearthed.

Over six thousand years ago, a new technology emergedwhen humans discovered that an extremely hot fire built oncertain greenish rocks produced nuggets of solid copper in thecooling ashes. The extraction of metals from ores by heating iscalled smelting. Smelting copper ore to separate the copperfrom copper-bearing ore minerals, such as green malachite,meant that copper became much more widely used. There isevidence that smelting was discovered around 4500 B.C.E. inthe Middle East because copper ornaments became morecommon after that time.

Since the ores of different metals often occur together,humans discovered that the metal smelted from a mixture ofores was much harder than pure copper. Smelting a mixture ofcopper and tin ores formed a hard copper–tin alloy known asbronze. Blades made from bronze could hold an edge very welland, if necessary, could be pounded back into shape. Bronzewas one of the first metallic alloys widely used for containers,tools, weapons (Figure 4), and armour. The technology ofsmelting copper and formulating bronze diffused acrossEurope during the period of history known as the Bronze Age.

Case StudyCase Study

The Bronze Age continued until humans developed thetechnology for smelting iron, which ushered in the Iron Age.

Case Study Questions

1. Is the discovery of metallurgy a scientific or technologicalachievement? Justify your answer.

2. What kinds of practical problems were solved or improvedwith the use of metals versus stone or wood? What newproblems were created with the development of metalimplements?

3. The following reaction equation represents the ancientprocess for smelting copper from malachite,Cu2CO3(OH)2(s), in a wood fire. Identify the reducingagent and the substance reduced.

Cu2CO3(OH)2(s) � 2 CO(g) →2 Cu(s) � H2O(g) � 3 CO2(g)

Extension

4. How was the process of trial and error used by early peoples to extract metals from their ores?

Figure 4Bronze spearheads fromthe third millennium B.C.E.


Purpose Design AnalysisProblem Materials Evaluation (1)Hypothesis ProcedurePrediction Evidence

To perform this investigation, turn to page 601.

Single Replacement ReactionsThis investigation is a review of single replacement reactions inpreparation for the development of a theory of oxidation andreduction. As part of the Design, include diagnostic tests (as inAppendix C.4) for the predicted products.

PurposeThe purpose of this investigation is to use the single replacementreaction generalization to predict and analyze the reactants andproducts.

ProblemWhat are the products of the single replacement reactions for thefollowing pairs of reactants? (a) copper and aqueous silver nitrate(b) aqueous chlorine and aqueous sodium bromide(c) magnesium and hydrochloric acid(d) zinc and aqueous copper(II) sulfate(e) aqueous chlorine and aqueous potassium iodide

INVESTIGATION 13.1 Introduction Report Checklist



Section 13.1

Figure 6A common single replacementreaction is the reaction of activemetals with an acid, such as zincwith hydrochloric acid.

WEB Activity

Canadian Achievers—Henry Taube and Rudolph MarcusHenry Taube (Figure 5(a)) was born and educated in Saskatchewan. Rudolph Marcus (Figure 5(b)) was born and educated in Quebec. Taube and Marcus both contributed to thetheory of electron transfer reactions. Each received a Nobel Prize in Chemistry, in 1983 and1992, respectively, for their work.

1. In one sentence, describe the work done by each chemist as cited by the NobelCommittee.

2. Describe one difference between Taube’s and Marcus’ work.


Figure 5(a) Henry Taube (1915–2005)

(b) Rudolph Marcus (1923– )

Electron Transfer Theory Single replacement reactions are reactions in which one element replaces another elementin a compound. These reactions are useful to investigate first because they provide arelatively simple introduction to the modern theoretical definitions of oxidation andreduction. Imagine that a reaction is a combination of two parts called half-reactions. Ahalf-reaction represents what is happening to one reactant in an overall reaction. It tellsonly part of the story. Another half-reaction is required to complete the description ofthe reaction. Splitting a chemical reaction equation into two parts not only makes theexplanations simpler, but also leads to some important applications, which are discussedlater in this unit.

For example, when zinc metal is placed into a hydrochloric acid solution, gas bub-bles form as the zinc slowly dissolves (Figure 6). Diagnostic tests show that the gas ishydrogen and that zinc ions are present in the resulting solution. Notice that zinc metalis oxidized to zinc chloride. This is a corrosion of zinc caused by the hydrochloric acid.

Zn(s) � 2 HCl(aq) → ZnCl2(aq) � H2(g)

What happens to the zinc and what happens to the hydrochloric acid? We can look athalf-reactions to answer these questions. The zinc atoms in the solid, Zn(s), are con-verted to zinc ions in solution, Zn2�(aq). Atomic theory requires that the zinc atomslose two electrons, as shown by the following half-reaction equation:

Zn(s) → Zn2�(aq) � 2 e�

Simultaneously, hydrogen ions in the solution gain electrons and are converted intohydrogen gas, as shown below:

2 H�(aq) � 2 e� → H2(g)

Notice that both of these half-reaction equations, or half-reactions, are balanced by mass(same number of atoms/ions of each element on both sides) and by charge (same totalcharge on both sides). A half-reaction is a balanced chemical equation that representseither a loss or gain of electrons by a substance.

In a laboratory, single replacement reactions in aqueous solution are easier to studythan the metallurgy or corrosion reactions discussed earlier in this chapter. However,all of these reactions share a common feature: ions are converted to atoms and atoms are converted to ions. For example, consider the reduction of aqueous silver nitrate tosilver metal in the presence of solid copper (Figure 7 on page 562). According to atomictheory, silver atoms are electrically neutral particles (47p�, 47e�) and silver ions are

(a)

(b)


562 Chapter 13 NEL

charged particles (47p�, 46e�). Therefore, in this reaction, a single electron is requiredto convert a silver ion into a silver atom. The following half-reaction equation explainsthe reduction of silver ions using the theoretical rules for atoms and ions:

Ag�(aq) � e� → Ag0(s) (reduction)

According to modern theory, the gain of electrons is called reduction.

Although this theoretical definition of reduction agrees with current atomic theory,it does not explain where the electrons come from. As crystals of silver metal are produced,the solution becomes blue, indicating that copper atoms are being converted to copper(II)ions. According to atomic theory, copper atoms (29p�, 29e�) must each be losing twoelectrons as they form copper(II) ions (29p�, 27e�):

Cu0(s) → Cu2�(aq) � 2 e� (oxidation)

According to modern theory, the loss of electrons is called oxidation.

Evidence shows that the silver-coloured solid and the blue colour of the solution aresimultaneously formed near the surface of the copper metal. Scientists believe, there-fore, that the electrons required by the silver ions are supplied when silver ions collidewith copper atoms on the metal surface.

The theory of electron transfer requires that the total number of electrons gained in areaction must equal the total number of electrons lost, and that oxidation and reductionare separate processes. This theoretical description requires oxidation and reduction tooccur simultaneously rather than sequentially. Oxidation–reduction reactions are oftensimply called “redox” reactions. A redox (reduction–oxidation) reaction is a chemicalreaction in which electrons are transferred between entities.

The equations for reduction and oxidation half-reactions, and the overall (net) ionicequation summarize the electron transfer that is believed to take place during a redox reaction. As in other chemical reactions, the net equation must be balanced. We willlearn how to write balanced half-reactions and net equations in Sample Problem 13.1.

Redox MnemonicsA mnemonic is a word or group ofwords used to help rememberinformation. “LEO the lion saysGER” is a mnemonic to helpremember that “Loss of Electronsis Oxidation” and “Gain ofElectrons is Reduction.” Anothermnemonic is “OIL RIG,” whichtranslates as “Oxidation Is Loss,Reduction Is Gain.”

DID YOU KNOW ??

Learning TipA common difficulty in writinghalf-reaction equations isdeciding on which side of theequation the electrons shouldappear. There are two ways toapproach this. You can use yourtheoretical knowledge of atomsand ions to determine whetherthe species is losing or gainingelectrons. The alternative is toadd up the electric charges onboth sides of the equationarrow. The total must be thesame on both sides. Try this forthe half-reaction equations onthis page.

Figure 7A piece of copper before it is placedinto a beaker of silver nitratesolution (left). Note the changesafter the reaction has occurred(right). The blue colour of thesolution indicates Cu2�(aq) ions arepresent.

Unit 7 - Ch 13 CHEM30 11/3/06 9:38 AM Page 562


Section 13.1

Write the balanced half-reaction equations and a balanced net equation for the reactionof copper metal with aqueous silver nitrate.

To show that the number of electrons gained equals the number of electrons lost in twohalf-reaction equations, it may be necessary to multiply one or both half-reactionequations by a coefficient to balance the electrons. In this example, the silver half-reactionequation must be multiplied by 2.

Cu(s) → Cu2�(aq) � 2 e� (two electrons lost by one atom)

2 [Ag�(aq) � e� → Ag(s)] (two electrons gained by two ions)

Now, add the half-reaction equations and cancel the terms that appear on both sides ofthe equation to obtain the net ionic equation.

2 Ag�(aq) � 2 e�� Cu(s) → 2 Ag(s) � Cu2�(aq) � 2 e��

2 Ag�(aq) � Cu(s) → 2 Ag(s) � Cu2�(aq)

Silver ions are reduced to silver metal by reaction with copper metal. Simultaneously,copper metal is oxidized to copper(II) ions by reaction with silver ions (Figure 8).

2 Ag�(aq) � Cu(s) → 2 Ag(s) � Cu2�(aq)

SAMPLE problem 13.1

Write and label two balanced half-reaction equations to describe the reaction of zincmetal with aqueous lead(II) nitrate, as given by the following chemical equation:

Zn(s) � Pb(NO3)2(aq) → Pb(s) � Zn(NO3)2(aq)

Solution

Zn(s) → Zn2�(aq) � 2 e� (oxidation)

Pb2�(aq) � 2 e� → Pb(s) (reduction)

COMMUNICATION example 1

Ag

AgCu Cu

Cu Cu

Cu

Cu Cu

Cu Cu

Cu2�

Cu Cu

Cu Cu

Cu Cu

Cu Cu

Cu Cu

Ag�

Ag�

Figure 8A model of the reaction of coppermetal and silver nitrate solutionillustrates aqueous silver ionsreacting at the surface of a copperstrip.

Learning TipWhen you are cancelling termsfor a net ionic equation, theterms must be identical,including their states of matter.Electrons must always cancelcompletely. This is because theelectrons that appear in eachhalf-reaction equation are thesame electrons. They are theelectrons that transfer from oneentity to another.

oxidized to metal ion

reduced to metal

To evaluate the theory of oxidation and reduction, you should look at its logical con-sistency with other accepted theories and definitions. The theoretical definitions of oxi-dation and reduction are consistent with the historical, empirical definitions presentedearlier in this chapter; for example, a compound is reduced to a metal and a metal isoxidized to form a compound. Redox theory is also consistent with accepted atomictheory and the collision–reaction theory. Most importantly, redox theory explains theobservations made by scientists.

WEB Activity

Simulation—Redox ReactionUse the computer simulation of the reaction between lead atoms and silver ions to determinewhich entity loses electrons and which gains electrons. Note the number of electrons lost orgained by each kind of atom/ion. Write the oxidation and reduction half-reaction equations.



564 Chapter 13 NEL

• A redox reaction is a chemical reaction in which electrons are transferred between entities.

• The total number of electrons gained in the reduction equals the total number of electrons lost in the oxidation.

• Reduction is a process in which electrons are gained by an entity.

• Oxidation is a process in which electrons are lost by an entity.

• Both reduction and oxidation are represented by balanced half-reaction equations.

SUMMARY Electron Transfer Theory

Practice7. Write a theoretical definition for each of the following terms:

(a) redox reaction (b) reduction (c) oxidation

8. Write a pair of balanced half-reaction equations—one showing a gain of electrons andone showing a loss of electrons—for each of the following reactions: (a) Zn(s) � Cu2�(aq) → Zn2�(aq) � Cu(s)(b) Mg(s) � 2 H�(aq) → Mg2�(aq) � H2(g)

9. For each of the following, write and label the oxidation and reduction half-reactionequations. Ignore spectator ions. (a) Ni(s) � Cu(NO3)2(aq) → Cu(s) � Ni(NO3)2(aq)(b) Pb(s) � Cu(NO3)2(aq) → Cu(s) � Pb(NO3)2(aq)(c) Ca(s) � 2 HNO3(aq) → H2(g) � Ca(NO3)2(aq)(d) 2 Al(s) � Fe2O3(s) → 2 Fe(l) � Al2O3(s) (Figure 9)

10. We have only looked at one type of single replacement reaction—a metal displacinganother metal from an ionic compound. A nonmetal can also displace anothernonmetal from an ionic compound. For example,

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

Using your knowledge of atoms and ions and the ideas presented in this chapter,write a pair of balanced half-reaction equations for this reaction—one showing a gainof electrons and one showing a loss of electrons.

11. Ionic compounds can react in double replacement reactions. For example,

FeCl3(aq) � 3 NaOH(aq) → Fe(OH)3(s) � 3 NaCl(aq)

According to ideas discussed in this chapter, has a redox reaction taken place in thereaction above? Explain your answer.

Figure 9The reduction of iron(III) oxide byaluminium is called the "thermite"reaction. Because this reaction israpid and very exothermic, moltenwhite-hot iron is produced. Here afalling aluminium wrenchmomentarily sparks a thermitereaction when it strikes a rusted ironblock.

DID YOU KNOW ??CompartmentalizationThe tradition in Western science isgenerally to break large systemsdown into smaller and smallerpieces to describe and explain them.Science is broken down into variousareas such as biology, chemistry,and physics, which can then befurther subdivided. Onecompartment of chemistry is redox,which includes severalsubcategories.

In contrast, Aboriginal science ismore holistic: the whole of nature,from the smallest parts to thecosmos, is connected. Viewingnature this way gives us a differentperspective on current problemssuch as pollution and climatechange. Recently, there is a trend inWestern science toward a moreholistic view after decades ofcompartmentalization.

Writing Complex Half-Reaction EquationsAlthough most metals and nonmetals have relatively simple half-reaction equations,polyatomic ions and molecular compounds undergo more complicated oxidation andreduction processes. In most of these processes, the reaction takes place in an aqueoussolution that is very often acidic or basic. Experimental evidence shows that water mol-ecules, hydrogen ions, and hydroxide ions play an important role in these half-reactions.A method of writing half-reactions for polyatomic ions and molecular compoundsrequires that water molecules and hydrogen or hydroxide ions be included. This method,illustrated in the following sample problem, is sometimes called the “half-reaction” or“ion–electron” method.



In a basic solution, the concentration of hydroxide ions greatly exceeds that of hydrogenions. For basic solutions, we will develop the half-reaction as if it occurred in an acidicsolution, and then convert the hydrogen ions into water molecules using hydroxide ions.This trick works because a hydrogen ion and a hydroxide ion react in a 1:1 ratio to forma water molecule. The following sample problem illustrates the procedure for writing half-reaction equations that occur in basic solutions.

Section 13.1

Nitrous acid can be reduced in an acidic solution to form nitrogen monoxide gas. What isthe reduction half-reaction for nitrous acid?

The first step is to write the reactants and products.

HNO2(aq) → NO(g)

If necessary, you should balance all atoms other than oxygen and hydrogen in this partial equation. In this example, there is only one nitrogen atom on each side.

Next, add water molecules, present in an aqueous solution, to balance the oxygenatoms.

HNO2(aq) → NO(g) � H2O(l)

Because the reaction takes place in an acidic solution, hydrogen ions are available.These can be used to balance the hydrogen on both sides of the equation.

H�(aq) � HNO2(aq) → NO(g) � H2O(l)

At this stage, all of the atoms should be balanced, but the charge on both sides will notbe balanced. Add an appropriate number of electrons to balance the charge. Becauseelectrons carry a negative charge, they are always added to the less negative, or morepositive, side of the half-reaction.

e� � H�(aq) � HNO2(aq) → NO(g) � H2O(l)

This balanced half-reaction equation represents a gain of electrons, or a reduction of thenitrous acid. Check to make sure that both the atom symbols and the charge are balanced.

SAMPLE problem 13.2


BIOLOGY CONNECTION

Metabolic Redox

There are many biologicalprocesses, including metabolismand cellular respiration, that areredox reactions. You will learnabout them in a biology course.

Copper metal can be oxidized in a basic solution to form copper(I) oxide. What is the half-reaction for this process?

Following the same steps as before, we write the formula and balance the atoms, otherthan oxygen and hydrogen. Here, the copper atoms must be balanced.

2 Cu(s) → Cu2O(s)

Next, balance the oxygen using water molecules and balance the hydrogen usinghydrogen ions, assuming, for the moment, an acidic solution. Balance the charge usingelectrons.

H2O(l) � 2 Cu(s) → Cu2O(s) � 2 H�(aq) � 2 e�

SAMPLE problem 13.3

Learning TipYou can also balance half-reaction equations in a basicsolution by following thesesteps.1. Balance atoms other than

oxygen and hydrogen.2. Add hydroxide ions to

balance the oxygen.3. Add water to balance the

hydrogen.4. Add electrons to balance the

charge.

The difficulty with this methodis that adding hydroxide ionsand water simultaneouslyaffects both the oxygen andhydrogen count.


566 Chapter 13 NEL

Because the half-reaction occurs in a basic solution, add the same number of hydroxideions as there are hydrogen ions to both sides of the equation. This is done to maintain thebalance of mass and charge.

2 OH�(aq) � H2O(l) � 2 Cu(s) → Cu2O(s) � 2 H�(aq) � 2 e� � 2 OH�(aq)

Combine equal numbers of hydrogen ions and hydroxide ions to form water molecules.

2 OH�(aq) � H2O(l) � 2 Cu(s) → Cu2O(s) � 2 H2O(l) � 2 e�

Finally, cancel equal amounts of H2O(l) and anything else that is the same from bothsides of the equation. Check that the atom symbols and charge are balanced.

2 OH�(aq) � H2O(l)� � 2 Cu(s) → Cu2O(s) � 2 H2O(l)� � 2 e�

2 OH�(aq) � 2 Cu(s) → Cu2O(s) � H2O(l) � 2 e�

Chlorine is converted to perchlorate ions in an acidic solution. Write the half-reactionequation. Is this half-reaction an oxidation or a reduction?

Solution

8 H2O(l) � Cl2(aq) → 2 ClO4�(aq) � 16 H�(aq) � 14 e�

According to redox theory, this half-reaction is an oxidation.


Aqueous permanganate ions are reduced to solid manganese(IV) oxide in a basic solution. Write the half-reaction equation. Is the half-reaction an oxidation or a reduction?

Solution

4 OH�(aq) � 4 H�(aq) � MnO4�(aq) � 3 e� → MnO2(s) � 2 H2O(l) � 4 OH�(aq)

4 H2O(l) � MnO4�(aq) � 3 e� → MnO2(s) � 2 H2O(l) � 4 OH�(aq)

MnO4�(aq) � 2 H2O(l) � 3 e� → MnO2(s) � 4 OH�(aq)

According to redox theory, this half-reaction is a reduction.


Learning TipRecall that loss of electrons isoxidation (LEO) and gain ofelectrons is reduction (GER).

Practice12. For each of the following, complete the half-reaction equation and classify it as an

oxidation or a reduction. (a) dinitrogen oxide to nitrogen gas in an acidic solution(b) nitrite ions to nitrate ions in a basic solution(c) silver(I) oxide to silver metal in a basic solution(d) nitrate ions to nitrous acid in an acidic solution(e) hydrogen gas to water in a basic solution

Redox ReactionsThis short movie and theaccompanying exercise will helpyou check your understanding ofredox reactions.


EXTENSION +



Section 13.1

Step 1: Write the chemical formulas for the reactants and products.Step 2: Balance all atoms, other than O and H.Step 3: Balance O by adding H2O(l).Step 4: Balance H by adding H�(aq).Step 5: Balance the charge on each side by adding e� and cancel anything

that is the same on both sides.

For basic solutions only:Step 6: Add OH�(aq) to both sides to equal the number of H�(aq) present.Step 7: Combine H�(aq) and OH�(aq) on the same side to form H2O(l).

Cancel equal amounts of H2O(l) from both sides.

SUMMARY Writing Half-Reaction Equations

Section 13.1 Questions1. Compare the historical empirical definitions of oxidation

and reduction with the modern theoretical definitions.

2. According to modern theory, explain what is meant by aredox reaction.

3. What is a half-reaction and how does it relate to the overallchemical reaction equation?

4. Describe what happens to the electrons that are lost byone entity in a redox reaction.

5. What common type of chemical reaction is generally not aredox reaction? Include a brief theoretical explanation.

6. Write and label a pair of balanced half-reaction equationsfor each of the following reactions. (a) Pb(s) � Cu2�(aq) → Pb2�(aq) � Cu(s)(b) Cl2(aq) � 2 Br�(aq) → 2 Cl�(aq)� Br2(l)(c) the removal of silver tarnish (Figure 10) in the single

replacement reaction of silver sulfide solid withaluminium

7. For each of the following applications, complete the half-reaction equation and classify it as an oxidation or areduction. (a) bacterial action in soil: ammonia to nitrite ions in an

acidic environment

(b) pulp and paper bleaching: hydrogen peroxide to waterin an acidic solution

(c) alkaline battery: manganese(IV) oxide to manganese(III)oxide in a basic environment

Extension

8. Photosynthesis is a complex electron-transfer process thatis essential for almost all life forms on Earth. Write and labelthe half-reaction equations for the “light reaction” and the“dark reaction.”

9. One of the earliest technologies is the use of fire. There arevarious stories in many cultures about how people learnedabout fire, such as the North American Aboriginal story ofhow the coyote steals fire. These stories serve as metaphorsfor the technology that we use today. Summarize the legendof the coyote and fire. How does this metaphor relate to thegoals and effects of technology today?

10. Archaeometallurgy is the study of ancient metallurgy usingmodern analytical techniques (Figure 11). Give someexamples of research in the field. What metals and timeperiods have been studied? Can a metal from one mine bedistinguished from the same metal from another mine?How is this information used in archaeological studies?

Figure 11Aslihan Yener (left)pioneered the use ofmodern X-ray techniques toidentify metals from as early as8000 B.C.E.

Figure 10If tarnished silver is placed on aluminium foil in a hotelectrolyte solution, a redox reaction converts the tarnishback into silver metal.





568 Chapter 13 NEL

13.213.2 Predicting Redox ReactionsA redox reaction may be explained as a transfer of valence electrons from one substanceto another. Evidence indicates that the majority of atoms, molecules, and ions are stableand do not readily release electrons. Since two entities must be involved in an electrontransfer, chemists explain this transfer as a competition for electrons. Using a tug-of-waranalogy, each entity pulls on the same electrons. If one entity is able to pull electrons awayfrom the other, a spontaneous reaction occurs (Figure 1). Otherwise, no reaction occurs(Figure 2). In the spontaneous reaction of copper(II) ions and zinc metal, the Cu2� ionis electron deficient and pulls electrons from a Zn atom. We can explain the reaction evi-dence using the idea that Cu2� “pulls harder” on Zn’s electrons than Zn does. Cu2� “wins”the two valence electrons from a Zn atom. A successful electron transfer has occurred.

Without mixing all possible reactants and observing any evidence of reaction, how canwe predict if a reaction will occur? If a reaction occurs, what will be produced? Theanswers to these questions cannot be obtained easily using redox theory. By observingmany successful and unsuccessful reactions, patterns emerge and empirical generaliza-tions can be made.

Oxidizing and Reducing AgentsBefore we look at these patterns, we need to define some terms commonly used bychemists. When discussing possible reactants and comparing their reactivities, it is cus-tomary and convenient to classify the reactants in a redox reaction. This classification orig-inated historically, but is now defined theoretically in terms of an ability to lose or gainelectrons. In any redox reaction, an electron transfer occurs. This means that one reac-tant is oxidized and one reactant is reduced.

e�

X � Y → productselectron change loses gains

(is oxidized) (is reduced)

Examples: Zn(s) � Cu2�(aq) → Zn2�(aq) � Cu(s)2 Br�(aq) � Cl2(g) → Br2(l) � 2 Cl�(aq)

Rather than saying “the reactant that is oxidized” and “the reactant that is reduced,”chemists use the terms reducing agent (RA) and oxidizing agent (OA). These terms orig-inated in the early history of metallurgy and corrosion. For example, to “reduce” a largervolume of iron(III) oxide to a smaller volume of pure iron, a substance called a reducingagent was required, such as CO(g).

reducing agent � Fe2O3(s) → Fe(s) � other products

Similarly, oxidation was originally associated with an oxidizing agent. For example, ametal could be oxidized by certain substances called oxidizing agents. At first, oxygen wasthe only known oxidizing agent, but others (e.g., halogens) can also oxidize, or corrode,metals.

oxidizing agent � Mg(s) → MgO(s) � other products

Figure 1Copper(II) ions react spontaneouslywith zinc metal. A copper(II) ion hasa stronger attraction for the valenceelectrons of a zinc atom than zincdoes.

Figure 2When a strip of copper is placed in asolution of nickel(II) ions, the greennickel(II) ion colour remains and thecopper metal does not react.Collisions between copper atomsand nickel(II) ions apparently do notresult in the transfer of electrons.

reduction

oxidation



The terms oxidizing and reducing agents developed separately, long before any redoxtheory of electron transfer emerged. Today, chemists routinely think in terms of electrontransfer to explain redox reactions. A redox reaction is recognized as an electron transferbetween an oxidizing agent and a reducing agent (Figure 3). Based on this idea, chemistssay that a reducing agent causes reduction by donating (losing) electrons to another sub-stance in a redox reaction. During this process, the reducing agent is oxidized. An oxi-dizing agent causes oxidation by removing (gaining) electrons from another substancein a redox reaction. During this process, the oxidizing agent is reduced. It is importantto note that oxidation and reduction are processes, and oxidizing agents and reducing agentsare substances. For example,

2 Ag�(aq) � Cu(s) → 2 Ag(s) � Cu2�(aq)OA RA

Section 13.2

OA + RA

e—

Figure 3In all redox reactions, electrons aretransferred from a reducing agent(RA) to an oxidizing agent (OA).

Table 1 Reactivities of Metal Ions with Metals

Ions Ag�(aq) Cu2�(aq) Pb2�(aq) Zn2�(aq)

reacted with Cu(s), Pb(s), Zn(s) Pb(s), Zn(s) Zn(s) none

number ofreactions 3 2 1 0

reactivity order most least

Purpose Design AnalysisProblem Materials Evaluation (1, 2, 3)Hypothesis ProcedurePrediction Evidence


Spontaneity of Redox ReactionsIn previous units in this textbook, we assumed that all chemicalreactions are spontaneous; that is, they occur once the reactantsare placed in contact, without a continuous addition of energy tothe system. Spontaneous redox reactions in solution generallyprovide visible evidence of a reaction within a few minutes.

PurposeThe purpose of this investigation is to test the assumption that allsingle replacement reactions are spontaneous.

ProblemWhich combinations of copper, lead, silver, and zinc metals andtheir aqueous metal ion solutions produce spontaneous reactions?

DesignA drop of each solution is placed in separate locations on a cleanarea of each of the four metal strips.


Development of a Redox TableSome redox reactions, such as single replacement reactions, are easy to study experi-mentally. The evidence of a reaction is immediately obvious and the interpretation of anelectron transfer is relatively simple. In the past, you have generally assumed that allsingle replacement reactions are spontaneous. However, the evidence you obtained inInvestigation 13.2 clearly shows that this assumption is unacceptable as only six of thecombinations led to a reaction. The question that arises is, “How do you know when achemical reaction will occur spontaneously without actually doing the reaction?”

Let’s look at the combinations of metals and metal ions used in Investigation 13.2.Copper, lead, silver, and zinc metals were combined one at a time with each of copper(II),lead(II), silver, and zinc ion solutions. Based on the evidence collected, we can rank theability of the metal ions to react with the metals (Table 1).

reduced to metal

oxidized to metal ion


570 Chapter 13 NEL

The most reactive metal ion, Ag�(aq), has the greatest tendency to gain electrons. Onthe other hand, Zn2�(aq) shows no tendency to gain electrons in the combinationstested. Therefore, the order of reactivity is also the order of strength as oxidizing agents.

strongest oxidizing agent Ag�(aq) � e� → Ag(s)

Cu2�(aq) � 2 e� → Cu(s)

Pb2�(aq) � 2 e� → Pb(s)

weakest oxidizing agent Zn2�(aq) � 2 e� → Zn(s)

The order of reactivity of the four metals can be obtained in a similar way (Table 2).

The most reactive metal, Zn(s), has the greatest tendency to lose electrons and Ag(s)shows no tendency to lose electrons in the combinations tested. Metals behave as reducingagents and so Zn(s) is the strongest reducing agent among those tested.

strongest reducing agent Zn(s) → Zn2�(aq) � 2 e�

Pb(s) → Pb2�(aq) � 2 e�

Cu(s) → Cu2�(aq) � 2 e�

weakest reducing agent Ag(s) → Ag�(aq) � e�

In these four reactions, the metal ions are the oxidizing agents and the silver ion isthe strongest oxidizing agent (SOA) because it is the most reactive in our group. Thetwo lists can be summarized using a single set of half-reactions (Table 3).

In Table 3, the metal ions are on the left side of the equations and the metal atoms areon the right side. For metal ions (the oxidizing agents), the half-reaction equations areread from left to right in the table. For metal atoms (the reducing agents), the half-reac-tion equations are read from right to left.

Table 2 Reactivities of Metals with Metal Ions

Metals Zn(s) Pb(s) Cu(s) Ag(s)

reacted with Ag�(aq), Cu2�(aq), Pb2�(aq) Ag�(aq), Cu2�(aq) Ag�(aq) none

number ofreactions 3 2 1 0

reactivity order most least

Learning TipThe double arrows indicate thathalf-reactions may be readfrom left to right (top arrow) orfrom right to left (bottomarrow).

SOA

SRA

Table 3 Relative Strengths of Oxidizing and Reducing Agents

OA � n e�0 RA

decreasing Ag�(aq) � e�0 Ag(s) decreasing

reactivity of Cu2�(aq) � 2 e�0 Cu(s) reactivity of

oxidizing Pb2�(aq) � 2 e�0 Pb(s) reducing

agents Zn2�(aq) � 2 e�0 Zn(s) agents

Learning TipIf a positively charged metal ionreacts, then it is usuallyconverted to a metal atom.According to redox theory, thisrequires a gain of electrons andhence the metal ion is behavingas an oxidizing agent. Similarly,if a metal atom reacts, then it isalways converted to a positivelycharged ion by losing electrons.Metals always behave asreducing agents.

WEB Activity

Web Quest—Piercings: A Rash DecisionMost Canadians wear some type of jewellery. However, some individuals react to their jewellery.This Web Quest explores the chemistry behind these reactions. What can you do to protectyourself from your jewellery?




Section 13.2

Practice1. Oxidation and reduction are processes, and oxidizing agents and reducing agents are

substances. Explain this statement.

2. If a substance is a very strong oxidizing agent, what does this mean in terms ofelectrons?

3. If a substance is a very strong reducing agent, what does this mean in terms ofelectrons?

4. The terms, oxidizing agent and reducing agent, may be confusing especially whenused in conjunction with the terms, oxidation and reduction. The key to avoidingconfusion is to focus on the word “agent.” Use a dictionary or write a definition inyour own words of the term “agent.” How can a banker (loans agent) be used as ananalogy for a reducing agent?

5. List the metal(s) that react spontaneously with a copper(II) ion solution.

6. Which metal(s) did not appear to react with a copper(II) ion solution?

7. Start with the position of Cu2�(aq) in Table 3 and note the position of the metal(s)that reacted and the metal(s) that did not react. For a metal that reacts spontaneouslywith Cu2�(aq), where does the metal appear on a table of reduction half-reactions (Table 3)?

8. Repeat questions 5 to 7 for the Pb2�(aq) ion.

9. Your answers to questions 7 and 8 form an empirical hypothesis that can be tested bymaking predictions for the other metal ions. Use Table 3 to predict which of thereactions should be spontaneous. According to the evidence from Investigation 13.2,are your predictions correct? Is your hypothesis verified?

10. An experiment similar to the example of metals and metal ions was conducted usinghalogens and halide ions. Prepare a redox table of half-reaction equations similar toTable 3 for the halogens.

Evidence

Only three combinations produced evidence of a reaction (Figure 4, Table 4).

Table 4 Reactions of Halogens with Solutions of Halides

Br2(aq) Cl2(aq) I2(aq)

Br�(aq) no reaction yellow-brown no reaction

Cl�(aq) no reaction no reaction no reaction

I�(aq) pink/purple pink/purple no reaction

Learning TipIf a nonmetal reacts, then it isusually converted to anegatively charged ion bygaining electrons. Nonmetalsalways behave as oxidizingagents. Negatively chargednonmetal ions lose electronswhen they react; therefore,nonmetal ions behave asreducing agents.

• An oxidizing agent causes oxidation by removing (gaining) electrons from anothersubstance in a redox reaction. In this process, the oxidizing agent is reduced.

• A reducing agent promotes reduction by donating (losing) electrons to another sub-stance in a redox reaction. In this process, the reducing agent is oxidized.

• A table of relative strengths of oxidizing and reducing agents—more simply knownas a redox table—is, by convention, listed as reductions (from left to right) in theform: OA � n e�

0RA, with the strongest oxidizing agent at the top left and strongestreducing agent at the bottom right of the table.

SUMMARY Oxidizing and Reducing Agents

Figure 4None of the combinations ofaqueous solutions of chlorine,bromine, and iodine with theircorresponding halides show anyevidence of reaction except for thereaction between (a) bromine andiodide ions, (b) chlorine andbromide ions, and (c) chlorine andiodide ions.

(a)

(b)

(c)

Refer to Table 1 (on page 569), and Tables 2 and 3 (on page 570) to answer questions 5 to 9.


572 Chapter 13 NEL

The Spontaneity RuleEvidence obtained from the study of many redox reactions has been used to establish ageneralization, called the redox spontaneity rule. The redox spontaneity rule statesthat a spontaneous redox reaction occurs only if the oxidizing agent (OA) is above thereducing agent (RA) in a table of relative strengths of oxidizing and reducing agents.Figure 5 illustrates how you can use the rule, along with a table of oxidizing and reducingagents, to predict whether a reaction is spontaneous or nonspontaneous (no reaction).

Another Method for Building Redox TablesOnce a spontaneity rule is developed from experimental evidence, the rule may be usedto generate redox tables. The evidence to be analyzed in this case is a net ionic equa-tion, accompanied by observations of spontaneity. In the following method, the spon-taneity rule, rather than the number of reactions observed, is used to order the oxidizingand reducing agents to produce a redox table. The procedure for this type of analysisand synthesis is illustrated by Sample Problem 13.4.

OA

RA

+spontaneousreaction

OA

RA

+nonspontaneousreaction

Figure 5The redox spontaneity rule

Purpose Design AnalysisProblem Materials EvaluationHypothesis ProcedurePrediction Evidence

Building a Redox TableSuppose that a research team is developing a table of relativestrengths of oxidizing and reducing agents. One team memberhad completed an investigation summarized in Table 3, page 570,and another had completed the investigation reported in Practicequestion 10, page 571. A third member used the combination ofmetals, nonmetals, and solutions shown below. By completing thisexercise, you will see how scientists have developed extensivetables of relative strengths of oxidizing and reducing agents.

PurposeThe purpose of this lab exercise is to create an extended table ofrelative strengths of oxidizing and reducing agents.

ProblemWhat is the table of relative strengths of oxidizing and reducingagents for the combined results from three experiments?

Evidence

Analysis(a) Use Table 5 to construct a mini-redox table of just those

substances.(b) Compare Table 3, your analysis table from Practice

question 10, and your table. Note that there are severalsubstances that appear in two of these tables. Combine allthree tables in one larger table showing the order of oxidizingand reducing agents.

LAB EXERCISE 13.A Report Checklist

Table 5 Reactions of Metals and Nonmetals with Solutions of Ions

I2(aq) Cu2�(aq) Ag�(aq) Br2(aq)

I�(aq) X X ✓ ✓

Cu(s) ✓ X ✓ ✓

Ag(s) X X X ✓

Br�(aq) X X X X

X no evidence of a redox reaction✓ evidence redox reaction occurred

DID YOU KNOW ??Redox TableThe format of the redox table usedin this textbook is very common.However, some books andreferences show a redox table asreversed (top to bottom), or list theoxidation and reduction half-reactions in alphabetical order or inorder of strength of the agents.



Section 13.2

Three reactions among indium, cobalt, palladium, and copper were investigated. Thereaction equations below indicate that two spontaneous reactions occurred and only onecombination did not react. Using these equations, construct a redox table of half-reactionequations showing the relative strengths of the oxidizing and reducing agents.

3 Co2�(aq) � 2 In(s) → 2 In3�(aq) � 3 Co(s) (spontaneous)

Cu2�(aq)� Co(s) → Co2�(aq) � Cu(s) (spontaneous)

Cu2�(aq) � Pd(s) → no evidence of reaction (nonspontaneous)

To construct a redox table from this information, work with one equation at a time. Identifythe oxidizing and reducing agents for the first reaction, and arrange them in two columnsusing the spontaneity rule. For the first reaction, this step is shown in Figure 6(a).Co2�(aq) is the oxidizing agent and In(s) is the reducing agent. Since the reaction isspontaneous, the oxidizing agent is above the reducing agent in the list.

In the second reaction, Cu2�(aq) is the oxidizing agent and Co(s) is the reducing agent.This reaction is also spontaneous; therefore, Cu2�(aq) is above Co(s) in the list. Since ametal appears on the same line as its ion in a redox table, add Co(s) and extend the list asshown in Figure 6(b).

No reaction occurs for the third pair of reactants. If a reaction had occurred, Cu2�(aq)would be the oxidizing agent and Pd(s) would be the reducing agent. As this reaction isnot spontaneous, the oxidizing agent appears below the reducing agent. Figure 6(c)shows the list extended to include Pd(s). To complete the table, write balanced half-reaction equations for each oxidizing/reducing agent pair.

SOA Pd2�(aq) � 2e�0 Pd(s)

Cu2�(aq) � 2e�0 Cu(s)

Co2�(aq) � 2e�0 Co(s)

In3�(aq) � 3e�0 In(s) SRA

SAMPLE problem 13.4

OA RA

Co2+(aq)

Co2+(aq)

In(s)

In(s)

Cu2+(aq)

Cu2+(aq)

Co(s)

Co(s)

In(s)

Pd(s)

(a)

(b)

(c)

Co2+(aq)

Figure 6The relative position of a pair ofoxidizing and reducing agents indicates whether a reaction will bespontaneous.

Learning TipThe nonspontaneity of areaction is communicated inseveral ways: “no evidence ofreaction,” “nonspontaneous,”“no reaction,” or “nonspont,”written over the equation arrow.

Practice11. A student performed the following reactions. Construct a table of relative strengths of

oxidizing and reducing agents.

Co2�(aq) � Zn(s) → Co(s) � Zn2�(aq)

Mg2�(aq) � Zn(s) → no evidence of reaction

12. In a school laboratory, four metals were combined with each of four solutions and thefollowing reactions occurred:

Be(s) � Cd2�(aq) → Be2�(aq) � Cd(s)

Cd(s) � 2 H�(aq) → Cd2�(aq) � H2(g)

Ca2�(aq) � Be(s) → no evidence of reaction

Cu(s) � H�(aq) → no evidence of reaction

Construct a table of relative strengths of oxidizing and reducing agents.

13. Is the redox spontaneity rule empirical or theoretical? Justify your answer.

14. Use the relative strengths of nonmetals and metals as oxidizing and reducing agents,as indicated in the following unbalanced equations, to construct a table of half-reactions.

Ag(s) � Br2(l) → AgBr(s)

Ag(s) � I2(s) → no evidence of reaction

Cu2�(aq) � I�(aq) → no redox reaction

Br2(l) � Cl�(aq) → no evidence of reaction

Explaining RelativeStrengths of AgentsWhy are some metals moreeffective as reducing agents thanothers? The strength of a reducingagent depends on several factors.This audio clip provides a bit morebackground on what makes somemetals stronger reducing agentsthan others.

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574 Chapter 13 NEL

Practice

15. Arrange the following metal ions in order of decreasing strength as oxidizing agents:lead(II) ions, silver ions, zinc ions, and copper(II) ions. How does this order comparewith Table 3 on page 570?

16. What classes of substances (e.g., metals, nonmetals, acidic, basic) usually behave as (a) oxidizing agents? (b) reducing agents?

17. Use atomic theory to explain why nonmetals behave as oxidizing agents and metalsbehave as reducing agents. Is there logical consistency between atomic theory andthe empirically determined table of oxidizing and reducing agents?

18. Trends in the reactivity of elements show that fluorine is the most reactive nonmetal.How does this relate to the position of fluorine in the redox table of oxidizing andreducing agents? State one reason why this element is the most reactive nonmetal.Why is your reason an explanation? (Keep asking a series of “why” questions untilyour theoretical knowledge is expended. Does your theory pass the test of being ableto explain the empirically determined table?)

19. (a) Repeat your answer to question 16 using a Venn diagram (two large intersectingcircles).

(b) Identify three oxidizing agents (other than Fe2�(aq), shown in Figure 7) from theredox table that can also act as reducing agents and record their symbols on yourVenn diagram.

(c) Try to explain this surprising behaviour.

20. Use the redox spontaneity rule to predict whether the following mixtures will showevidence of a reaction; that is, predict whether the reactions are spontaneous. (Donot write the equations for the reaction.) (a) nickel metal in a solution of silver ions(b) zinc metal in a solution of aluminium ions(c) an aqueous mixture of copper(II) ions and iodide ions(d) chlorine gas bubbled into a bromide ion solution(e) an aqueous mixture of copper(II) ions and tin(II) ions(f) copper metal in nitric acid

21. Complete the Prediction, Design, and Materials (including safety precautions) for thefollowing investigation.

PurposeThe purpose of this investigation is to test the order of strengths of oxidizing andreducing agents given on the redox table (Appendix I).

ProblemWhat is the relative order of strengths of oxidizing and reducing agents for aluminium,nickel, lead, cobalt, and their respective aqueous ions?

22. Describe two designs or methods that can be used to build a redox table.

23. From your knowledge, list two metals that are found as elements and two that arenever found as elements in nature. Test your answer by referring to the position ofthese metals in the table of oxidizing and reducing agents.

24. Has empirical or theoretical knowledge been the most useful to you in predicting thespontaneity of redox reactions? Explain.

Fe2+

Fe3+

GER/OA

LEO/RA

Fe

Figure 7Iron(II) ions can either lose or gainelectrons and, therefore, can act aseither reducing agents or oxidizingagents.

Learning TipAlthough you can consult theredox table in Appendix I, it ismuch more efficient tomemorize which classes ofsubstances tend to be oxidizingagents or reducing agents(question 15). This generalpattern helps to speed up theprocess of recognizing oxidizingand reducing agents and isnecessary to classifysubstances that don’t appear inAppendix I.

An Extended Redox TableChemists have analyzed evidence collected in many experiments, to produce an extendedredox table of oxidizing and reducing agents, such as the one found in Appendix I. Thistable represents the combined efforts of many people over many years. A redox table isan important reference for chemists. You can use this table to compare oxidizing andreducing agents, and to predict spontaneous redox reactions.

DID YOU KNOW ??Science and Belief SystemsTraditional knowledge of Aboriginalpeoples is based on observation,experience, testing, teaching, andrecording. Although themethodology may differ, thisempirical basis is not that muchdifferent from Western science.Traditional knowledge, such astraditional ecological knowledge, issometimes criticized within Westernculture because of its connection toAboriginal beliefs. Yet Westernscience also operates under a beliefsystem that sometimes makes itdifficult for novel ideas to becomeaccepted. For example, Arrhenius’ideas were disputed and overruledby the prominent scientists of theday in spite of Arrhenius’ solid evidence.

Use the redox table in Appendix I to answer the following questions.



Predicting Redox Reactions in SolutionArrhenius’ ideas about solutions provide an important starting point for predicting redoxreactions. In solutions, molecules and ions behave approximately independently of eachother. A first step in predicting redox reactions is to list all entities that are present. (Somehelpful reminders are listed in Table 6.) For example, when copper metal is placed intoan acidic potassium permanganate solution, copper atoms, potassium ions, permanganateions, hydrogen ions, and water molecules are all present. Next, using your knowledge of oxi-dizing and reducing agents, and the redox table in Appendix I, label all possible oxidizingand reducing agents in the starting mixture. The permanganate ion is listed as an oxi-dizing agent only in an acidic solution. To indicate this combination, draw an arc betweenthe permanganate and hydrogen ions as shown below, and label the pair as an oxidizingagent. This procedure of listing and identifying entities present is a crucial step in pre-dicting redox reactions.

We can use a redox table to identify the strongest oxidizing and reducing agents in amixture, and then predict which reactions will occur. If we assume that collisions arecompletely random, the strongest oxidizing agent and the strongest reducing agent willreact. (In some cases, further reactions may occur as well, but we will consider only theinitial reaction, unless otherwise specified.) The following instructions allow you tomake the correct and most efficient use of a redox table, such as the one in Appendix I.

• Choose the strongest oxidizing agent present in your mixture by starting at the topleft corner of a redox table and going down the list until you find the oxidizingagent that is in your mixture.

• Choose the strongest reducing agent in your mixture by starting at the bottomright corner of the table and going up the list until you find the reducing agent thatis in your mixture.

• Read reduction half-reaction equations from left to right (following the forwardarrow).

• Read oxidation half-reaction equations from right to left (following the reversearrow).

• Assume that any substances not present in the table are spectator ions. You do notneed to label or consider these substances.

Section 13.2

Getting Rid of Skunk OdourThe smell of a skunk (Figure 8) iscaused by a thiol compound (R—SH). To deodorize a petsprayed by a skunk, you need toconvert the smelly thiol to anodourless compound. Hydrogenperoxide in a basic solution(usually from sodium bicarbonate)acts as an oxidizing agent tochange the thiol to a disulfidecompound (RS—SR), which isodourless.

DID YOU KNOW ??

Figure 8A skunk’s only defence is its abilityto spray a smelly liquid a distanceof 3 m.

Table 6 Hints for Listing andLabelling Entities

• Aqueous solutions containH2O(l) molecules.

• Acidic solutions contain H�(aq) ions.

• Basic solutions containOH�(aq) ions.

• Some oxidizing and reducingagents are combinations, forexample, MnO4

�(aq) andH�(aq).

• H2O(l) , Fe2�(aq), Cu�(aq),Sn2�(aq), and Cr2�(aq) may actas either oxidizing or reducingagents. Label both possibilitiesin your list.

OAOA

OA OA

Cu(s) K�(aq) MnO4�(aq) H�(aq) H2O(l)

RA RA

Practice25. List all entities initially present in the following mixtures, and identify all possible

oxidizing and reducing agents. (a) A lead strip is placed in a copper(II) sulfate solution.(b) A gold coin is placed in a nitric acid solution.(c) A potassium dichromate solution is added to an acidic iron(II) nitrate solution.(d) An aqueous chlorine solution is added to a phosphorous acid solution.(e) A potassium permanganate solution is mixed with an acidified tin(II) chloride

solution.(f) Iodine solution is added to a basic mixture containing manganese(IV) oxide.


576 Chapter 13 NEL

Suppose a solution of potassium permanganate is slowly poured into an acidified iron(II)sulfate solution. Does a redox reaction occur and, if it does, what is the reaction equation?Describe two diagnostic tests of your prediction.

To make a prediction, the entities initially present are identified as oxidizing agents,reducing agents, or both, as shown below.

Use the redox table in Appendix I to choose the strongest oxidizing agent and thestrongest reducing agent from your list and indicate them as SOA and SRA.

Now, write the half-reaction equation for the reduction of the SOA from the redox table.

MnO4�(aq) � 8 H�(aq) � 5 e� → Mn2�(aq) � 4 H2O(l)

Write the half-reaction equation for the oxidation of the SRA. Remember to read fromright to left in the table.

Fe2�(aq) → Fe3�(aq) � e�

Before combining the half-reaction equations, balance the number of electronstransferred by multiplying one or both half-reaction equations by an integer so that thenumber of electrons gained by the oxidizing agent equals the number of electrons lost bythe reducing agent.

In this case, the iron ion half-reaction must be multiplied by 5. Add the two equations,but remember to cancel any common terms. You can cancel terms as you add (e.g., 5e�)or after you add the two half-reactions.

MnO4�(aq) � 8 H�(aq) � 5 e� → Mn2�(aq) � 4 H2O(l)

5 [Fe2�(aq) → Fe3�(aq) � e� ]

MnO4�(aq) � 8 H�(aq) � 5 Fe2�(aq) → 5 Fe3�(aq) � Mn2�(aq) � 4 H2O(l)

Finally, use the spontaneity rule to predict whether the net ionic equation represents aspontaneous redox reaction. Indicate this by writing “spont.” or “nonspont.” over the equation arrow.

spont.

MnO4�(aq) � 8 H�(aq) � 5 Fe2�(aq) → 5 Fe3�(aq) � Mn2�(aq) � 4 H2O(l)

In this case, we predict that the reaction is spontaneous. We can test this prediction bymixing the solutions (Figure 9) and performing some diagnostic tests. If the solutions aremixed and the purple colour of the permanganate ion disappears, then it is likely that thepermanganate ion reacted. If the pH of the solution is tested before and after reaction,and the pH has increased, then the hydrogen ions likely reacted.

RA

OA

SRA

OAOAOASOAOA

K�(aq) MnO4�(aq) H�(aq) Fe2�(aq) SO4

2�(aq) H2O(l)

RA

OA

RA

OAOAOAOAOA

K�(aq) MnO4�(aq) H�(aq) Fe2�(aq) SO4

2�(aq) H2O(l)

SAMPLE problem 13.5

Figure 9A solution of potassiumpermanganate is being added to anacidic solution of iron(II) ions. Thedark purple colour of MnO4

�(aq)ions instantly disappears. Theinterpretation is that MnO4

�(aq)ions react with Fe2�(aq) ions toproduce the yellow-brown Fe3�(aq)and Mn2�(aq) ions.

Learning Tip• If the half-reaction equation

shows two or more entitiespresent, then both must be inyour list. If there is only oneentity, then leave it unlabelledas a spectator ion.

• Be careful with a few entitiesthat can act either as an OAor an RA, for example, theiron(II) ion in Sample Problem13.5.



Section 13.2

In a chemical industry, could copper pipe be used to transport a hydrochloric acidsolution? To answer this question,(a) predict the redox reaction and its spontaneity, and(b) describe two diagnostic tests that could be done to test your prediction.

Solution

(a) SOA OA

Cu(s) H�(aq) Cl�(aq) H2O(l)

SRA RA RA RA

2 H�(aq) � 2 e� → H2(g)

Cu(s) → Cu2�(aq) � 2 e�

nonspont.

2 H�(aq) � Cu(s) H2(g) � Cu2�(aq)

Since the reaction is nonspontaneous, it should be possible to use a copper pipe tocarry hydrochloric acid.

(b) If no gas is produced when the mixture is observed, then it is likely that nohydrogen gas was produced (Figure 10). If the colour of the solution did notchange to blue, then copper probably did not react to produce copper(II) ions. (Ifthe solution is tested for pH before and after adding the copper, and the pH did notincrease, then the hydrogen ions probably did not react.)


Figure 10Copper in hydrochloric acid doesnot appear to react.

DisproportionationChemists believe that redox reactions are electron-transfer reactions. One reactant (OA)removes electrons from a second reactant (RA) if a spontaneous reaction takes place.Although the oxidizing and reducing agents that react are usually different entities, thisis not a requirement. A reaction in which a species is both oxidized and reduced is calleddisproportionation. This type of reaction is often a redox reaction and occurs when asubstance can act either as an oxidizing agent or as a reducing agent.

For example, we know that an iron(II) ion can behave either as an oxidizing agent ora reducing agent (Figure 7, page 574). What happens if two iron(II) ions in a solutioncollide? Will a spontaneous reaction occur as a result of an electron transfer from oneiron(II) ion to another iron(II) ion?

Fe2�(aq) � 2 e– → Fe(s)

2[Fe2�(aq) → Fe3�(aq) � e�]

3 Fe2�(aq) → Fe(s) � 2 Fe3�(aq)

Using a redox table and the spontaneity rule with iron(II) as the strongest oxidizingagent and iron(II) as the strongest reducing agent, we see that this reaction is nonspontaneous.

Not all disproportionation reactions are nonspontaneous. For example, check thedisproportionation of the copper(I) ion on the redox table in Appendix I. Copper(I) asan oxidizing agent appears above copper(I) as a reducing agent. Therefore, an aqueoussolution of copper(I) ions will spontaneously, but slowly, disproportionate into copper(II)ions and copper metal.

Learning TipDisproportionation reactionsare also referred to by the more descriptive terms of“self oxidation–reduction” or“autoxidation.”


578 Chapter 13 NEL



Predicting the Reaction of Sodium Metal(Demonstration)The process of developing theories, laws, and generalizationsrequires that they be tested numerous times in as many differentsituations as possible. This process is necessary not only todetermine their validity, but also to identify exceptions that maylead to new knowledge.

As part of the Design, include a list of diagnostic tests usingthe “If [procedure] and [evidence], then [analysis]” format forevery product predicted. (This format is described in AppendixC.4.)

PurposeThe purpose of this demonstration is to test the five-step methodfor predicting redox reactions.

ProblemWhat are the products of the reaction of sodium metal withwater?


Step 1: List all entities present and classify each as a possible oxidizing agent, reducingagent, or both. Do not label spectator ions.

Step 2: Choose the strongest oxidizing agent as indicated in a redox table, and writethe equation for its reduction.

Step 3: Choose the strongest reducing agent as indicated in the table, and write theequation for its oxidation.

Step 4: Balance the number of electrons lost and gained in the half-reaction equationsby multiplying one or both equations by a number. Then add the two balancedhalf-reaction equations to obtain a net ionic equation.

Step 5: Using the spontaneity rule, predict whether the net ionic equation represents a spontaneous or nonspontaneous redox reaction.

SUMMARYFive-Step Method for Predicting RedoxReactions

Will a solution of chromium(II) chloride be stable? Predict the redox reaction and itsspontaneity.

SolutionSOA OA

Cr2�(aq) Cl�(aq) H2O(l)

SRA RA RA RA

Cr2�(aq) �2 e� → Cr(s)

2 [Cr2�(aq) → Cr3�(aq) �e�]

nonspont.

3 Cr2�(aq) Cr(s) � 2 Cr3�(aq)

According to the redox table and the spontaneity rule, a chromium(II) chloride solutionshould not react and, therefore, should be stable.




Predicting Redox Reactions by Constructing Half-ReactionsA redox reaction includes both an oxidation and a reduction. In other words, one substancehas to lose electrons as another substance gains electrons. Choosing and writing half-reaction equations is commonly done using a table of relative strengths of oxidizing andreducing agents. But what if this table does not provide the half-reaction equationsneeded? In this case, use your knowledge about constructing your own half-reactionequations (refer to the Summary on page 567), and then balance electrons to obtain theoverall redox reaction equation.

For a particular reaction, chemists know the main starting materials and the reactionconditions (such as acidic or basic). A chemical analysis of the products determines theoxidized and reduced species produced in the reaction. This provides a skeleton equationshowing only the main reactants and products. Chemists can then determine the detailsof the final redox equation by looking at the individual balanced half-reaction equations.

Section 13.2

Practice26. Use the five-step method to predict the most likely redox reaction in each of the

following situations. For any spontaneous reaction, describe one diagnostic test toidentify a primary product.(a) During a demonstration, zinc metal is placed in a hydrochloric acid solution. (b) A gold ring accidentally falls into a hydrochloric acid solution.(c) Nitric acid is painted onto a copper sheet to etch a design.

27. In your previous chemistry course, you made predictions of reactions according to thesingle replacement generalization assuming the formation of the most common ion.(a) Use the generalization about single replacement reactions to predict the reaction

of iron metal with a copper(II) sulfate solution.(b) Use the redox theory and table to predict the most likely redox reaction of iron

metal with a copper(II) sulfate solution.(c) Can both predictions be correct? Which do you think is most likely correct and

why?(d) Write one qualitative and one quantitative experimental design to test the two

different predictions made for the reaction between iron metal and the copper(II) sulfate solution.

28. Write a Prediction, with your reasoning, and a Design (including safety precautions,diagnostic tests, and disposal instructions) for the following experiment.

ProblemWhat are the products of the reaction of tin(II) chloride with an ammoniumdichromate solution acidified with hydrochloric acid?

29. When aluminium pots are used for cooking, small pits often develop in the metal. Useyour knowledge of redox reactions to explain the formation of these pits. Suggest whythis might be a slow process.

30. Oxygen gas is bubbled into an aqueous solution of iron(II) iodide containing excesshydrochloric acid. Predict all spontaneous reactions in the order in which they willoccur.

Figure 11The space shuttle

Aluminium Oxide CloudsThe solid rocket boosters of thespace shuttle contain the mainreactants ammonium perchlorateand aluminium powder.Ammonium perchlorate is apowerful oxidizing agent andaluminium is a relatively strongreducing agent. Their veryexothermic reaction producesfinely divided aluminium oxide,which forms the billows of whitesmoke you see in Figure 11.

DID YOU KNOW ??


580 Chapter 13 NEL

For reactions that occur in basic solutions, you can follow the procedure outlined inSample Problem 13.6, and then convert to a basic solution. In other words, create the bal-anced redox equation for an acidic solution, and then add OH�(aq) to convert theH�(aq) to water molecules. An example for a basic solution is shown in the followingCommunication Example.

Figure 12A breathalyzer is a technologicaldevice that measures the alcoholcontent in exhaled air based on thecolour change from orange(dichromate ion) to green(chromium(III) ion) when an acidicdichromate ion solution reacts withethanol in a breath sample.

A common example of the application of redox reactions is the technology of abreathalyzer (Figure 12). A person suspected of being intoxicated blows into this deviceand the alcohol in the person’s breath reacts with an acidic dichromate ion solution toproduce acetic acid (ethanoic acid) and aqueous chromium(III) ions. Predict the balancedredox reaction equation.

The information provided tells you the skeleton equation for the major reactants andproducts.

C2H5OH(aq) � Cr2O72�(aq) → CH3COOH(aq) � Cr3�(aq)

The first step is to separate the entities into the start of two half-reaction equations,keeping related entities together.

C2H5OH(aq) → CH3COOH(aq)

Cr2O72�(aq) → Cr3�(aq)

Now you can complete each half-reaction equation using the same procedure you usedin Section 13.1: Balance atoms other than O and H; balance O by adding H2O(l); balance Hby adding H+(aq); and finally balance the charge by adding electrons. For the ethanolhalf-reaction,

H2O(l) � C2H5OH(aq) → CH3COOH(aq) � 4 H�(aq) � 4 e�

Follow the same procedure to construct the second half-reaction equation.

6 e– � 14 H�(aq) � Cr2O72�(aq) → 2 Cr3�(aq) � 7 H2O(l)

Recall that the total number of electrons lost must equal the total number of electronsgained. Using the simplest whole numbers, multiply one or both half-reaction equationsso that the electrons will be balanced. In this example, the simplest solution is a total oftwelve electrons transferred.

3 [H2O(l) � C2H5OH(aq) → CH3COOH(aq) � 4 H�(aq) � 4 e�]

2 [6 e� � 14 H�(aq) � Cr2O72�(aq) → 2 Cr3�(aq) � 7 H2O(l)]

Add the two half-reaction equations. Cancel the electrons and anything else that is thesame on both sides of the equation. Note that it is not unusual to have unequal chemicalamounts of some entities, such as hydrogen ions and water, on the two sides of theequations. The lower amount will always cancel completely. For example, 3 H2O(l) from thefirst half-reaction equation cancels 3 mol out of the 14 H2O(l) in the second equationleaving 11 H2O(l) in the net equation.

3 C2H5OH(aq) � 2 Cr2O72�(aq) � 16 H�(aq) → 3 CH3COOH(aq) � 4 Cr3�(aq) � 11 H2O(l)

Check the final redox equation to make sure that both the symbols and the charges arebalanced.

SAMPLE problem 13.6

CAREER CONNECTION

ConservatorThe career of conservator is highlyvaried, allowing theseprofessionals to tailor their work totheir own special interests.Conservators work with a widerange of objects, including antiquemetals, furniture, and ceramics.The conservator makesrecommendations on how theobjects should be preserved, aswell as performs chemicalanalyses to treat existing damage.Conservators often use chemistryto restore historical artifacts. Find out more about therequirements for this interestingand challenging career.




Section 13.2

Permanganate ions and oxalate ions react in a basic solution to produce carbon dioxideand manganese(IV) oxide.

MnO4�(aq) � C2O4

2�(aq) → CO2(g) � MnO2(s)

Write the balanced redox equation for this reaction.

Solution

2 [3 e� � 4 H�(aq) � MnO4�(aq) → MnO2(s) � 2 H2O(l)]

3 [C2O42�(aq) → 2 CO2(g) � 2 e�]

8 H�(aq) � 2 MnO4�(aq) � 3 C2O4

2�(aq) → 2 MnO2(s) � 4 H2O(l) � 6 CO2(g)

8 OH�(aq) � 8 H�(aq) � 2 MnO4�(aq)� 3 C2O4

2�(aq) →2 MnO2(s) � 4 H2O(l) � 6 CO2(g) � 8 OH�(aq)

4 H2O(l) � 2 MnO4�(aq) � 3 C2O4

2�(aq) → 2 MnO2(s) � 6 CO2(g) � 8 OH�(aq)


Step 1: Use the information provided to start two half-reaction equations.Step 2: Balance each half-reaction equation.Step 3: Multiply each half-reaction equation by simple whole numbers to

balance the electrons lost and gained.Step 4: Add the two half-reaction equations, cancelling the electrons and

anything else that is exactly the same on both sides of the equation.

For basic solutions onlyStep 5: Add OH�(aq) to both sides equal in number to the number of H�(aq)

present.Step 6: Combine H�(aq) and OH�(aq) on the same side to form H2O(l), and cancel

the same number of H2O(l) on both sides.

SUMMARYPredicting Balanced Redox Equationsby Constructing Half-Reactions

Practice31. Balance the following redox equations using the half-reaction method. All reactions

occur in an acidic solution. (a) Zn(s) � NO3

�(aq) → NH4�(aq) � Zn2�(aq)

(b) Cl2(aq) � SO2(g) → Cl�(aq) � SO42�(aq)

32. Balance the following skeleton redox equations using the half-reaction method. Allreactions occur in a basic solution. (a) MnO4

�(aq) � I�(aq) → MnO2(s) � I2(s)(b) CN�(aq) � IO3

�(aq) → CNO�(aq)� I�(aq)(c) OCl�(aq) → Cl�(aq) � ClO3

�(aq)

33. Balance the following redox equation.

KMnO4(aq) � H2S(aq) � H2SO4(aq) → K2SO4(aq) � MnSO4(aq) � S(s)

Learning TipAlways check your final answerby counting the number of eachkind of atom on both sides ofthe equation, and by checkingthe net charge on each side.


582 Chapter 13 NEL

Section 13.2 Questions1. What is the key idea used to explain a redox reaction?

2. Distinguish between oxidation and oxidizing agent in termsof modern theory.

3. Distinguish between reduction and reducing agent in termsof modern theory.

4. Write and label two half-reaction equations to describeeach of the following reactions: (a) Co(s) � Cu(NO3)2(aq) → Cu(s) � Co(NO3)2(aq)(b) Cd(s) � Zn(NO3)2(aq) → Zn(s) � Cd(NO3)2(aq)(c) Br2(l) � 2 KI(aq) → I2(s) � 2 KBr(aq)

5. Using the redox table in Appendix I, predict the spontaneityof each of the reactions shown in 4(a) to (c).

6. Prepare a redox table showing the relative strengths ofoxidizing and reducing agents in Table 7.

7. What is the relative strength of oxidizing and reducingagents for strontium, cerium, nickel, hydrogen, platinum,and their aqueous ions? Use the following information toconstruct a table of relative strengths of oxidizing andreducing agents.

3 Sr(s) � 2 Ce3�(aq) → 3 Sr2�(aq) � 2 Ce(s)

Ni(s) � 2 H�(aq) → Ni2�(aq) � H2(g)

2 Ce3�(aq) � 3 Ni(s) → no evidence of reaction

Pt(s) � 2 H�(aq) → no evidence of reaction

8. Write an experimental design to determine a mini-redoxtable for the first three metals and metal ions of Group 12.Include safety and disposal information.

9. For each of the following mixtures, use the complete five-step method to predict the most likely redox reaction.Include one diagnostic test to test your predicted reaction. (a) Solutions of nickel(II) nitrate and iron(II) chloride are

mixed.(b) Oxygen gas is bubbled over a solid silver mesh

immersed in a solution of sodium iodide.(c) An acidic solution of potassium dichromate is added to

a sodium iodide solution.

10. A chemical technician prepares several solutions for use ina chemical analysis. Will each of the solutions listed belowbe stable if stored for a long time? Justify your answer. (a) acidic tin(II) chloride in an inert glass container(b) copper(II) nitrate in a tin can

11. In the industrial production of iodine, chlorine gas isbubbled into seawater. Using only water and iodide ions inseawater as the possible reactants, predict the most likelyredox reaction, including equations for the half-reactions.

12. The steel of an automobile fender is exposed to acid rain.(Assume that steel is made mainly of iron.) Predict the mostlikely redox reactions, including the equations for therelevant half-reactions.

13. Define disproportionation and illustrate it using half-reaction equations of tin(II) ions.

14. An excess of cobalt metal was left in an aqueous mixturecontaining silver ions, iron(III) ions, and copper(II) ions foran extended time. Write a balanced redox equation forevery reaction that occurs.

15. Balance the following equation representing a reaction thatoccurs in an acidic solution: Mn2�(aq) � HBiO3(aq) → Bi3�(aq) � MnO4

�(aq)

16. Balance the following equations representing reactions thatoccur in a basic solution: (a) Cr(OH)3(s) � IO3

�(aq) → CrO42�(aq) � I�(aq)

(b) Ag2O(s) � CH2O(aq) → Ag(s) � CHO2�(aq)

17. Many commercially available drain cleaners contain a basicsolution of sodium hydroxide, which helps to remove greasein the drains. Some solid drain cleaners contain solidsodium hydroxide and finely divided aluminium metal. Whenmixed with water, this produces a very vigorous, exothermicreaction shown by the following skeleton equation:

Al(s) � H2O(l) → Al(OH)4�(aq) � H2(g)

(a) Complete the balanced redox equation for this reaction.(b) How does this chemical technology provide a solution to

a practical problem?(c) Describe and discuss some possible health and safety

issues associated with the use of solid drain cleaners.

18. What is the WHMIS symbol for oxidizing materials? WhatHousehold Hazardous Product Symbols would be used foroxidizing materials present in a consumer product?

Extension

19. The subject of antioxidants may be controversial, but thereare many groups and companies that want you to useantioxidant products. What does the term antioxidantsuggest? Define antioxidant and list three importantexamples. Briefly summarize some of the controversy that isassociated with antioxidants.

20. Many biologically important molecules contain a metal ionsurrounded by and bonded to a large organic molecule(such as porphyrin). One very important example ischlorophyll. Prepare a report (paper, poster, or electronic)that contains information about the structure of the mainform of chlorophyll, the general role that chlorophyll playsin electron transfer reactions, the close relatives ofchlorophyll that contain iron and cobalt, and atechnological application inspired by the chlorophyllmolecule.

Table 7 Reactions of Group 13 Elements and Ions

Al3�(aq) Tl�(aq) Ga3�(aq) In3�(aq)

Al X ✓ ✓ ✓

Tl X X X X

Ga X ✓ X ✓

In X ✓ X X

X no evidence of a redox reaction ✓ a spontaneous reaction





13.313.3Oxidation StatesHistorically, oxidation and reduction were considered to be separate processes, more ofinterest for technology than for science. With modern atomic theory came the idea ofan electron transfer involving both a gain of electrons by one entity and a loss of electronsby another entity. This theory of redox reactions is most easily understood for atoms ormonatomic ions. Metals and monatomic anions tend to lose electrons (become oxidized),whereas nonmetals and monatomic cations tend to gain electrons (become reduced).

More complex redox reactions, such as the reduction of iron(III) oxide by carbonmonoxide in iron production, the oxidation of glucose in cellular respiration, and the useof dichromate ions in chemical analysis, are not adequately described or explained withsimple redox theory.

To describe the oxidation and reduction of molecules and polyatomic ions, chemistsdeveloped a method of “electron bookkeeping” to keep track of the loss and gain of elec-trons. The idea is similar to determining the electric charge of a simple ion by countingelectrons and protons; for example, a sodium ion (11p�, 10e�) has an ion charge of 1�.For atoms in molecules and polyatomic ions, chemists count a shared pair of electronsin a covalent bond as if it belongs entirely to the more electronegative atom in the bond.For example, in a water molecule (Figure 1), the shared pairs of electrons are assignedto the oxygen atom because oxygen has a higher electronegativity (3.5) compared withhydrogen (2.1). Now we can count the electrons around the oxygen and compare this withthe number of protons in the nucleus (just like we do for simple ions). In this system,the oxidation state of an atom in an entity is defined as the apparent net electric chargethat an atom would have if electron pairs in covalent bonds belonged entirely to themore electronegative atom. An oxidation state is a useful idea for keeping track of elec-trons, but it does not represent an actual charge on an atom—oxidation states are arbi-trary charges and should not be confused with actual electric charges.

An oxidation number is a positive or negative number corresponding to the oxida-tion state assigned to an atom in a covalently bonded entity. For example, in a watermolecule, the oxidation number of the oxygen atom is �2 and the oxidation number ofeach hydrogen atom is �1.

To distinguish oxidation numbers from actual electrical charges, oxidation numbersare written in this textbook as positive or negative numbers; that is, with the sign pre-ceding the number. Chemists use this method to assign oxidation numbers to manycommon atoms and ions (Table 1), which can then be used to determine the oxidationnumbers of other atoms.

O H

HFigure 1An oxygen atom has 8 p� and 8 e�.If the oxygen atom gets to count thetwo hydrogen electrons in the twoshared pairs of electrons, then 8 p�

and 10 e� results in an apparent netcharge of 2�. Each hydrogen atomwith 1 p� has no additionalelectrons. Its one electron hasalready been counted by the oxygenatom. Therefore, the hydrogen hasan apparent net charge of 1� .

Learning TipAlthough the meaning of theterms oxidation state andoxidation number are slightlydifferent, some people usethese terms interchangeably.

Table 1 Common Oxidation Numbers

Atom or ion Oxidation number Examples

all atoms in elements 0 Na is 0Cl in Cl2 is 0

hydrogen in all compounds, �1 H in HCl is �1except hydrogen in hydrides �1 H in LiH is �1

oxygen in all compounds, �2 O in H2O is �2except oxygen in peroxides �1 O in H2O2 is �1

all monatomic ions charge on ion Na� is �1S2� is �2

Learning TipOxidation numbers are simplypositive or negative numbersassigned on the basis of a setof arbitrary rules. It is importantfor you to realize that these arenot electric charges. For thisreason, chemists use the termoxidation number. For example,we assign oxidation numbers of�2 and �1 to the oxygen andhydrogen atoms in a watermolecule.


584 Chapter 13 NEL

What is the oxidation number of carbon in methane, CH4?

This is determined by assigning an oxidation number of �1 to hydrogen (Table 1).

x �1

CH4

Now solve for x. Since a methane molecule is electrically neutral, the oxidation numbersof the one carbon atom and the four hydrogen atoms (4 times �1) must equal zero.

x � 4(�1) � 0

x � �4

x �1 �4 �1

CH4 becomes CH4

Carbon in methane has an oxidation number of �4.

SAMPLE problem 13.7

Learning Tip1. The sum of the oxidation

numbers of any entity mustequal the net charge on thatentity: zero for neutralcompounds, the ion chargefor polyatomic ions.

2. The method only works ifthere is just one unknownafter referring to Table 1. Ifthere are two or moreunknowns, a Lewis formulaand electronegativities arerequired.

What is the oxidation number of manganese in a permanganate ion, MnO4�.

The oxidation number of manganese in the permanganate ion, MnO4�, is determined

using the oxidation number of oxygen as �2 (Table 1) and the knowledge that the chargeon the ion is 1�. The total of the oxidation numbers of the one manganese atom (x) andthe four oxygen atoms (4 times �2) must equal the charge on the ion (1�).

x � 4(�2) � �1

x � �7

x �2 �7 �2

MnO4� becomes MnO4

�

The oxidation number of manganese in MnO4� is �7.

SAMPLE problem 13.8

What is the oxidation number of sulfur in sodium sulfate?

Solution

�1 x �2 2(�1) � x � 4(�2) � 0Na2SO4 x � �6

According to the concept of oxidation states, the oxidation number of sulfur in sodiumsulfate is �6.


Oxidation numbers are simply a systematic way of counting electrons. Therefore, thesum of the oxidation numbers in a compound or ion must equal the total charge—zerofor neutral compounds and the ion charge for ions.

Learning TipAlternatively, you can alwayssplit an ionic formula into thecation and anion before solvingfor an unknown oxidationnumber.

For example, if you want toknow the oxidation number forsulfur in sodium sulfate, startwith the sulfate ion:

x �2

SO42�

x � 4(�2) � �2 x � �6



Section 13.3

Step 1: Assign common oxidation numbers (Table 1 on page 583).Step 2: The total of the oxidation numbers of atoms in a molecule or ion equals the value

of the net electric charge on the molecule or ion.(a) The sum of the oxidation numbers for a compound is zero.(b) The sum of the oxidation numbers for a polyatomic ion equals the charge

on the ion.Step 3: Any unknown oxidation number is determined algebraically from the sum of

the known oxidation numbers and the net charge on the entity.

SUMMARY Determining Oxidation Numbers

Oxidation Numbers and Redox ReactionsAlthough the concept of oxidation states is somewhat arbitrary, because it is based onassigned charges, it is self-consistent and allows predictions of electron transfer. Chemistsbelieve that if the oxidation number of an atom or ion changes during a chemical reac-tion, then an electron transfer (that is, an oxidation–reduction reaction) occurs. Basedon oxidation numbers, an increase in the oxidation number is defined as an oxidationand a decrease in the oxidation number is a reduction. If oxidation numbers are listedas positive and negative numbers on a number line (Figure 3), then the process ofoxidation involves a change to a more positive value (“up” on the number line) andreduction is a change to a more negative value (“down” on the number line). If the oxidation numbers do not change, this is interpreted as no transfer of electrons. A reac-tion in which all oxidation numbers remain the same is not a redox reaction.

Coal is a fossil fuel that is burned in huge quantities in some electrical power gener-ating stations. If we assume pure carbon and complete combustion, carbon is convertedto carbon dioxide. In this reaction, the oxidation number of carbon changes from 0 in

0

�1

�2

+1

+2

oxid

atio

n

reduction

oxidationnumber

Figure 3In a redox reaction, both oxidationand reduction occur.

Figure 2The bleaching of wood pulp toproduce white paper is now mostlydone using chlorine dioxide in placeof the more environmentallydamaging elemental chlorine.

Practice1. Determine the oxidation number of

(a) S in SO2 (c) S in SO42� (e) I in MgI2

(b) Cl in HClO4 (d) Cr in Cr2O72� (f) H in CaH2

2. Determine the oxidation number of nitrogen in (a) N2O(g) (d) NH3(g) (g) N2(g)(b) NO(g) (e) N2H4(g) (h) NH4Cl(s)(c) NO2(g) (f) NaNO3(s)

3. Determine the oxidation number of carbon in (a) graphite (elemental carbon) (c) sodium carbonate(b) glucose (d) carbon monoxide

4. Bruderheim, Alberta, is the site of several companies that produce sodium chlorate.Almost all of the sodium chlorate produced is sold to pulp and paper mills to producechlorine dioxide as a bleaching agent (Figure 2). Determine the oxidation number ofevery atom or ion in the following chemical equation for the industrial production ofchlorine dioxide.

2 ClO3�(aq) � 2 Cl�(aq) � 4 H�(aq) → 2 ClO2(g) � Cl2(g) � 2 H2O(l)

5. Carbon can be progressively oxidized in a series of organic reactions. Determine theoxidation number of carbon in each of the compounds in the following series ofoxidations:

methane→ methanol→ methanal→ methanoic acid→ carbon dioxideCH4 CH3OH CH2O HCOOH CO2


586 Chapter 13 NEL

C(s) to �4 in CO2(g). Simultaneously, oxygen is reduced from 0 in O2(g) to �2 inCO2(g).

The main purpose of assigning oxidation numbers is to see how these numbers changeas a result of a chemical reaction. In any redox reaction, like the combustion of carbon, therewill always be both an oxidation and a reduction. We will use these changes to balanceredox equations, but first we will look at some additional examples.


Oxidation States of VanadiumVanadium is a transition metal that forms many different ions(Table 2). Vanadium and its compounds have many differentuses, including colouring for glass, ceramics, and plastics.

PurposeThe purpose of this lab exercise is to use the concept of oxidationstates to investigate some redox chemistry of vanadium compounds.

ProblemWhat are the oxidation numbers and changes in oxidationnumber for vanadium ions?

Analysis(a) Using Table 2, identify the vanadium ions in the sequence of

reactions in Table 3.(b) In each case, is the vanadium being oxidized or reduced?

Justify your answer, using oxidation numbers.(c) Explain the observations made in (3) to (6) in Table 3.

Suggest what is causing these changes.

LAB EXERCISE 13.B Report Checklist

Table 2 Colours of Vanadium Ions

Ion name Ion formula Colour

vanadate(V) VO3�(aq)

vanadate (IV) VO2�(aq)

vanadium(III) V3�(aq)

vanadium(II) V2�(aq)

EvidenceTable 3 Reactions of Vanadium Ions

Procedure Final solution colours

(1) ammonium vanadate(V)dissolved in sulfuric acid yellow

(2) yellow solution withthree subsequent additions yellow turned blue,of small quantities of zinc dust then green, then violet

(3) violet solution leftsitting in an open container slowly turned green

(4) yellow solution mixed withpotassium iodide solution very dark blue, almost black

(5) blue solution mixed withpotassium iodide solution stayed blue; no change

(6) violet solution slowly mixedwith acidic potassium violet to green to blue to permanganate yellow

C(s) + O2(g) → CO2(g) 0 0 +4 –2

oxidation

reduction



Section 13.3

You have seen the reaction of active metals such as zinc with an acid. Identify theoxidation and reduction in the reaction of zinc metal with hydrochloric acid.

First, you need to write the chemical equation, as it is not provided. Net ionic equationsare best, but the procedure will still work if you write a nonionic equation.

Zn(s) � 2 H�(aq) → Zn2�(aq) � H2(g)

After writing the equation, determine all oxidation numbers.

0 �1 �2 0

Zn(s) � 2 H�(aq) → Zn2�(aq) � H2(g)

Now look for the oxidation number of an atom/ion that increases as a result of thereaction and label the change as oxidation. There must also be an atom/ion whoseoxidation number decreases. Label this change as reduction.

Zn(s) � 2 H�(aq) → Zn2�(aq) + H2(g)

SAMPLE problem 13.9

0 +1 +2 0

oxidation

reduction

When natural gas burns in a furnace, carbon dioxide and water form. Identify oxidationand reduction in this reaction.

First, write the equation.

CH4(g) � 2 O2(g) → CO2(g) � 2 H2O(g)

Now we can insert the oxidation numbers and arrows.

Carbon is oxidized from �4 in methane to �4 in carbon dioxide as it reacts with oxygen.Simultaneously, oxygen is reduced from 0 in oxygen gas to �2 in both products.

Notice that the oxygen atoms in the reactant are distributed between the two products.This does not change our procedure because we are only looking for the change fromreactant to product. We say that “oxygen is reduced” in this reaction and it does notmatter where the reduced oxygen appears in the products.

SAMPLE problem 13.10

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) �4 �1 0 �4 �2 �1 �2

oxidation

reduction

Redox in Biological SystemsBiologists often classify oxidationand reduction in terms of theaddition or removal of oxygen orhydrogen.

• Removal of oxygen decreasesthe oxidation number of carbon(i.e., a reduction), e.g.,

HCOOH → HCHO

C is �2 C is 0

The reverse also applies.Addition of oxygen correspondsto an oxidation.

• Removal of hydrogen increasesthe oxidation number of carbon(i.e., an oxidation), e.g., in thenicotinamide coenzyme

NADH → NAD�

C is �2 C is �1

The reverse also applies.Addition of hydrogencorresponds to a reduction.

DID YOU KNOW ??


588 Chapter 13 NEL

Redox Reactions in LivingOrganismsThe ability of carbon to take ondifferent oxidation states isessential to life on Earth.Photosynthesis involves a series ofreduction reactions in which theoxidation number of carbonchanges from �4 in carbondioxide to an average of 0 insugars such as glucose.

Silicon, which resemblescarbon, has been postulated as afoundation for life elsewhere in theuniverse. The most commonoxidation number for silicon is �4.

DID YOU KNOW ??

Practice6. Methanol reacts with acidic permanganate ions as shown below:

5 CH3OH(l) � 2 MnO4�(aq) � 6 H�(aq) → 5 CH2O(l) � 2 Mn2�(aq) � 8 H2O(l)

(a) Assign oxidation numbers to all atoms/ions.(b) Which atom/ion is oxidized? Label the oxidation above the equation.(c) Which atom/ion is reduced? Label the reduction below the equation.

• According to current theory, a redox reaction is a chemical reaction in which elec-trons are transferred and the oxidation numbers change.

• Oxidation is the increase in oxidation number and corresponds to a loss ofelectrons.

• Reduction is the decrease in oxidation number and corresponds to a gain ofelectrons.

SUMMARY Electron Transfer and Oxidation States

The determination of blood alcohol content from a sample of breath (Figure 12, page 580)or blood involves the reaction of the sample with acidic potassium dichromate solution. Ifethanol is present, chromium(III) ions, water, and acetic acid are produced. Identify theoxidation and reduction in the following chemical reaction:

2 Cr2O72�(aq) � 16 H�(aq) � 3 C2H5OH(aq) → 4 Cr3�(aq) � 11 H2O(g) � 3 CH3COOH(aq)

Solution

According to the concept of oxidation states, chromium atoms in Cr2O72� are reduced

(�6 to �3). Carbon atoms in C2H5OH are oxidized (�2 to 0).


2 Cr2O7 (aq) + 16 H (aq) + 3 C2H5OH(aq) → 4 Cr (aq) + 11 H2O(g) + 3 CH3COOH(aq) +6 –2

oxidation

reduction

2– + 3+

+3+1 –2 +1 –2 +1 +1 –2 0 +1 0–2–2 +1

WEB Activity

Case Study—Catalytic ConvertersThe catalytic converter is a technological development designed to reduce the amount ofpollutants in an automobile’s exhaust. Describe the main components of the converter andexplain the three main reactions. Use oxidation numbers in your explanation.




Section 13.3

Balancing Redox Equations Using OxidationNumbersKnowing the ratio of reacting chemicals is necessary in many applications. Chemistsuse the mole ratio from a balanced chemical equation to study the nature of the reac-tion. The stoichiometry of a reaction is essential in many types of chemical analysis,such as the breathalyzer. Finally, chemical industries need to know the quantities of reac-tants to mix and the yield of a desired product. In this and previous courses, you havealready seen many examples of the use of balanced chemical equations.

Simple redox reaction equations can be balanced by inspection or by a trial-and-errormethod. You have done this often in previous courses for reactions such as single replace-ment and combustion. More complex redox reactions may be very difficult to balancethis way because of the number and complexity of the reactants and products. As youwill see, oxidation numbers and half-reaction equations can be used to balance anyredox equation.

One way of recognizing a redox reaction is to assign oxidation numbers to each atomor ion, and then look for any changes in the numbers. Any change in the oxidation numberof a particular atom or ion is believed to be related to a change in the number of electrons.Because electrons are transferred in a redox reaction, the total number of electrons lostby one atom/ion must equal the total number of electrons gained by another atom/ion.In terms of oxidation numbers, this means that the changes in oxidation numbers mustalso be balanced.

The total increase in oxidation number for a particular atom/ion must equal the totaldecrease in oxidation number of another atom/ion.

Let’s look at a simple example first. You could easily balance this equation by inspec-tion, but we will use it to illustrate the main points of the oxidation number method.

Learning TipAll redox reactions are electrontransfer reactions. This meansthat the electrons that are lostby one entity are the sameelectrons that are gained byanother. This is like you givingfive dollars to a friend. You losefive dollars and your friendgains five dollars: five dollarshas been transferred. Obviously,the money lost must equal themoney gained. The same is truefor electrons in a redox reaction.

7. For each of the following chemical reactions, assign oxidation numbers to eachatom/ion and indicate whether the equation represents a redox reaction. If it does,identify the oxidation and reduction. (a) Cu(s) � 2 AgNO3(aq) → 2 Ag(s) � Cu(NO3)2(aq)(b) Pb(NO3)2(aq) � 2 KI(aq) → PbI2(s) � 2 KNO3(aq)(c) Cl2(aq) � 2 KI(aq) → I2(s) � 2 KCl(aq)(d) 2 NaCl(l) → 2 Na(l) � Cl2(g)(e) HCl(aq) � NaOH(aq) → H2O(l) � NaCl(aq)(f) 2 Al(s) � 3 Cl2(g) → 2 AlCl3(s)(g) 2 C4H10(g) � 13 O2(g) → 8 CO2(g) � 10 H2O(l)(h) 2 H2O2(l) → 2 H2O(l) � O2(g)

8. Where possible, classify the chemical equations in question 7 (a�h) using the fivereaction types: formation, simple decomposition, single replacement, doublereplacement, and complete combustion. Which reaction type does not appear to be aredox reaction?

9. Hydrogen peroxide, H2O2(l) , can either be oxidized or reduced depending on thesubstance with which it reacts. Use oxidation numbers to explain why this is possible.

10. Earth has an oxidizing atmosphere of oxygen. The planet Saturn has a reducingatmosphere of hydrogen and methane. Describe the two types of atmospheres interms of changes in oxidation numbers of carbon and of the likely reactions.


590 Chapter 13 NEL

Learning TipYou can adjust the number ofelectrons per atom to thenumber per molecule bymultiplying the number peratom by the subscript of theatom in the chemical formula.

Hydrogen sulfide is an unpleasant constituent of “sour” natural gas. Hydrogen sulfide isnot only very toxic, but it also smells terrible, similar to rotting eggs. It is common practiceto “flare,” or burn, relatively small quantities of sour natural gas that occur with oil deposits(Figure 4). The gas is burned because it is not worth recovering and treating a smallquantity of gas. When this gas is burned, hydrogen sulfide is converted to sulfur dioxide.Use oxidation numbers to balance this equation.

H2S(g) � O2(g) → SO2(g) � H2O(g)

The first step is to assign oxidation numbers to all atoms/ions and look for the numbersthat change. Circle or highlight the oxidation numbers that change.

Notice that a sulfur atom is oxidized from �2 to �4. This is a change of 6 and means 6 e�

have been transferred. An oxygen atom is reduced from 0 to �2, a change of 2 or 2 e�

transferred. Because the substances in the equation are molecules, not atoms, we need tospecify the change in the number of electrons per molecule.

�1 �2 0 �4 �2 �1 �2

H2S(g) � O2(g) → SO2(g) � H2O(g)

6 e�/S atom 2 e�/O atom

6 e�/H2S 4 e�/O2

One H2S molecule contains one sulfur atom. Therefore, the number of electronstransferred per sulfur atom is the same number per H2S molecule. An O2 moleculecontains two O atoms. Therefore, when one O2 molecule reacts, two oxygen atomstransfer 2 e� each for a total of 4 e�.

The next step is to determine the simplest whole numbers that will balance the numberof electrons transferred for each reactant. The numbers become the coefficients for thereactants.

�1 �2 0 �4 �2 �1 �2

H2S(g) � O2(g) → SO2(g) � H2O(g)

6 e�/S atom 2 e�/O atom

6 e�/H2S 4 e�/O2

Now you have the coefficients for the reactants.

2 H2S(g) � 3 O2(g) → SO2(g) � H2O(g)

The coefficients of the products can easily be obtained by balancing the atoms whoseoxidation numbers have changed, and then any other atoms. The final balanced equationis shown below:

2 H2S(g) � 3 O2(g) → 2 SO2(g) � 2 H2O(g)

� 3�12

� 2�12


�1�2

H2S(g) + O2(g) SO2(g) + H2O(g)�4 �2 �1 �20

oxidation

reduction

Figure 4Since sour natural gas oftencontains hydrogen sulfide, burningit can be a significant source ofpollutants. It is also a waste ofenergy when many of these flaresoperate in a large oil field.



Section 13.3

Chlorate ions and iodine react in an acidic solution to produce chloride ions and iodateions. Balance the equation for this reaction.

ClO3�(aq) � I2(aq) → Cl�(aq) � IO3

�(aq)

Assign oxidation numbers to each atom/ion and note which numbers change.

�5 �2 0 �1 �5 �2


�(aq)

A chlorine atom is reduced from �5 to �1, a change of 6. Simultaneously, an iodine atomis oxidized from 0 to �5, a change of 5. Record the change in the number of electrons peratom, and per molecule or polyatomic ion.

�5 �2 0 �1 �5 �2


�(aq)

6 e�/Cl 5 e�/I

6 e�/ClO3� 10 e�/I2

The total number of electrons transferred by each reactant must be the same. Multiplythe numbers of electrons by the simplest whole numbers to make the totals equal, in thiscase, 30 e�. You can now write the coefficients for the reactants and the products.

�5 �2 0 �1 �5 �2

5 ClO3�(aq) � 3 I2(aq) → 5 Cl�(aq) � 6 IO3

�(aq)

6 e�/Cl 5 e�/I

6 e�/ClO3� 10 e�/I2

� 5 � 3

Although the chlorine and iodine atoms are now balanced, notice that the oxygen atomsare not; 15 on the left versus 18 on the right. Because this reaction occurs in an aqueoussolution, we can add H2O molecules to balance the O atoms. The reactant side requiresthree oxygen atoms (from three water molecules) to equal the total of 18 oxygen atoms onthe product side.

3 H2O(l) � 5 ClO3�(aq) � 3 I2(aq) → 5 Cl�(aq) � 6 IO3

�(aq)

In adding water molecules, we are also adding H atoms. Because this reaction occurs inan acidic solution, we will add H�(aq) to balance the hydrogen.

3 H2O(l) � 5 ClO3�(aq) � 3 I2(aq) → 5 Cl�(aq) � 6 IO3

�(aq) � 6 H�(aq)

The redox equation should now be completely balanced. Check your work by checking thetotal numbers of each atom/ion on each side and checking the total electric charge,which should also be balanced.


Sometimes you may not know all of the reactants and products of a redox reaction.The main reactants and oxidized/reduced products will always be given and you will knowif the reaction took place in an acidic or basic solution. Experimental evidence shows thatwater molecules, hydrogen ions, and hydroxide ions play important roles in reactions insuch solutions. The procedure for balancing such equations is initially the same as the oneused in Sample Problem 13.11, but you will need to add water molecules, hydrogen ions,and/or hydroxide ions to finish the balancing of the overall equation. The following twosample problems illustrate this procedure.

Figure 5Pemmican

Learning TipA balanced chemical reactionequation includes both a massand charge balance. Mass isbalanced using the atomicsymbols. If the symbols balance,but not the charge, the equationis not balanced. Be sure tocheck both the symbols andcharges.

Aboriginal FoodPreservation TechnologyPemmican, from the Cree word“pimikan,” is a mixture of dried,ground red meat, dried berries,and fat or grease from bonemarrow (Figure 5). When storedin skin or intestine bags(essentially becoming vacuum-sealed), the pemmican can bekept unspoiled for months or evenyears. This Aboriginal technologysolved a problem that Westernscientists recognize today as theoxidation of lipids and proteins byaerobic bacteria, the main culpritin the spoilage of meat.

DID YOU KNOW ??


592 Chapter 13 NEL

Methanol reacts with permanganate ions in a basic solution. The main reactants andproducts are shown below. Balance the equation for this reaction.

CH3OH(aq) � MnO4�(aq) → CO3

2�(aq) � MnO42�(aq)

We will follow the same procedure as in the previous problem, adjusting for a basicsolution at the end: assign oxidation numbers; note which ones change and by how muchper reactant; and then balance the total number of electrons to obtain the coefficients forthe main reactants and products.

�2 �1 �2 �1 � 7 �2 � 4 �2 �6 �2

1 CH3OH(aq) � 6 MnO4�(aq) → 1 CO3

2�(aq) � 6 MnO42�(aq)

6 e�/C 1 e�/Mn

6 e�/CH3OH 1 e�/MnO4�

� 1 � 6

Just as before, add H2O(l) to balance the O atoms. The reactant side requires twooxygen atoms (from two water molecules) to equal the 27 oxygen atoms on the productside. Next, balance the H atoms using H�(aq). The product side requires eight hydrogenions to balance the eight hydrogen atoms on the reactant side (four in water and four inmethanol).

2 H2O(l) � CH3OH(aq) � 6 MnO4�(aq) → CO3

2�(aq) � 6 MnO42�(aq) � 8 H�(aq)

If this reaction occurred in an acidic solution, you would now be finished. For a basicsolution, however, we must add enough OH�(aq) to both sides to equal the number ofH�(aq) present. The hydrogen and hydroxide ions on the same side of the equation arethen combined to form water.

8 OH�(aq) � 2 H2O(l) � CH3OH(aq) � 6 MnO4�(aq) →

CO32�(aq) � 6 MnO4

2�(aq) � 8 H�(aq) � 8 OH�(aq)

8 H20

Finally, cancel the same number of H2O molecules on both sides. In this case, the H2Oon the reactant side can be cancelled by also removing 2 H2O from the product side,leaving the extra 6 H2O in the final equation.

8 OH�(aq) � CH3OH(aq) � 6 MnO4�(aq) → CO3

2�(aq) � 6 MnO42�(aq) � 6 H2O(l)


Learning TipOnce the main reactants andproducts are balanced, analternate procedure is to:• balance the charge using

OH�(aq)• balance the hydrogen using

H2O(l)• use the oxygen count as a

check on your work

Learning TipFor disproportionationreactions, start with twoidentical entities on thereactant side and follow theusual procedure for balancingequations.

Household bleach contains sodium hypochlorite. Some of the hypochlorite ionsdisproportionate (react with themselves) to produce chloride ions and chlorate ions. Writethe balanced redox equation for the disproportionation.

Solution

�1 �2 �1 �2 �1 �5 �2

2 ClO�(aq) � ClO�(aq) → 2 Cl�(aq) � ClO3�(aq)

2 e�/Cl 4 e�/Cl

2 e�/ClO� 4 e�/ClO�

� 2 � 1

3 ClO�(aq) → 2 Cl�(aq) � ClO3�(aq)




Section 13.3

Step 1: Assign oxidation numbers and identify the atoms/ions whose oxidation num-bers change.

Step 2: Using the change in oxidation numbers, write the number of electrons trans-ferred per atom.

Step 3: Using the chemical formulas, determine the number of electrons transferred per reactant. (Use the formula subscripts to do this.)

Step 4: Calculate the simplest whole number coefficients for the reactants that will bal-ance the total number of electrons transferred. Balance the reactants and products.

Step 5: Balance the O atoms using H2O(l), and then balance the H atoms using H�(aq).

For basic solutions onlyStep 6: Add OH�(aq) to both sides equal in number to the number of H�(aq) present.Step 7: Combine H�(aq) and OH�(aq) on the same side to form H2O(l), and cancel

the same number of H2O(l) on both sides.

SUMMARYBalancing Redox Equations UsingOxidation Numbers

Balance the chemical equation for the oxidation of ethanol by dichromate ions in abreathalyzer (Figure 6) to form chromium(III) ions and acetic acid in an acidic solution.

Solution

�6 �2 �2�1 �2 �1 �3 0 �1 0 �2�2 �1

16 H�(aq) � 2 Cr2O72�(aq) � 3 C2H5OH(aq) → 4 Cr3�(aq) � 3 CH3COOH(aq) � 11 H2O(l)

3 e�/Cr 2 e�/C

6 e�/Cr2O72� 4 e�/C2H5OH

� 2 � 3


Practice11. Why is the change in oxidation number of an atom the same as the number of

electrons transferred?

12. Balance the following chemical equations for reactions in an acidic solution: (a) Cr2O7

2�(aq) � Cl�(aq) → Cr3�(aq) � Cl2(aq)(b) IO3

�(aq) � HSO3�(aq) → SO4

2�(aq)� I2(s)(c) HBr(aq) � H2SO4(aq) → SO2(g) � Br2(l)

13. Balance the following chemical equations for reactions in a basic solution: (a) MnO4

�(aq) � SO32�(aq) → SO4

2�(aq) � MnO2(s)(b) ClO3

�(aq) � N2H4(aq) → NO(g) � Cl�(aq)

14. For the reactions in question 13, identify the oxidizing agents and the reducingagents.

15. Ammonia gas undergoes a combustion to produce nitrogen dioxide gas and watervapour. Write and balance the reaction equation.

Figure 6The original breath alcohol testingdevice uses the reaction of ethanolwith acidic dichromate solution.More modern devices, like the oneshown, use the reaction of ethanolin a fuel cell to determine the bloodalcohol count.

Learning TipCompare the solution inCommunication Example 4 withthe solution in Sample Problem13.6 (page 580). Theserepresent different approachesto the same question. In thissection, you should practise theoxidation number method. Ontests, you will usually chooseyour method.


594 Chapter 13 NEL

Bleaching Wood PulpThe production of pulp and paper is one of Canada’s majorindustries (Figure 7), employing several thousand peopleacross the country and contributing billions of dollars toCanada’s export market. The processes used in the pulp andpaper industry illustrate how a successful technology can haveboth intended and unintended consequences.

Modern chemistry and chemical technology make itpossible for cellulose fibres from wood pulp to be bleached,dyed, coated, and treated to manufacture paper, as well ascellophane and explosives. The production of white paperinvolves a bleaching process in which a strong oxidizing agentoxidizes coloured organic compounds. For many years, theoxidizing agent was elemental chlorine, which breaks downand removes lignin, an organic polymer that binds the woodfibres together. Over 300 reaction by-products, includingchloroform, carbon tetrachloride, chlorophenols, and furans,are produced during the process.

Many of these by-products are potentially harmful.Research indicates that, although only one of 75 dioxinisomers is extremely toxic, some dioxins can cause immunesystem suppression and severe reproductive disorders,including birth defects and sterility. Also, certain dioxins arepotent carcinogens. Once in the ecosystem, dioxins resistbreakdown and bioaccumulate in animal tissues. Traces ofdioxins have been found in bleached paper products such asdiapers, sanitary products, paper plates, toilet paper, coffeefilters, food packaging, and writing paper.

In the 1980s, growing concern over toxic chemicals enteringthe ecosystem from pulp mill effluent prompted the Canadiangovernment to implement stricter environmental regulations.As a result, the pulp and paper industry underwent the largestenvironmental upgrade in its history. The use of elementalchlorine for bleaching was reduced or replaced, whichproduced a corresponding reduction in the emission of dioxinsand furans. In addition, the recycling of paper increased andthe water consumption per tonne of paper decreased. Overall,the effluent quality has vastly improved, although its adverseeffects on fish and other aquatic organisms are still beingseen.

Industry researchers proposed multiple solutions, many ofwhich were implemented to reduce the quantities of toxicchemicals produced by pulp mills. Each solution involveddifferent designs, materials, and processes. Some of themeasures introduced to reduce the emission oforganochlorines include: more efficient washing of pulp;oxygen pre-bleaching; using chlorine dioxide, hydrogenperoxide, or ozone instead of elemental chlorine; andbleaching to a lesser degree. While these new technologiesare reducing the negative impact of pulp mills on theenvironment, there is concern that they, too, will have someunintended effects on the environment.



1. What were the unintended consequences of usingchlorine to bleach wood pulp?

2. List some new processes that were developed to reducethe problems created by bleaching wood pulp with chlorine.

3. Describe the perspectives that need to be consideredwhen selecting the chemical to replace elemental chlorinein the bleaching process.

4. The most common alternative to the use of elementalchlorine as a bleaching agent is chlorine dioxide gas. Inaddition to having a less negative environmental impact,chlorine dioxide has a greater oxidizing ability. Assumingboth chlorine and chlorine dioxide are converted tochloride ions, use oxidation numbers to show the greateroxidizing ability of chlorine dioxide.

Extension

5. Investigate at least two pulp and paper companies (one inCanada and one outside of North America) to see whatactions they have taken to reduce their impact on theenvironment. Present your findings in a medium of yourchoice.

6. Research several reactions involved in the bleaching ofpaper, and balance or verify the balancing of the reactionequations using oxidation numbers.

Figure 7There are over 150 pulp and paper mills in Canada. Canada isthe fourth largest producer in the world of pulp and paperproducts, and the world's largest supplier of newsprint.





Section 13.3

Section 13.3 Questions1. Copy and complete the following table to distinguish

between oxidation and reduction:

2. Define an oxidation number.

3. State two ways in which you can recognize a redoxreaction, using a chemical reaction equation.

4. Write the oxidation number of each atom/ion in thefollowing substances: (a) carbon monoxide, CO(g) , a toxic gas(b) ozone, O3(g) , ozone layer(c) ammonium chloride, NH4Cl(s), used in dry cells

(batteries)(d) phosphoric acid, H3PO4(aq), in cola soft drinks(e) sodium thiosulfate, Na2S2O3(s), antidote for cyanide

poisoning(f) sodium tripolyphosphate, Na5P3O10(s), in laundry

detergents

5. Assigning oxidation numbers using the rules we haveestablished may occasionally produce some unusualresults. For example, consider Fe3O4. (a) Determine the oxidation number of iron in Fe3O4.(b) What is unusual about your answer? Suggest a reason

for your answer.

6. Redox reactions are common in organic chemistry. Forexample, carboxyl groups can be oxidized to form carbondioxide. In the following chemical equation, permanganateions convert oxalate ions to carbon dioxide in an acidicsolution.

2 MnO4�(aq) � 5 C2O4

2�(aq) � 16 H�(aq) →2 Mn2�(aq) � 8 H2O(l) � 10 CO2(g)

(a) Assign oxidation numbers to all atoms/ions.(b) Which atom is oxidized? State the change.(c) Which atom is reduced? State the change.(d) Identify the oxidizing and reducing agents.

7. When carbon dioxide is released into the atmosphere fromnatural or human activities, some of it reacts with water toform carbonic acid. This accounts for the natural acidity ofrainwater and may also contribute to acid rain.(a) Write the balanced chemical equation for the reaction

of carbon dioxide with water to form carbonic acid.(b) Is this a redox reaction? Justify your answer.

8. Balance the following equations representing reactions thatoccur in an acidic solution: (a) Cu(s) � NO3

�(aq) → Cu2�(aq) � NO2(g)(b) H2O2(aq) � Cr2O7

2�(aq) → Cr3�(aq)� O2(g) � H2O(l)(c) Mn2�(aq) � HBiO3(aq) → Bi3�(aq) � MnO4

�(aq)

9. Balance the following equations representing a reactionthat occurs in a basic solution: (a) Cr(OH)3(s) � IO3

�(aq) → CrO42�(aq) � I�(aq)

(b) Ag2O(s) � HCHO(aq) → Ag(s) � HCO2�(aq)

(c) S2O42�(aq)� O2(g) → SO4

2�(aq)

10. Hydrogen peroxide decomposes in the presence of acatalyst to form water and oxygen gas. Use oxidationnumbers to show that this reaction is a disproportionation.

11. State two general experimental designs that could helpdetermine the balancing of the main species in a redoxreaction.

12. Evidence shows that iron(III) sulfide reacts according to thefollowing balanced redox equation:

2 Fe2S3(s) � 6 H2O(l) � 11 O2(g) →4 FeO(s) � 6 H2SO4(aq)

(a) List the oxidation numbers for all entities in thisreaction equation.

(b) Identify the entities that are oxidized and reduced.(c) How does this reaction equation illustrate the

limitations of the method presented to balance redoxreaction equations?

Extension

13. How is the making of pemmican by First Nations peoplesan example of technology providing solutions to practicalproblems? List some advantages of pemmican and state itsimportance to early European explorers and settlers.

14. Most organisms derive their metabolic energy from cellularrespiration, making this one of the most important biologicalredox processes. Outline the chemistry of cellular respiration.Your response should include:• the overall chemical reaction equation for cellular

respiration, with the oxidation, reduction, oxidizing agent,reducing agent, and overall direction of electron transferindicated

• brief descriptions of the three main stages of aerobicrespiration

15. Science terms and concepts are often used to helppromote a variety of new technologies marketed toconsumers. One recent example is the titanium necklace orbracelet. Refer to Appendix B.4 and use the Internet toanswer the following questions.(a) List some science terms and concepts that are

mentioned in the promotion of this product.(b) Briefly summarize the claims implied by the

manufacturer.(c) What kind of evidence is presented to justify the

claims?(d) Write a brief experimental design to conduct a scientific

test of the claims and collect more reliable evidence.

Electron transfer Oxidation states

oxidation

reduction





596 Chapter 13 NEL

13.413.4 Redox StoichiometryThe stoichiometric method can be used to predict or analyze the quantity of a chemicalinvolved in a chemical reaction. You encountered many applications of stoichiometry inChapter 7 involving masses, volumes, and concentrations of reactants and products.For the stoichiometry calculations in Chapter 7, you assumed that all the reactions werespontaneous, fast, stoichiometric, and quantitative. These same assumptions apply toredox stoichiometry.

There are many industrial and laboratory applications of redox stoichiometry. Forexample, a mining engineer must know the concentration of iron in a sample of iron orein order to decide whether or not a mine would be profitable. Chemical technicians inindustry, monitoring the quality of their companies’ products, must determine the con-centration of substances such as sodium hypochlorite (NaClO) in bleach, or hydrogenperoxide (H2O2) in disinfectants. Hospital laboratory technicians and environmentalchemists detect tiny traces of chemicals by a variety of methods. Although much analyt-ical chemistry involves sophisticated equipment, the basic technological process of titra-tion still has an important role (Appendix C.4).

In a titration, one reagent (the titrant) is slowly added to another (the sample) untilan abrupt change in a solution property (the endpoint) occurs (Figure 1). In acid–basetitrations, the titrant is generally a strong acid or base. In redox titrations, the titrant isalways a strong oxidizing or reducing agent. Two oxidizing agents commonly used inredox titrations are acidic solutions of permanganate ions or dichromate ions. They areboth strong oxidizing agents and undergo a colour change when they oxidize a reducingagent in a sample being titrated. The permanganate ion, which has an intense purple-pink colour in solution, changes to the essentially colourless manganese(II) ion in areaction with a reducing agent that is usually colourless (Figure 2).

MnO4–(aq) � 8 H �(aq) � 5 e– → Mn2�(aq) � 4 H2O(l)

purple-pink colourless

Once the reducing agent in the sample has completely reacted, the next drop of per-manganate added remains unreacted and causes a pink colour in the mixture. The colourchange of the reaction mixture (colourless to pink) is the endpoint and corresponds toa slight excess of unreacted permanganate ion. The volume of permanganate solutionadded when the endpoint is reached is a measurement of the point at which stoichiometricquantities of reactants have been combined.

The dichromate ion is also commonly used in redox titrations; however, its colourchange is not as easy to see: the orange dichromate solution gradually changes to a greenchromium(III) solution. A redox indicator is usually added to produce a sharper endpoint.

Cr2O72–(aq) � 14 H�(aq) � 6 e– → 2 Cr3�(aq) � 7 H2O(l)

orange green

When a titration is used to analyze the concentration of a sample, the concentrationof the titrant used must be accurately known. If the titrant is not a standard solution, thetitrant is standardized by calculating its concentration using evidence from an analysiswith a primary standard. A primary standard is a chemical that can be used directly to prepare a standard solution—a solution of precisely known concentration (refer toUnit 4).

Figure 1Titration is a common experimentaldesign for quantitative chemicalanalysis.

Figure 2A solution of potassiumpermanganate is being added to anacidic solution of iron(II) ions. Thedark purple-pink colour ofMnO4

�(aq) ions instantlydisappears as they react with iron(II)ions to produce the almostcolourless Mn2�(aq).



Section 13.4

Learning TipThe general stoichiometryprocedure is as follows:1. Write a balanced chemical

equation withmeasurements andconversion factors.

2. Convert givenmeasurements into achemical amount.

3. Calculate the amount of therequired substance usingthe mole ratio.

4. Convert this calculatedamount to the finalrequested quantity.

Scientific CredibilityFor credibility, scientific claimsmust be testable empirically. Theresults of any tests must bereplicated by furtherexperimentation. The same personmight repeat the measurements, ina titration, for example, ormembers of the same researchteam might repeat the experiment.The scientific community acceptsnew scientific discoveries only ifdifferent scientists in differentlaboratories are able to reproducethe results.

DID YOU KNOW ??

Standardized SolutionThe standardized potassiumpermanganate solution can beused as a strong oxidizing agent in further titrations. A laboratorytechnician might standardize thesolution in the morning, and thenre-standardize at noon and at the end of the day to increase thecertainty of the results.

DID YOU KNOW ??

A solution of potassium permanganate cannot be directly prepared with a preciselyknown concentration because the permanganate ion reacts with organic and inorganicimpurities in the water and with the water itself. Thus, potassium permanganate is notused as a primary standard. Complete the Analysis of the investigation report.

ProblemWhat is the concentration of the potassium permanganate solution?

DesignA freshly prepared solution of potassium permanganate is titrated against samples ofacidic tin(II) chloride solution, which has a known concentration. The tin(II) chloridesolution is the primary standard.

Evidence

AnalysisThe Analysis for a titration experiment follows the same general stoichiometry steps thatyou practiced in Unit 4. The main difference is that you have a much more sophisticatedway of writing the chemical reaction equation.

The first endpoint was overshot and was not used in the average for the analysis. At theendpoint, an average of 16.8 mL of permanganate solution was used.

OA SOA OA OA OA

K�(aq) MnO4�(aq) H�(aq) Sn2�(aq) Cl�(aq) H2O(l)

SRA RA RA RA

2 [MnO4�(aq) � 8 H�(aq) � 5 e� → Mn2�(aq) � 4 H2O(l)

5 [Sn2�(aq) → Sn4�(aq) � 2 e�]

2 MnO4�(aq) � 16 H�(aq) � 5 Sn2�(aq) → 2 Mn2�(aq) � 8 H2O(l) � 5 Sn4�(aq)

16.8 mL 10.00 mL

c 0.0500 mol/L

nSn2� � 10.00 mL � � 0.500 mmol

nMnO4� � 0.500 mmol � �

25

� � 0.200 mmol

[MnO4�] � � 0.0119 mol/L or 11.9 mmol/L

or [MnO4�] � 10.00 m�L Sn2� � � �

25m

mol

oMl S

nnO2�

4�

� ��16.8 m�L

1MnO4

��

� 0.0119 mol/L

According to the evidence gathered and the stoichiometric analysis, the amountconcentration of the potassium permanganate solution is 0.0119 mol/L or 11.9 mmol/L.

0.0500 mol Sn2�

��1 L Sn2�

0.200 mmol��

16.8 mL

0.0500 mol��

1 L


Table 1 Titration of 10.00 mL of Acidic 0.0500 mol/L SnCl2(aq) with KMnO4(aq)

Trial 1 2 3 4

final burette reading (mL) 18.4 35.3 17.3 34.1

initial burette reading (mL) 1.0 18.4 0.6 17.3

volume of KMnO4(aq) (mL) 17.4 16.9 16.7 16.8

endpoint colour dark pink light pink light pink light pink


598 Chapter 13 NEL


Analyzing for TinExtensive long-term research has found that treating children’steeth with fluoride significantly reduces tooth decay. When thiswas first discovered, toothpastes were produced containing tin(II)fluoride. Complete the Analysis of the investigation report.

PurposeThe purpose of this lab exercise is to use the stoichiometricmethod in a redox chemical analysis.

ProblemWhat is the amount concentration of tin(II) ions in a solutionprepared for research on toothpaste?

LAB EXERCISE 13.C Report Checklist

Evidence

Table 2 Titration of 10.00 mL of acidic Sn2�(aq) with 0.0832 mol/L KMnO4(aq)

Trial 1 2 3

final burette reading (mL) 15.8 28.1 40.6

initial burette reading (mL) 3.4 15.8 28.1

volume of KMnO4(aq) (mL) 12.4 12.3 12.5

WEB Activity

Canadian Achievers—Imants LauksDr. Imants Lauks (Figure 3) is a world leader in developing biochips for clinical diagnosticproducts. He invented the silicon chip blood analyzer in 1986. Biochips combine silicon chiptechnology with chemical reactions, many of which are electrochemical (redox) reactions.

1. Describe the FlexCard™ technology that Dr. Lauks’ current company, Epocal, is developing.

2. What practical problems does this technology solve?Figure 3Imants Lauks (1952– ) www.science.nelson.com GO

Practice1. Titration is a common experimental procedure for the quantitative analysis of

chemical substances. What are the four requirements for titration experiments?

2. Titration is one of several experimental procedures that can be used to determine thequantity of a chemical in a sample. What are some alternative designs available forthis purpose?

3. Silver metal can be recycled by reacting nickel metal with waste silver ion solutions.What volume of 0.10 mol/L silver ion solution will react completely with 25.0 g ofnickel metal?

4. In a chemical analysis of a chromium alloy, all of the chromium is first converted tochromate ions. A 50.0 mL sample of the chromate ion solution is then reduced in a basicsolution to chromium(III) hydroxide by reaction with 22.6 mL of 1.08 mol/L sodium sulfite.In this reaction, the sulfite ions are oxidized to sulfate ions. What is the amountconcentration of the chromate ion solution?

5. Pure iron metal may be used as a primary standard for permanganate solutions. A1.08 g sample of pure iron wire was dissolved in acid, converted to iron(II) ions, anddiluted to 250.0 mL. In the titration, an average volume of 13.6 mL of permanganatesolution was required to react with 10.0 mL of the acidic iron(II) solution. Calculatethe amount concentration of the permanganate solution.



Section 13.4

Figure 4The blue Cr2�(aq) solution is oxidizedto a green Cr3�(aq) solution.


Analyzing for Chromium in SteelStainless steel is a corrosion-resistant, esthetically pleasing alloy,normally composed of nickel, chromium, and iron. Complete theAnalysis of the investigation report.

PurposeThe purpose of this lab exercise is to use the stoichiometricmethod in a redox chemical analysis.

ProblemWhat is the amount concentration of chromium(II) ions in a solution obtained in the analysis of a stainless steel alloy?

DesignA standard potassium dichromate solution is used as an oxidizingagent to oxidize chromium(II) ions to chromium(III) ions in anacidic solution (Figure 4).

LAB EXERCISE 13.D Report Checklist

Evidence

Table 3 Titration of 10.00 mL of acidic Cr2�(aq) with 0.125 mol/L K2Cr2O7(aq)

Trial 1 2 3



volume of K2Cr2O7(aq) (mL) 17.4 17.4 17.4



Analyzing a Hydrogen Peroxide SolutionIn this investigation, you assume the role of a laboratorytechnician working in a consumer advocacy laboratory, testingthe concentration of a hydrogen peroxide solution.

PurposeThe technological purpose of this investigation is to test andevaluate the percent concentration of the consumer solution ofhydrogen peroxide.

ProblemWhat is the percent concentration of hydrogen peroxide in aconsumer product?

DesignAn acidic solution of the primary standard, iron(II) ammoniumsulfate–water (1/6), is prepared and the potassium permanganate

solution is standardized by a titration with this primary standard.A 25.0 mL sample of a consumer solution of hydrogen peroxide isdiluted to 1.00 L with water (that is, it is diluted by a factor of 40).The standardized potassium permanganate solution is used totitrate the diluted and acidified hydrogen peroxide. The amountconcentration of the original hydrogen peroxide is obtained byanalysis of the titration evidence, and by using a graph, preparedwith a graphing calculator or computer spreadsheet program, ofthe amount and percent concentration of aqueous hydrogenperoxide.



600 Chapter 13 NEL

Section 13.4 Questions1. State the similarities and differences between the method

for redox stoichiometry and other examples ofstoichiometry.

2. In acid–base titrations, one reactant (usually the titrant) is astrong acid or strong base. Similarly, in a redoxstoichiometry, the reactant used to analyze a sample ofunknown concentration is either a strong oxidizing orstrong reducing agent. (a) State two common strong oxidizing agents commonly

used in a redox titration.(b) Using the redox table, suggest a strong reducing agent

that might be suitable for a redox titration analysis.(c) What are some examples of other strong reducing

agents that might be used in an analysis? What type ofexperimental design would be appropriate?

3. Why is it necessary to standardize a potassiumpermanganate solution to be used in a chemical analysis?

4. In a chemical analysis, 10.00 mL samples of aqueoushydrogen peroxide are acidified and then titrated with0.200 mol/L sodium perchlorate solution. From theevidence, an average volume of 24.0 mL of aqueous sodiumperchlorate is required to reach the endpoint. Calculate theamount concentration of the hydrogen peroxide solution.

5. Complete the Analysis and Evaluation (of the predictionand, thus, of the metallurgical process) of the investigationreport.

PurposeThe purpose of this lab exercise is to use redoxstoichiometry to evaluate a technological process.

ProblemWhat is the amount concentration of iron(II) ions in asolution obtained in an iron ore analysis?

PredictionAccording to the required standards for the metallurgicalprocess, the concentration of the iron(II) ions should be80.0 mmol/L.

DesignThe iron(II) solution is titrated to iron(III) with a standardcerium(IV) ion solution, which is reduced to cerium(III). Theindicator shows, as the endpoint, a sharp colour changefrom red to pale blue.

Evidence

6. A scientist used the titration method to analyze for tin(II)chloride. Complete the two steps of the Analysis.

ProblemWhat is the amount concentration of a tin(II) chloridesolution prepared from a sample of tin ore?

DesignThe potassium dichromate solution is first standardized bytitration with 10.00 mL of an acidified 0.0500 mol/L solutionof the primary standard, FeSO4•(NH4)2SO4•6 H2O(s). Thestandardized dichromate solution is then titrated against10.00 mL of the acidified tin(II) chloride solution.

Evidence

Extension

7. Potassium dichromate is a common reagent used in theanalysis of the iron content of iron ore samples. If eachanalysis begins with the same mass of the ore, a redoxtitration can be designed such that the volume ofdichromate required corresponds to the percent iron in theore. This design eliminates the need for any calculations, sorapid, efficient analyses can be carried out by technicians.Starting with a 1.00 g sample of iron ore, the sample istreated to convert all the iron into iron(II) ions, and thenacidified. Predict the concentration of potassiumdichromate required so that the volume (in millilitres)equals the percentage of iron in the original sample.

Table 4 Titration of 25.0 mL of Fe2�(aq) with 0.125 mol/L Ce4�(aq)

Trial 1 2 3 4



Table 5 Titration of 10.00 mL of 0.0500 mol/L Fe2�(aq)with K2Cr2O7(aq)

Trial 1 2 3 4



Table 6 Titration of 10.00 mL of Sn2�(aq) with K2Cr2O7(aq)

Trial 1 2 3 4




Chapter 13 INVESTIGATIONS


Single Replacement Reactions

This investigation is a review of single replacement reactionsin preparation for the development of a theory of oxidationand reduction. As part of the Design, include diagnostic tests(as in Appendix C.4) for the predicted products.

PurposeThe purpose of this investigation is to use the single replace-ment reaction generalization to predict and analyze the reac-tants and products.

ProblemWhat are the products of the single replacement reactionsfor the following pairs of reactants?

(a) copper and aqueous silver nitrate

(b) aqueous chlorine and aqueous sodium bromide

(c) magnesium and hydrochloric acid

(d) zinc and aqueous copper(II) sulfate

(e) aqueous chlorine and aqueous potassium iodide

Materialslab apron chlorine watereye protection magnesium ribbonfive small test tubes zinc striptwo test tube stoppers aqueous silver nitratetest tube rack aqueous sodium bromidesteel wool hydrochloric acidwash bottle aqueous copper(II) sulfatematches aqueous potassium iodidecopper strip hexane

Chapter 13

Procedure 1. Set up five test tubes, each filled to a depth of 2–3 cm

with one of the five aqueous solutions.2. Add the element indicated to each test tube.3. Perform diagnostic tests on each of the five mixtures.

Record your evidence.4. Dispose of the solutions as directed by your teacher.

Purpose Design AnalysisProblem Materials Evaluation (1)Hypothesis ProcedurePrediction Evidence

INVESTIGATION 13.1 Report Checklist

Toxic, corrosive, and irritant chemicals are used inthis investigation. Avoid skin contact. Wash anysplashes on the skin or clothing with plenty of water.If any chemical is splashed in the eye, rinse for atleast 15 min and inform your teacher.

Keep the hexane sealed to avoid evaporation.Dispose of the hexane as directed by your teacher.Hexane is highly flammable. Keep away from openflame. Make sure matches are extinguished bydipping in water. Do not inhale the vapours.


602 Chapter 13 NEL

Spontaneity of Redox Reactions

In previous units in this textbook, we assumed that all chem-ical reactions are spontaneous; that is, they occur once thereactants are placed in contact, without a continuous additionof energy to the system. Spontaneous redox reactions in solu-tion generally provide visible evidence of a reaction withina few minutes.

PurposeThe purpose of this investigation is to test the assumptionthat all single replacement reactions are spontaneous.

ProblemWhich combinations of copper, lead, silver, and zinc metalsand their aqueous metal ion solutions produce spontaneousreactions?

DesignA drop of each solution is placed in separate locations on aclean area of each of the four metal strips.

Materialslab aproneye protectionreusable strips of copper, lead, silver, and zinc metals

(Note that the lead strips bend much more easily than thezinc strips, which look similar.)

0.10 mol/L solutions of copper(II) nitrate, lead(II) nitrate,silver nitrate, and zinc nitrate in dropper bottles

steel wool or sandpaper



These chemicals are toxic—especially the leadsolution—and irritants. Avoid skin contact.Remember to wash your hands before leaving thelaboratory. Rinse all of the metal strips thoroughlyand return them so they can be used again.



Predicting the Reaction of SodiumMetal (Demonstration)

The process of developing theories, laws, and generalizationsrequires that they must be tested numerous times in as manydifferent situations as possible. This process is necessary notonly to determine their validity, but also to identify excep-tions that may lead to new knowledge.

As part of the Design, include a list of diagnostic tests usingthe “If [procedure] and [evidence], then [analysis]” formatfor every product predicted. (This format is described inAppendix C.4.)

PurposeThe purpose of this demonstration is to test the five-stepmethod for predicting redox reactions.

ProblemWhat are the products of the reaction of sodium metal withwater?

This reaction of sodium metal must be demonstratedwith great care, because a great deal of heat isproduced. Use only a piece the size of a small pea,use a safety screen, wear a lab apron, eye protection,and face shield, and keep observers at least twometres away.



Chapter 13

Analyzing a Hydrogen PeroxideSolution

In this investigation, you assume the role of a laboratory tech-nician working in a consumer advocacy laboratory, testing theconcentration of a hydrogen peroxide solution (Figure 1).

PurposeThe technological purpose of this investigation is to test andevaluate the percent concentration of the consumer solutionof hydrogen peroxide.

ProblemWhat is the percent concentration of hydrogen peroxide in aconsumer product?

DesignAn acidic solution of the primary standard, iron(II) ammo-nium sulfate–water (1/6), is prepared and the potassium per-manganate solution is standardized by a titration with thisprimary standard. A 25.0 mL sample of a consumer solutionof hydrogen peroxide is diluted to 1.00 L with water (that is,it is diluted by a factor of 40). The standardized potassiumpermanganate solution is used to titrate the diluted and acid-ified hydrogen peroxide. The amount concentration of theoriginal hydrogen peroxide is obtained by analysis of the titra-tion evidence, and by using a graph, prepared with a graphingcalculator or computer spreadsheet program, of the infor-mation in Table 1.

Materials

2. Measure the required mass of the iron(II) compoundin a clean, dry 100 mL beaker.

3. Dissolve the solid in about 40 mL of H2SO4(aq).

4. Transfer this solution into a clean 100 mL volumetricflask, rinsing and adding pure water to complete thepreparation of the standard solution.

5. Transfer 10.00 mL of the standard iron(II) solutionby pipette into a clean 250 mL Erlenmeyer flask.

6. Titrate the acidic iron(II) sample with KMnO4(aq).

7. Repeat steps 5 and 6 until three consistent volumes(within 0.1 mL) are obtained.

8. Transfer 10.00 mL of the diluted hydrogen peroxidesolution by pipette into a clean 250 mL Erlenmeyerflask.

9. Using a 10 mL graduated cylinder, add 5 mL ofH2SO4(aq) to the hydrogen peroxide solution.

10. Titrate the acidic hydrogen peroxide solution withKMnO4(aq).

11. Repeat steps 8 to 10 until three consistent volumes(within 0.1 mL) are obtained.

12. Dispose of all solutions into a labelled wastecontainer.



Figure 1In drugstores, hydrogenperoxide is usually sold as a3% solution. Hairdressersuse a 6% solution.

lab aproneye protectionFeSO4•(NH4)2SO4•6 H2O(s)2 mol/L H2SO4(aq)

diluted H2O2(aq)KMnO4(aq)(list to be completed by

student)

Procedure 1. (Pre-lab) Calculate the required mass of

FeSO4•(NH4)2SO4

•6H2O(s) to prepare 100.0 mL of a0.0500 mol/L solution.

Sulfuric acid is corrosive. Iron(II)ammonium sulfate, hydrogen peroxide,and potassium permanganate areirritants. Avoid inhaling any solid, andavoid skin or eye contact.

Table 1 H2O2(aq) Concentration

Amount Percentconcentration concentration(mol/L) (%)

0.73 2.5

0.76 2.6

0.79 2.7

0.82 2.8

0.85 2.9

0.88 3.0

0.91 3.1

0.94 3.2

0.97 3.3

1.0 3.4


Chapter 13 SUMMARY

604 Chapter 13 NEL

Outcomes

Knowledge

• define oxidation and reduction operationally (historically)and theoretically (13.1, 13.2, 13.3)

• define the following terms: oxidizing agent, reducing agent,oxidation number, half-reaction, disproportionation (13.1,13.2, 13.3)

• differentiate between redox reactions and other reactions byidentifying half-reactions and changes in oxidation number(13.1, 13.2, 13.3)

• identify electron transfer, oxidizing agents, and reducingagents in redox reactions that occur in everyday life in bothliving and nonliving systems (all sections)

• compare the relative strengths of oxidizing and reducingagents from empirical data (13.2)

• predict the spontaneity of a redox reaction based on a redoxtable, and compare predictions to experimental results (13.2)

• write and balance equations for redox reactions in acidic,basic, and neutral solutions, including disproportionationreactions, by using half-reaction equations, developingsimple half-reaction equations, and assigning oxidationnumbers (13.2, 13.3, 13.4)

• perform calculations to determine quantities of substancesinvolved in redox titrations (13.4)

STS

• state that a goal of technology is to solve practical problems(all sections)

• recognize that technological problems may require varioussolutions and have both intended and unintendedconsequences (13.1, 13.2, 13.3)

Skills

• initiating and planning: design an experiment to determinethe reactivity of various metals (13.1, 13.2); and describeprocedures for safe handling, storing, and disposal ofmaterials used in the laboratory, with reference to WHMISand consumer product labelling information (13.1, 13.2, 13.4)

• performing and recording: select and use appropriateequipment to perform a redox titration (13.4); use a standardredox table to predict the spontaneity of redox reactions(13.2, 13.4); and create charts, tables, or spreadsheets relatedto redox reactions (13.1, 13.2, 13.4)

• analyzing and interpreting: analyze evidence from anexperiment to derive a simple redox table (13.2); interpretpatterns and trends in redox reactions (all sections); andevaluate redox experiments, including identifying thelimitations of the evidence (13.1, 13.2, 13.4)

• communication and teamwork: work collaboratively inaddressing problems, and select and use appropriate modesof representation for redox reactions and answers to redoxproblems (all sections)

Key Terms

13.1half-reaction

reduction

oxidation

redox reaction

13.2reducing agent

oxidizing agent

redox spontaneity rule

disproportionation

13.3oxidation number

oxidation

reduction

MAKE a summary

1. Start with “redox” and make a flow chart or conceptmap that includes all of the Key Terms listed, plus anyimportant generalizations and procedures that will helpyou learn the material in this chapter.

2. Refer back to your answers to the Starting Pointsquestions at the beginning of this chapter. How hasyour thinking changed?

Go To

The following components are available on the Nelson Web site. Follow the links for Nelson Chemistry Alberta 20–30.

• an interactive Self Quiz for Chapter 13

• additional Diploma Exam-style Review questions

• Illustrated Glossary

• additional IB-related material

There is more information on the Web site wherever you seethe Go icon in this chapter.


Electric UniverseElectricity is everywhere. Only two centuries ago, peopledeveloped technologies to release and control electricity. Sincethen we have come to rely on it for almost every aspect of ourlives. A science writer talks about some of the characters whohelped to bring electricity into our lives.


EXTENSION +


Many of these questions are in the style of the DiplomaExam. You will find guidance for writing Diploma Exams inAppendix H. Exam study tips and test-taking suggestionsare on the Nelson Web site. Science Directing Words usedin Diploma Exams are in bold type.

DO NOT WRITE IN THIS TEXTBOOK.

Part 11. Historically,

��i��

meant producing ��

i�i��

from theirnaturally occurring compounds. According to moderntheory, this process involves a

��i�i�i��

of electrons.

The above statement is completed by the information inwhich row?

2. A reducing agent can be described as a substance thatA. loses electrons and causes reductionB. loses electrons and becomes reducedC. gains electrons and causes oxidationD. gains electrons and becomes reduced

3. Which one of the following general reaction types will notbe a redox reaction?A. combustionB. simple decompositionC. disproportionationD. double replacement

4. Some natural and technological processes that involveredox reactions are1. corrosion2. metallurgy3. cellular respiration4. photosynthesis5. rusting of iron6. magnesium metal flaresThe processes in which oxygen behaves as an oxidizingagent are, in numerical order,

___, ___, ___, and ___.

5. Which of the following combinations would produce aspontaneous redox reaction?A. nitric acid and iron(III) chloride solutionB. chromium metal and aqueous cobalt(II) chlorideC. oxygen gas bubbled into a sodium bromide solutionD. aqueous tin(II) nitrate and potassium iodide solutions

6. The strongest oxidizing agent isA. X2�(aq)B. Y2�(aq)C. Z2�(aq)D. W2�(aq)

7. The metal that has the weakest attraction for its electrons isA. W(s)B. Z(s)C. Y(s)D. X(s)

8. Which of the following solutions should not be stored in atin-plated container?I NaNO3(aq) III SnBr2(aq)II AgNO3(aq) IV Cl2(aq)

A. I onlyB. II and IIIC. II and IVD. III and IV

9. When the half-reaction equations are constructed usingcoefficients to balance electrons, the number of electronstransferred isA. 6B. 3C. 2D. 1

10. The coefficients of the net redox equation, in the order ofsubstances in the equation, are

___, ___, ___, and ___.

Chapter 13Chapter 13 REVIEW


Row i ii iii

A. oxidation metals gain

B. reduction metals gain

C. oxidation nonmetals loss

D. reduction nonmetals loss

NR Use this information to answer questions 9 and 10.

Ozone is a strong oxidizing agent that will oxidize aqueousiodide ions in an acidic solution to iodate ions. The unbalancedredox reaction equation is

____O3(g) � ____ I�(aq) → ____ IO3�(aq) � ____ O2(g)

Use this information to answer questions 6 and 7.

Four different metals and their corresponding metal ionsolutions are mixed to determine if a spontaneous reactionoccurs.

Table 1 Metal–Ion Reactions

X2�(aq) Y2�(aq) Z2�(aq) W2�(aq)

X(s) X X X ✓

Y(s) ✓ X ✓ ✓

Z(s) ✓ X X ✓

W(s) X X X X

✓ spontaneous reactionX no evidence of reaction


NR


606 Chapter 13 NEL

11. The oxidation number of the carbon atom in the carbonateion isA. �6B. �4C. –2D. 0

12. The oxidation numbers of the sulfur atoms in hydrogensulfide, sulfur dioxide, and sulfur trioxide are, in order ofthese compounds,A. 0, 0, 0B. �2, �4, �6C. �2, �4, �4D. �2, �4, �6

13. In the two-step conversion from hydrogen sulfide to sulfurtrioxide, the sulfur atoms areA. oxidized in both stepsB. reduced in both stepsC. oxidized first, then reducedD. reduced first, then oxidized

14. An unintended consequence of this process may beA. depletion of the ozone layerB. natural gas shortagesC. altering local climateD. acid rain (deposition)

15. In a standardization experiment, 25.0 mL of an acidic 0.100 mol/L tin(II) chloride solution required an averagevolume of 12.7 mL of potassium dichromate solution forcomplete reaction. The amount concentration of thepotassium dichromate solution is __________ mmol/L.

Part 216. Write a theoretical description of a redox reaction. Include

the following terms in your answer: electrons, oxidation,reduction, oxidizing agent, and reducing agent.

17. Use a table of relative strengths of oxidizing and reducingagents, such as the one in Appendix I, to answer thefollowing questions.(a) How can you predict whether or not a combination of

substances will react spontaneously?(b) If a spontaneous redox reaction occurs, what kinds of

evidence might be observed?

18. Define each of the following in terms of both electrons andoxidation numbers. (a) oxidation(b) reduction(c) redox reaction

19. For each of the following, complete the half-reactionequation and classify as an oxidation or reduction. (a) HClO2(aq) → HClO(aq) (acidic)(b) Al(OH)4

�(aq) → Al(s) (basic)(c) Br�(aq) → BrO4

�(aq) (acidic)(d) ClO�(aq) → Cl2(g) (basic)

20. Various pairs of metals and metal ions were combined andthe evidence interpreted, as shown below: 2 Ga(s) � 3 Cd2�(aq) → 2 Ga3�(aq) � 3 Cd(s) Ga(s) � Mn2�(aq) → no evidence of reaction3 Mn2�(aq) � 2 Ce(s) → 3 Mn(s) � 2 Ce3�(aq)(a) Use this information and the redox spontaneity rule to

develop a table of oxidizing and reducing agents forthese metals and their ions.

(b) Identify the strongest oxidizing and the strongestreducing agent in your table.

21. For the following solutions, list the entities believed to bepresent, and classify them as possible oxidizing or reducingagents. (a) aqueous chlorine solution(b) tin(II) nitrate solution(c) acidic potassium iodate solution

22. For each of the following mixtures, list and classify theentities present, predict the half-reaction and net ionicreaction equations, and predict whether or not aspontaneous reaction will be observed. (a) Chlorine gas is bubbled into an iron(II) sulfate solution.(b) Nickel(II) nitrate solution is mixed with a tin(II) sulfate

solution.(c) A zinc coating on a drain pipe is exposed to air and

water.(d) An acidic solution of sodium sulfate is spilled on a

steel laboratory stand. (Consider only the iron in thesteel.)

(e) For use in a titration, a sodium hydroxide solution isadded to a potassium sulfite solution to make it basic.

Use this information to answer questions 12 to 14.

Many natural gas wells, called “sour” gas wells, containconsiderable quantities of hydrogen sulfide gas as well asmethane. When this mixture burns (Figure 1), hydrogensulfide is converted to sulfur dioxide. Once in the atmosphere,sulfur dioxide may be converted to sulfur trioxide.

Figure 1A flaring gas well

NR

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NEL Redox Reactions 607

23. The reactivity of metals varies considerably from very lowreactivity (noble metals) to explosively reactive. Design anexperiment to study the reactivity of calcium metal in water.Your response should include:• Purpose• Problem• Prediction, including half-reaction and net ionic

equations• Design, including diagnostic tests• Materials, including WHMIS safety cautions• Procedure, including disposal instructions

24. Predict the balanced redox reaction equation byconstructing and labelling oxidation and reduction half-reaction equations.(a) Pt(s) � NO3

�(aq) � Cl�(aq) →PtCl6

2�(aq) � NO2(g) (acidic)(b) CN�(aq) � ClO2

�(aq) →CNO�(aq) � Cl�(aq) (basic)

(c) PH3(g) � CrO42�(aq) →

Cr(OH)4�(aq) � P4(s) (basic)

25. Assign oxidation numbers to all atoms/ions and indicatewhich atom/ion is oxidized and which is reduced. (a) 2 Al(s) � Fe2O3(s) → 2 Fe(s) � Al2O3(s)(b) In(s) � 3 Tl�(aq) → In3�(aq) � 3 Tl(s)(c) 2 Cr3�(aq) � Sn2�(aq) → 2 Cr2�(aq) � Sn4�(aq)(d) Cl2(aq) � 2 I�(aq) → 2 Cl�(aq) � I2(aq)(e) UCl4(s) � 2 Ca(s) → 2 CaCl2(s) � U(s)

26. Balance the following chemical equations using theoxidation number method. (a) C6H12O6(s) � O2(g) → CO2(g) � H2O(l)(b) Au3�(aq) � SO2(aq) → SO4

2�(aq) � Au(s) (acidic)(c) BrO3

�(aq) � C2H6O(aq) → CO2(g) � Br�(aq) (d) Ag(s) � NO3

�(aq) → Ag�(aq) � NO(g) (acidic)(e) HNO3(aq) � SO2(g) → H2SO4(aq) � NO(g)(f) Zn(s) � BrO4

�(aq) →Zn(OH)4

2�(aq) � Br� (aq) (basic)

27. Balanced redox equations can be obtained using threedifferent methods, other than trial-and-error for simpleequations. (a) Briefly describe each method.(b) Which method is also used to predict the products?

What other information can also be predicted?(c) For which methods do you need to know the primary

products?(d) Which method do you prefer to use? Why?

28. A commercial kit is available to clean silver by removing thetarnish using a redox reaction. (Assume that silver tarnish issilver sulfide.) A zinc strip is placed in a water softenersolution and the tarnished silver is placed so that it is incontact with the zinc strip. (a) Write the overall chemical equation and balance it

using the simplest possible method.(b) Verify, using oxidation numbers, that the chemical

equation is balanced.(c) Write oxidation and reduction half-reaction equations.

Chapter 13

29. Magnesium metal reacts rapidly in hot water. Predict themass of precipitate that will form if a 2.0 g strip ofmagnesium reacts completely with water.

30. A student uses a redox titration to determine theconcentration of iron(II) ions in an acidic solution. Theevidence in Table 2 shows the volume of 7.50 mmol/LMnO4

�(aq) that reacted with 10.0 mL of Fe2�(aq). Calculatethe amount concentration of the iron(II) ions.

31. Three chemistry teachers developed a problem to teststudents’ understanding of redox concepts. The challenge isto identify three unknown solutions (labelled A, B, and C)using any of the materials listed below in your procedure.Assuming all possible spontaneous reactions are rapid andthat the nitrate ion is a spectator ion, write a procedure toidentify which solution is sodium nitrate, which is lead(II)nitrate, and which is calcium nitrate. Describe theexpected results. 0.25 mol/L solutions of A, B, and Csilver, zinc, and magnesium stripsdropper bottles of 0.25 mol/L aqueous solutions of:

sodium sulfate; sodium carbonate; and sodium hydroxidesteel wooltest tubes and test tube rack50 mL beakers400 mL waste beaker

32. Photofinishing laboratories (Figure 2) often produce awaste solution containing silver ions. The CEO of a companywants to know if it is economical to recover this silver. Thefirst step in the study is to determine the concentration ofsilver ions in the waste solution. (a) Write an experimental design based on redox concepts

to determine the concentration of silver ions. (b) Using appropriate terms and chemical equations,

explain the redox chemistry in your proposed design.

Table 2 Titration of 10.0 mL of Fe2�(aq) with 7.50 mmol/L MnO4

�(aq)

Trial 1 2 3



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Figure 2In film processing,the film is placed ina developing agent.


608 Chapter 13 NEL

33. Methanol is used as a windshield-washer antifreeze;containers are usually labelled with the freezing point of thesolution (Figure 3). A chemical technician can test thevalidity of the claim using various experimental designs. Theexperimental design chosen below is the titration of a basicsolution of methanol with a standardized solution ofpotassium permanganate based on the following(unbalanced) chemical equation.CH3OH(aq) � MnO4

–(aq) →CO32–(aq) � MnO4

2–(aq)Use the information in Tables 3, 4, and 5 and complete the Analysis of the investigation report.

PurposeThe purpose of this investigation is to use redoxstoichiometry for a chemical analysis.

ProblemWhat is the freezing point of a sample of windshield-washerfluid?

DesignA potassium permanganate solution is prepared andstandardized against an acidic 0.331 mol/L solution ofiron(II) ammonium sulfate (Table 3). The standardizedpermanganate solution is then titrated against a basicmethanol solution, which has been diluted by a factor of1000 (Table 4).

Evidence

Extension

34. For the production of pulp from wood, a variety of methodsare used, including mechanical and chemical processes.These have advantages and disadvantages that have beenwidely debated. Prepare an argument for or against thefollowing statement: “The immediate economic value ofusing technology to produce a product far outweighs anypossible future adverse effects.”Your response should also include• researched information about a variety of mechanical

and chemical processes• an evaluation of these processes from technological,

economic, and ecological perspectives• reference to redox chemistry

35. Vanadium (Figure 4) is a very versatile element in terms ofits reactivity. Vanadium metal reacts with fluorine to formVF5, with chlorine to form VCl4, with bromine to form VBr3,with iodine to form VI2, with oxygen to form V2O5, and withhydrochloric acid to form VCl2.(a) Identify the oxidation states of vanadium in each of

these compounds.(b) What interpretation can be made about the oxidizing

power of chemicals that react with vanadium metal?(c) Describe how the oxidation state of vanadium relates

to the colours of the compounds formed.(d) Briefly describe some technological applications of

vanadium and its compounds.

Figure 3This methanol windshield-washerantifreeze can be usedat temperatures above �45 °C.

Table 5 Concentrations and Freezing Points of AqueousSolutions of Methanol

Amount concentration Percent by Freezing point(mol/L) mass (%) (°C)

0 0 0

6.035 20.00 �15.0

11.672 40.00 �38.6

16.754 60.00 �74.5

www.science.nelson.com GOTable 3 Titration of Potassium Permanganate Solution with10.00 mL Acidic FeSO4•(NH4)SO4(aq)

Trial 1 2 3 4



Table 4 Titration of Standardized PotassiumPermanganate Solution with 10.00 mL CH3OH(aq)

Trial 1 2 3



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Figure 4Vanadinite is the most common vanadium ore found in largereserves in many countries, including Canada.

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36. The earliest metallurgy would be classified aspyrometallurgy. Other processes such as hydrometallurgyand electrometallurgy are more recent inventions. Defineeach of these types of metallurgy and provide one commonexample of each. Why are all three of these processesexamples of redox reactions? What do these processesillustrate about the goal of technology and the interactionbetween science and technology?

37. The nitrogen cycle (Figure 5) is a very important andcomplex biological system that includes nitrogen fixation,nitrification, and denitrification. Describe the chemicalreactions in the nitrogen cycle.Your response should include:• definitions of the terms nitrogen fixation, nitrification,

and denitrification.• a summary of the main changes of nitrogen-containing

entities and their oxidation numbers in the nitrogencycle

• some examples of the half-reaction equations, labelledas oxidation or reduction

• descriptions of some positive and negativeenvironmental impacts of the processes in the nitrogencycle.

Chapter 13

eutrophication

gaseousatmospheric

nitrogen store

bacterianitrogen fixation

lightningfixation

runoff

fertilizers

organic matter

nitrogen fixationdenitrification

nitrites(NO2

–)nitrates(NO3

–)

gaseouslosses

N2 & N2O

plantconsumption

nitrification

nitrification

ammonium(NH4

+)

precipitation

fossil fuelemissions

leaching

leaching


Figure 5The nitrogen cycle


DE


In this chapter

Career Connections: Materials Engineering Technologist; Chemical Technologist

1414 Electrochemical Cells

chapter

Electrochemical Cells

Exploration: A SimpleElectric Cell

Investigation 14.1:Designing an Electric Cell

Web Activity: Hydrogen:Wonderfuel or Hype?

Web Activity: Lewis Urry

Case Study: The BallardFuel Cell

Investigation 14.2: AVoltaic Cell(Demonstration)

Web Activity: Voltaic CellsUnder StandardConditions

Investigation 14.3: TestingVoltaic Cells

Biology Connection:Reduction Potentials

Lab Exercise 14.A:Developing a Redox Table

Mini Investigation: HomeCorrosion Experiment

Web Activity: GalvanizingSteel

Investigation 14.4: APotassium IodideElectrolytic Cell

Investigation 14.5:Electrolysis(Demonstration)

Web Activity: ElectrolyticCell Stoichiometry

610 Chapter 14 NEL

Since their invention in 1888, vehicles powered entirely by electricity have drifted in andout of fashion (Figure 1). Many experts predict that in the next decade, electric vehicleswill finally make a breakthrough. Electric power is slowly becoming a viable alternativeto gasoline power, thanks to a combination of political, economic, and environmentalfactors. One advantage of electric cars over gasoline-fuelled cars is that they produce lesspollution. Also, while cars powered by gasoline engines are about 15% efficient, manyelectric cars are 90% efficient. (Of course, overall efficiency and environmental impactsdepend on how the electricity and gasoline are produced.) Other attractive features ofelectric vehicles are that they are nearly silent and require minimal maintenance.

The biggest obstacle to the widespread use of electric cars is the lack of a powerful, lightweight, inexpensive battery. Scientists and engineers are researching alter-natives to the common lead–acid battery. Perhaps the most promising alternative is a bat-tery that runs continuously as fuel is supplied. One such alternative is the aluminium–airfuel cell, which uses aluminium metal as the fuel and oxygen from the air to produce elec-tricity. Another possibility is a fuel cell in which a hydrogen-rich fuel and oxygen fromthe air produce electricity.

Redox reactions can produce electricity and, conversely, electricity can cause redoxreactions. Many materials that we take for granted were virtually unknown until theprocess of electrolysis made their production possible. Aluminium, chlorine, hydrogen,sodium hydroxide, magnesium, and copper are produced in large quantities by elec-trolytic processes. In this chapter, you will learn how batteries are made, how electricitycan be used to produce chemicals, and how science and technology work together inthe development of electrochemical processes.

Answer these questions as best you can with your current knowledge. Then, usingthe concepts and skills you have learned, you will revise your answers at the end ofthe chapter.

1. What concepts can we use to explain how electrochemical cells work?

2. Describe the relationship between science and technology in the development ofelectrochemical cells.

3. List the types and uses of a variety of common electric cells. Include an assessmentof the impact of each one on our lives.

STARTING Points


Electrochemical Cells 611NEL

Exploration A Simple Electric Cell

A cell that produces electricity can be amazingly simplebecause it uses very common materials and requires notechnical expertise to construct. Anyone can make one andthen improve its efficiency without much understanding of thescientific principles involved. That is why the electric cell wasused for more than 100 years before scientists understood howit worked.

Materials: copper and zinc metal strips (or any two differentmetals); steel wool; orange, apple, and potato (and other fruitsor vegetables); LCD clock; voltmeter (or multimeter) with leads

• Clean the metal strips with steel wool to remove any coatingor oxides.

• Stick both metal strips into the orange. Make sure that themetal strips are not in contact inside the orange.

• Momentarily touch the leads (red—positive; black—negative)from the voltmeter, one to each metal strip. Now reverse theleads and test again.

(a) Record and describe what happened in each case.• Connect the leads to the LCD clock, paying attention to

positive and negative connections.(b) Does the clock work? If it does not, suggest a solution to

make it work. Try it.(c) Explain, in your own words, what you think happened in (b).• Repeat the process using other fruits and vegetables.(d) Which fruit or vegetable seemed to be the best at producing

electricity?(e) What do all fruits and vegetables have in common?(f) How could you improve upon your electric cell?

Figure 1Today’s electric cars are a viable alternative togasoline-powered cars in urban environments.An electric car has electric motors that arepowered by rechargeable batteries.


612 Chapter 14 NEL

14.114.1 Technology of Cells and BatteriesBefore 1800, scientists knew that static electricity was produced by the friction createdby two moving objects in contact. They discovered ways of storing the charges tem-porarily, but when the energy was released in the form of an electrical spark, it could notbe put to practical use. Practical applications of electricity were developed after 1800, theyear in which Alessandro Volta announced his invention of the electric cell.

Volta invented the first electric cell but he got his inspiration from the work, almost30 years earlier, of the Italian physician Luigi Galvani. Galvani noticed that the musclesin a frog’s leg would twitch when a spark hit the leg. Galvani’s crucial observation wasthat two different metals could make the muscle twitch. Unfortunately, Galvani thoughthis discovery was due to some mysterious “animal electricity.” It was Volta who recog-nized that this effect had nothing to do with animals or muscle tissue, and everything todo with conductors and electrolytes, as you observed in the Exploration at the beginningof this chapter.

Cells and BatteriesAlthough an electric cell is a device that continuously converts chemical energy into elec-trical energy, the electric cells that Volta invented produced very little electricity. Eventually,he came up with a better design by joining several cells together. A battery is a group oftwo or more electric cells connected to each other in series, like railway cars in a train.Volta’s first battery consisted of several bowls of brine (aqueous sodium chloride) con-nected by metals that dipped from one bowl into the next (Figure 1). This arrangementof metal strips and electrolytes produced a steady flow of electric current.

Zn Cu

Figure 1A version of Volta’s first battery. Each beaker contains two different metals, copper and zinc, inan electrolyte, salt water. A series of beakers forms a series of cells (a battery) whose totalvoltage is the sum of the individual voltages of all cells.

Figure 2Volta’s revised cell design, simplerthan the first, consisted of asandwich of two metals separatedby paper soaked in salt water (theelectrolyte). A cell consisted of alayer of zinc metal separated from alayer of copper metal by the brine-soaked paper. A large pile of cellscould be constructed to give moreelectrical energy.

cell

battery

coppermetal

zincmetal

papersoakedin a saltsolution

+

+

–

–

Shocking PersonalExperiments“I introduced into my ears twometal rods with rounded ends andjoined them to the terminals of theapparatus. At the moment thecircuit was completed, I received ashock in the head—and began tohear a noise—a crackling andboiling. This disagreeablesensation, which I feared might bedangerous, has deterred me sothat I have not repeated theexperiment.”

Alessandro Volta (1745–1827)

DID YOU KNOW ??

Volta improved the design of this battery by replacing the strips of metal with flatsheets, and replacing the bowls with paper or leather soaked in brine. This producedmore electric current for a longer time. As shown in Figure 2, Volta stacked cells on topof each other to form a battery, known as a voltaic pile. When a loop of wire was attachedto the top and bottom of this voltaic pile, a steady electric current flowed. Volta assem-bled voltaic piles containing more than 100 cells.

Volta’s invention was an immediate technological success because it produced an electric current more simply and more reliably than methods that depended on staticcharges. It also produced a steady electric current—something no other device coulddo. The development of this technology led to many advances in physics (for example,



the theory and description of current electricity), in chemistry (for example, the dis-covery of Groups 1 and 2 metals), and in electrical and chemical engineering.

Basic Cell Design and PropertiesEach electric cell is composed of two electrodes (solid electrical conductors) and oneelectrolyte (aqueous electrical conductor) (Figure 3). In the cells we buy for home use,the electrolyte is usually a moist paste, containing only enough conducting solution tomake the cell function. The electrodes are usually two metals, or graphite and a metal.In some designs, one of the electrodes is the container of the cell. One of the electrodesis marked positive (�) and the other is marked negative (�). In an electric cell or battery,the cathode is the positive electrode and the anode is the negative electrode.

According to theory, electricity is the flow of electrons. Electrons move from the anodeof a battery through the external circuit to the cathode. A battery produces electricity onlywhen there is an external conducting path, such as a wire, through which electrons canmove. Disconnecting the wire from the battery immediately stops the flow of electrons.

A voltmeter is a device that can be used to measure the energy difference, per unitelectric charge, between any two points in an electric circuit (Figure 4(a)). The energydifference per unit charge is called the electric potential difference or the voltage, andis measured in volts (V). For example, the electrons transferred via a 1.5 V cell release onlyone-sixth as much energy compared with the electrons from a 9 V battery.

Since the voltage is a ratio of energy to charge, it is not dependent on the size of thecell. You may have noticed that you can buy the same type and brand of 1.5 V cells in avariety of sizes, such as AA, B, C, and D. All are rated at 1.5 V. The larger cells can storemore energy at the same time as transferring more charge, but the ratio of energy to chargeis the same as the smaller cells. The voltage of a cell depends mainly on the chemical com-position of the reactants in the cell.

Electric current, measured by an ammeter in amperes (A), is a measure of the rate offlow of charge past a point in an electrical circuit (Figure 4(b)). The larger the electriccell of a particular type, the greater the current that can be produced by the cell.The charge transferred by a cell or battery is measured in coulombs (C) and expressesthe total charge transferred by the movement of charged particles.

Section 14.1

voltmeter

electrolyte

+ —

anode(—)

cathode(+)

Figure 3A cell always contains twoelectrodes—an anode and acathode—and an electrolyte. Whentesting the voltage of a cell or battery,the red (�) lead of the voltmeter isconnected to the positive electrode(cathode), and the black (�) lead isconnected to the negative electrode(anode).

Figure 4(a) A dam built across a stream or river stops the flow of water. There is a potential energy

difference between a kilogram of water at the top of the dam and a kilogram of water atthe bottom of the dam. A disconnected cell or battery is like a dam holding back electrons.There is a potential energy difference between the electrons at the anode and theelectrons at the cathode. A voltmeter measures this potential energy difference.

(b) If water is released from behind the dam, it naturally flows from the region of higherpotential energy (behind the dam) to a lower potential energy below the dam. Similarly,when a circuit is connected to a cell or battery, the electrons naturally flow because thereis a difference in potential energy between the anode and the cathode.

(b)(a)

Electric Charge and CurrentAnalogyWe can extend the analogy inFigure 4. The mass (kg) of waterstored behind the dam is similar tothe available charge (C) stored inthe chemicals of a cell. Whenreleased, the current, or flow, ofwater (kg/s) is similar to thecurrent, or flow, of electrons (C/s).

DID YOU KNOW ??


614 Chapter 14 NEL

Technological Problem SolvingThe initial development of cells and batteries preceded much of the current scientificunderstanding of these devices. Cells and batteries existed almost 100 years before theelectron was discovered. The study of electric cells is a good illustration of tremendousadvances in technology based on very limited scientific knowledge. Technological devel-opment or problem solving is similar in some ways to scientific problem solving, but

• An electric cell must have two electrodes and an electrolyte.

• An electrode is a solid conductor.

• An electrolyte is an aqueous conductor.

• The cathode is the electrode labelled positive.

• The anode is the electrode labelled negative.

• The electrons flow through the external circuit from the anode to the cathode.

Components of an Electric CellSUMMARY

Aboriginal Science andTechnologyAboriginal peoples have a longhistory of technologicaldevelopment in many areas suchas agriculture, food preservation,medicine, and transportation. Justas in Western societies, Aboriginaltechnology had a practicalpurpose, was developed through atrial-and-error process, and wasinterdependent on Aboriginalscience using traditional orindigenous knowledge.

DID YOU KNOW ??

Practice1. What are the components of a simple electric cell?

2. Write an empirical definition of electrode and electrolyte, and a conventionaldefinition of anode and cathode.

3. If a DVD player requires 9 V to operate, how many 1.5 V “dry” cells connected in serieswould it need?

4. Differentiate between electric current and voltage.

5. Why do manufacturers of battery-operated devices print a diagram showing thecorrect orientation of the batteries? (Supply two answers to this question: one from ascientific perspective and one from a technological perspective.)

6. List several examples illustrating how a new technology (electric cells) led to new scientific discoveries.

Table 1 Electrical Quantities and SI Units

Quantity Symbol Meter Unit Unit symbol

charge Q — coulomb C

current I ammeter ampere A (1 A � 1 C/s)

potential difference V voltmeter volt V (1V � 1J/C)

power P — watt W (1 W � 1J/s)

energy density — — joules per J/kgkilogram

The power of a cell or battery is the rate at which it produces electrical energy. Power ismeasured in watts (W), and is calculated as the product of the current and the voltage ofthe battery. The energy density, or specific energy of a battery, is a measure of the quantityof energy stored or supplied per unit mass. Energy density may be measured in joules perkilogram (J/kg). Table 1 summarizes some important electrical quantities and their unitsof measurement.



its purpose differs. The purpose of technological problem solving is to find a realisticway around a practical difficulty—to make something work—while the purpose ofscientific problem solving is to describe, explain, and ultimately understand naturaland technological phenomena. Technology and science are dependent on each other.Although scientific knowledge can be used to guide the creation of a technology, the tech-nology may create new scientific understanding.

A systematic trial-and-error process is often used in technological problem solving(Appendix C.1). This is not a new process. Aboriginal peoples used a systematic trial-and-error process to ensure the survival of the tribe. An example of a trial-and-error processis as follows:

• Develop a general design for problem-solving trials; for example, select which vari-ables to manipulate and which to control.

• Follow several prediction–procedure–evidence–analysis cycles, manipulating andsystematically studying one variable at a time.

• Complete an evaluation based on criteria such as efficiency, reliability, cost, andsimplicity.

This technological problem-solving model was important in the early development ofpractical electric cells.

Section 14.1

Purpose Design AnalysisProblem Materials Evaluation (1, 3)Hypothesis ProcedurePrediction Evidence


Designing an Electric CellIn this cell, an aluminium soft-drink can is one of the electrodes.The other electrode is a solid conductor, such as a piece ofcopper wire or pipe, an iron nail, or graphite from a pencil. Theelectrolyte may be a salt solution, or an acidic or basic solution.Although many characteristics of a cell are important, only onecharacteristic, voltage, is investigated here. Check with yourteacher if you want to evaluate other designs and materials.

When evaluating the Purpose (Part 3), include your opinionabout the reliability, cost, and simplicity of your final electric cell.

PurposeThe purpose of this investigation is to make an electric cell.

ProblemWhat combination of electrodes and electrolyte gives the largestvoltage for an aluminium-can cell?

DesignIn a trial-and-error procedure, different variables (secondelectrode, electrolyte) are modified one at a time while all othervariables are held constant. The voltage of each cell is measuredas the responding variable.


A “Not Quite Dry” CellThe electrolyte in the “dry cell” isactually a moist paste (Figure 5).If the cell were completely dry, itwould not work because the ionsin the electrolyte must be able tocarry the electric current tocomplete the circuit. Just enoughwater is added so that the ions canmove, but not enough to make themixture liquid.

DID YOU KNOW ??

Consumer, Commercial, and Industrial CellsSince Volta’s invention of the electric cell and battery, there have been many advancesin electrochemistry and technology. Invented in 1865, the zinc chloride cell is commonly referred to as a dry cell because this design was the first to use a sealedcontainer. The first 9 V battery was made up of a series of 1.5 V dry cells (Figure 6).Both the 1.5 V dry cell and the 9 V battery are simple, reliable, and relatively inex-pensive. Other cells, such as the alkaline dry cell and the mercury cell (Table 2 onpage 616), were developed to improve the performance of the original dry cell. Oneproblem with all of these cells is that the chemicals eventually become depleted andirreversible reactions prevent these cells from being recharged. Cells that cannot berecharged are called primary cells.

Figure 5

carbonelectrode

1.5 V cell 9 V battery

MnO2 andNH4Cl

electrolytepaste

zincelectrode

Figure 6Like a flashlight D cell, the zincchloride dry cell on the left has avoltage of 1.5 V. The 9 V battery on theright is made up of six 1.5 V dry cells inseries.


616 Chapter 14 NEL

Secondary CellsSecondary cells can be recharged by using electricity to reverse the chemical reactionthat occurs when electricity is produced by the cell. Secondary cells and batteries includethe nickel–cadmium (Ni–Cd) cell and the lead–acid battery (Table 2 and Figure 7). A rel-atively recently developed secondary cell with a unique design is a lithium-ion cell, calledthe Molicel.

Table 2 Primary, Secondary, and Fuel Cells

Type Name of cell Half-reactions Characteristics and uses

primary dry cell (1.5 V) 2 MnO2(s) � 2 NH4�(aq) � 2 e� → Mn2O3(s) � 2 NH3(aq) � H2O(l) • inexpensive, portable, many

cells Zn(s) → Zn2�(aq) � 2 e� sizes• flashlights, radios, many

other consumer items

alkaline dry 2 MnO2(s) � H2O(l) � 2 e� → Mn2O3(s) � 2 OH�(aq) • longer shelf life; highercell (1.5 V) Zn(s) � 2 OH�(aq) → ZnO(s) � H2O(l) � 2 e� currents for longer

periods compared withdry cell

• same uses as dry cell

mercury cell HgO(s) � H2O(l) � 2 e� → Hg(l) � 2 OH�(aq) • small cell; constant voltage(1.35 V) Zn(s) � 2 OH�(aq) → ZnO(s) � H2O(l) � 2 e� during its active life

• hearing aids, watches

secondary Ni–Cd cell 2 NiO(OH)(s) � 2 H2O(l) � 2 e� → 2 Ni(OH)2(s) � 2 OH�(aq) • can be completely sealed;cells (1.25 V) Cd(s) � 2 OH�(aq) → Cd(OH)2(s)� 2 e� lightweight but expensive

• all normal dry cell uses, aswell as power tools, shavers, portable computers

lead–acid cell PbO2(s) � 4 H�(aq) � SO42�(aq) � 2 e� → PbSO4(s) � 2 H2O(l) • very large currents; reliable

(2.0 V) Pb(s) � SO42�(aq) → PbSO4(s) � 2 e� for many recharges

• all vehicles

fuel cells hydrogen– O2(g) � 2 H2O(l) � 4 e� → 4 OH�(aq) • lightweight; high efficiency; oxygen cell 2 H2(g) � 4 OH�(aq) → 4 H2O(l) � 4 e� can be adapted to use (1.2 V) hydrogen-rich fuels

• vehicles and space shuttle

aluminium–air 3 O2(g) � 6 H2O(l) � 12 e� → 12 OH�(aq) • very high energy density;cell (2 V) 4 Al(s) → 4 Al3�(aq) � 12 e� made from readily available

aluminium alloys• designed for electric cars

one cell

negative plates:lead screen filledwith spongy lead

+

—

—

—+ +

cathode

positive plates:lead screen filledwith PbO2(s)

cell spacer

anode

H2SO4(aq)electrolyte

in each cell

Figure 7The anodes of a lead–acid carbattery are composed of spongylead and the cathodes arecomposed of lead(IV) oxide on ametal screen. The large electrodesurface area is designed to deliversufficient current to start a carengine.

The MolicelFind out about a made-in-Canadatechnological development thathas had a big impact on portablepower.


EXTENSION +



Section 14.1

One of the most common and reliable secondary cells is the lead–acid cell in a typicalcar battery. The discharging of this cell (see lead–acid cell, Table 2) produces approxi-mately 2.0 V based on the following net equation.

discharging

Pb(s) � PbO2(s) � 2 H2SO4(aq) 2 PbSO4(s) � 2 H2O(l)

This cell requires the input (from the car’s alternator) of at least 2.0 V to force theproducts to change back to the reactants to recharge it. The half-reactions for the lead–acidcell listed in Table 2 must both be reversed to obtain the following net equation.

charging

2 PbSO4(s) � 2 H2O(l) Pb(s) � PbO2(s) � 2 H2SO4(aq)

A battery can be recharged if the products are stable with no further reactions occurring andif the products are able to travel through the electrolyte toward the appropriate electrode.

Fuel CellsA fuel cell is a different solution to the problem of the limited life of a primary cell. Fuelcells produce electricity by the reaction of a fuel that is continuously supplied to keepthe cell operating. In principle, the fuel cell could be used forever, provided the fuel is con-tinuously supplied. The fuel cell offers several advantages over methods that produceelectricity by the combustion of fossil fuels. For example, fuel cells generate electricitymore efficiently (Table 3), without producing greenhouse gases or substances that con-tribute to acid rain. The development of a cost-effective fuel cell is currently the focusof much scientific study and technological research and development.

William Grove accidentally invented the first fuel cell in 1839 using platinum elec-trodes, hydrogen, and oxygen as fuels, and sulfuric acid as the electrolyte. Grove wasactually studying the reverse process—using electricity to convert water into hydrogenand oxygen. After one experiment, he reconnected the two electrodes without a powersupply attached, and found that a small current was produced spontaneously as hydrogenand oxygen combined to form water. Grove continued to work on this cell, but eventu-ally decided it was not a practical device because the characteristics, such as voltage,current, and energy capacity, were quite poor. Although many scientists, including NobelPrize winners Fritz Haber and Walther Nernst, worked on improving the cell, they werelargely unsuccessful. They manipulated many variables, such as different electrodes andelectrolytes, but the reaction rates were too low and the electrodes became corroded.

Finally, in 1955, Francis Bacon succeeded where many others had failed. He produceda practical hydrogen–oxygen fuel cell using an alkaline electrolyte and electrodes con-structed of porous nickel (Figure 8). Although the idea had been around for a long time,Bacon’s cell was really the first practical fuel cell. NASA quickly adopted thehydrogen–oxygen fuel cell as an electrical power source for space flights because hydrogenand oxygen are already available for propulsion systems and the product, water, can bepurified for drinking. NASA’s fuel cell, a modification of the original Bacon cell, is an alka-line cell using potassium hydroxide as the electrolyte (Table 2 on page 616). It produces12 kW of electricity and operates at 70% efficiency. Unfortunately, NASA’s fuel cell isexpensive and has a relatively short working life, primarily due to the corrosive elec-trolyte. As a result, NASA’s cell is not economically viable for general or commercialapplications.

Learning TipDischarging a cell or battery islike letting the waterspontaneously run out from thehigher level behind a dam.Charging (or recharging) is likepumping the water up behind thedam. This is not a spontaneousprocess and requires energy.

Table 3 Efficiencies of Different Technologies*

Technology Efficiency*

fuel cells 40�70%

electric powerplants 30�40%

automobile engines 17�23%

gasoline lawn mower about 12%

*Efficiency is the fraction of the maximumavailable energy that is actually usable.

O2(g)in

cathode (+) anode (—)

O2(g)out

H2(g) in

H2(g) and H2O(g)out

electrolyte

Figure 8Hydrogen and oxygen gases arecontinuously pumped into thishydrogen–oxygen fuel cell. Each gasreacts at a different electrode.Unused gases are recycled.


618 Chapter 14 NEL

Aluminium–Air CellAnother type of fuel cell is the metal–air fuel cell, the most common of which is thealuminium–air cell (Table 2). This is actually an aluminium–oxygen cell and has beendeveloped for possible use in electric cars. Air is pumped into the cell and oxygen reactsat the cathode while a replaceable mass of aluminium reacts at the anode (Figure 9).The fuel is solid aluminium metal and the product, aluminium hydroxide, can be recy-cled back to aluminium metal. The simple design means that this cell can be assembledin almost any size. The high energy density of these cells results from the fact that threemoles of electrons are released from every mole of aluminium, and aluminium is a verylightweight metal. Unlike hydrogen, storage and transportation of a solid fuel do notpose a problem. Estimates from prototypes suggest that the aluminium anode will needreplacement every 2500 km in an electric car.

Large-Scale Commercial and Industrial Fuel CellsThe requirements for electrical power fuel cells for large-scale use in businesses andindustry are similar with regard to the fuel, but there is less concern about volume,weight, or energy density. However, businesses and industries need cells with muchlonger lifetimes. Fuel cells for large-scale commercial and industrial use are almost alwaysco-generation units. This means that they produce electricity as well as heat for spaceheating. Co-generation means that the overall efficiencies can be as high as 90%.Commercially viable fuel cells today are usually acid electrolyte cells such as the phosphoric acid fuel cell, which can produce 400 MW of power, sufficient for meetingthe electrical energy needs of a small city (Figure 10). These cells usually use naturalgas as a source of hydrogen for the fuel cell and operate at temperatures of 200 °C.

oxygen(air)

electrolyte

aluminium anode

oxygen(air)

gasdiffusioncathode

Figure 9Several Canadian companies havedone extensive research in thedevelopment of the aluminium–airsolid fuel cell.

Figure 10The world’s first commercialphosphoric acid fuel cell wasproduced by ONSI/InternationalFuel Cells. It has been availablesince 1992 and uses natural gas,waste methane, propane, orhydrogen as fuels. This unitproduces 200 kW electricity and 200kW heat at a total system efficiencyof 80%.

WEB Activity

Web Quest—Hydrogen: Wonderfuel or Hype?Taking the role of a research company, you and your group will investigate and prepare anillustrated presentation on the subject of hydrogen as a fuel. Is it clean? Is it inexpensive? Is it widely available? Will hydrogen become the wonderfuel of the 21st century or is hydrogenjust power hype?


Fuel CellsCar designers are working onfuturistic cars, powered by fuelcells, that would cut down onpollution. This video showspractical applications of scienceand technology that may, someday, bring us cheap, clean,efficient transportation.


EXTENSION +



Section 14.1

Practice7. Compare scientific and technological problem solving.

8. What steps are involved in technological problem solving?

9. Suppose you decided to develop and market an aluminium-can cell (see Investigation14.1.) How and why would you alter the electrolyte?

10. Distinguish between primary and secondary cells, including a common example ofeach.

11. What are some advantages and disadvantages of the zinc chloride dry cell?

12. What do the designs of a dry-cell container and an ice-cream cone have in common?

13. One of the most successful batteries has been the lead–acid car battery. (a) Identify the anode, cathode, and electrolyte.(b) How are the large currents produced that are necessary to start a car?(c) What has been the social impact of this battery?(d) What are some possible environmental impacts of this battery?

14. For both the hydrogen–oxygen fuel cell and the aluminium–air fuel cell,(a) write the two half-reaction equations (refer to Table 2, page 616)(b) label each equation from (a) as an oxidation or a reduction(c) write the net ionic equation for each cell

15. Using several perspectives, state some advantages and disadvantages of a fuel cell.

16. Assess the possible future importance of fuel cells in society.

17. Experimental cell phones that run on miniature hydrogen–oxygen fuel cells existtoday. What must be in place before the average consumer can buy one?

Figure 11Lewis Urry (1927–2004), inventor ofthe alkaline battery

WEB Activity

Canadian Achievers—Lewis UrryAfter graduating in chemical engineering from the University of Toronto in 1950, Lewis Urry(Figure 11) went to work for Eveready Battery Company where he developed the first practical,long-life electric cell.

1. Identify the unique feature in Urry’s cell (“battery”) and explain how it works.

2. List three familiar devices that depend on the alkaline cell.

3. Name one other cell developed by Urry.


Success or Failure?Thomas Edison, the Americaninventor, set out to invent asecondary battery. His plan was touse an alkaline electrolyte and ironas the anode, but he needed tofind a suitable cathode material.After thousands of experiments,his friend and associate WalterMallory commented on Edison’slack of results. Edison replied thathe had lots of results and he knewfifty thousand things that did notwork! Eventually he discoveredthat a thin metallic nickel filmproduced a battery that waslighter and more powerful than theexisting lead–acid batteries.

DID YOU KNOW ??


620 Chapter 14 NEL

The Ballard Fuel CellA variation of a hydrogen–oxygen fuel cell, also known as thehydrogen fuel cell, was developed for commercial applicationsby Ballard Power Systems in Vancouver, BC. The Ballard fuelcell uses a proton exchange membrane (PEM) in place of aliquid electrolyte. Normal electric cells use the ions in theliquid electrolyte to transfer electric charge within the cell. In ahydrogen fuel cell, the PEM is made from a solid proton-conducting polymer that transfers charge within the cell(Figure 12). The PEM is simple, robust, eliminates corrosiveliquids, and permits a high energy density.

The Ballard fuel cell consists of an anode and a cathodeseparated by a PEM. Hydrogen fuel admitted through a porousanode is converted into hydrogen ions (protons) and freeelectrons in the presence of a catalyst at the anode. Anexternal circuit conducts the free electrons and produces thedesired electrical current. Water and heat are produced whenthe protons, after migrating through the polymer membrane tothe cathode, react both with oxygen molecules from the airand with the free electrons from the external circuit.

Fuel cells can be connected in series (stacked) to increasethe voltage and power output. For example, an experimentaltransit bus uses an electric motor powered by a Ballard fuelcell that is capable of 205 kW (or 250 hp).

Ballard has development agreements with most major carmanufacturers to use its cells in future electric cars. The zero-emission engines convert hydrogen, or hydrogen-rich fuelssuch as natural gas and even methanol, into electricity,producing water and heat as the main byproducts.

Although the Ballard hydrogen fuel cell looks verypromising, there are several problems yet to be solved. Cost isa major factor, which may be partially solved by massproduction. The fuel is also under debate. If hydrogen gas isused, where does it come from? Electrolysis of water uses a lotof electrical energy and, if this energy comes from a fossil-fuelgenerating plant, a lot of pollution is produced along with thehydrogen gas. How would the hydrogen gas be distributedand stored on board the vehicle? There are important safetyconcerns associated with the handling and storage ofhydrogen, which is flammable. Many scientists and engineersbelieve that the solution is not to use hydrogen gas directly,but to use hydrogen-rich fuels. We have a lot of knowledge ofreforming hydrocarbons to produce hydrogen. If natural gas oreven gasoline were reformed as needed on board the vehicle,then we would have a familiar fuel source and aninfrastructure in place to supply this fuel. Not everyone agreesthat this is a good solution.

Hydrogen fuel cells may change the way that we useenergy, especially for transportation. Fuel cells may be usednot only to power cars, buses, and boats, but also manyportable devices.



1. Is hydrogen a “clean” fuel? Explain your answer.

2. Make a list, in order of importance, of the advantages of aBallard hydrogen fuel cell for urban buses compared withthose of a diesel engine for buses. Make a similar list ofpossible disadvantages.

3. How has the development of fuel cells been influenced bysociety and how might they influence society in thefuture? Include a variety of perspectives. Use informationfrom various print and electronic sources.

Extension

4. One solution to the hydrogen storage problem is to makehydrogen as it is needed using a reformer. Research howreformers can be used to produce hydrogen.

5. Iceland has ambitious plans to be the first nation toreplace all its fossil fuels with “clean” hydrogen. Researchhow Iceland plans to produce the hydrogen it will need.

air(oxygen) fuel (hydrogen)

fuelrecirculated

cathode anode

exhaustwater vapour(no pollution)

heat (90 °C)

proton exchangemembrane (PEM)

Figure 12The hydrogen fuel cell has the same design as Volta’s originalcell, but the electrolyte is a conducting solid.





Section 14.1

Section 14.1 Questions1. Draw a simple diagram of an electric cell and label:

electrodes, electrolyte, cathode, anode, signs for cathodeand anode, and direction of electron flow through anexternal wire.

2. State the evidence that an electric cell involves a redoxreaction.

3. List three types of electric cells used in consumer andcommercial operations. Briefly describe the main feature ofeach cell.

4. State two common examples of consumer cells and wherethey may be used.

5. A silver oxide cell is often used when a miniature cell orbattery is required, as in watches, calculators, and cameras.The following half-reaction equations occur in the cell:Ag2O(s) � H2O(l) � 2 e� → 2 Ag(s) � 2 OH�(aq) Zn(s) � 2 OH�(aq) → Zn(OH)2(s) � 2 e�

(a) In which direction does the electric current flow: silverto zinc or zinc to silver?

(b) Which is the anode and which is the cathode?(c) Write the net redox equation for the discharging of the

silver oxide cell.

6. Describe an example where a new technology led toscientific discoveries and an example in which scientificknowledge led to a better technology.

7. Possible solutions to technological problems often involvedifferent designs, materials, and processes. List severalexamples that solve the problem of cells becomingdepleted and unusable.

8. What are some unintended consequences of thewidespread use of many different types of cells andbatteries?

9. Describe how societal needs and expectations haveaffected the development of the electric cells availabletoday.

10. Suppose cells and batteries did not exist. What impactwould that have on your life?

11. State several reasons why it is important to recycle thecomponents of cells and batteries when they have reachedthe end of their useful life.

12. (a) Why is there a great deal of interest in electric cars?(b) Suggest some reasons why we don’t use lead–acid

batteries as the only power source for electric cars.(c) How have advances in hydrogen fuel cells facilitated

the development of electric cars?

Extension

13. Most people associate technological development with“progress” in Western societies and think that technologytransfer to Aboriginal peoples only occurs in one direction.In fact, although not often recognized and valued, there hasbeen significant technology transfer from Aboriginalpeoples to Western society. Summarize several examples of

significant Aboriginal technological knowledge that wastransferred to early European settlers. What area is still asignificant source of technological transfer from Aboriginalto Western societies? Why is this an area of dispute?

14. Portable electronic devices can be found everywhere.Laptop computers, cellular telephones, mobile radios, cordless phones, portable disc and MP3 players, and digital cameras all require an electric cell.(a) What are some of the requirements for cells used in

these applications?(b) Why are some rechargeable batteries used in various

portable devices supposed to be totally “drained”(discharged) before recharging?

15. People whose heart occasionally beats too slowly or tooquickly often have a pacemaker fitted to keep the heartbeating regularly (Figure 13). Pacemakers use a battery forelectric power. What kind of battery is commonly usedtoday? How long does it last? How does the doctor knowwhen the battery is nearing the end of its life and needs tobe replaced? Why are rechargeable batteries generally notused?

16. Plastic batteries were the dream of the 1980s, thedisappointment of the 1990s, and the subject of the 2000Nobel Prize for Chemistry. Now it appears that somecommercial products will eventually result from theresearch and development invested in plastic batteries.Briefly describe the electrodes and electrolyte for a plasticbattery. How is this battery similar to and different from anordinary battery? What are some advantages anddisadvantages?

Figure 13A pacemaker containselectronics and a built-inbattery. The whole unit isonly a few centimetres insize and is implantedunder the skin near thecollarbone.






622 Chapter 14 NEL

14.214.2 Voltaic CellsElectric cells were invented for practical purposes in about 1800, but they were notexplained scientifically until about 100 years after their invention. Their use, however, con-tributed to the scientific understanding of redox reactions. Later this knowledge helpedexplain reactions inside the cell itself. Electric cells adapted for scientific study are oftencalled galvanic cells (in recognition of Luigi Galvani), or voltaic cells (in recognition ofAlessandro Volta).

You learned in Chapter 13 that a redox reaction involves a transfer of electrons fromthe reducing agent to the oxidizing agent. In a voltaic cell, electrons pass from the reducingagent to the oxidizing agent through an external circuit rather than passing directly fromone substance to another. You have seen that the individual components of a cell—elec-trodes and electrolytes—determine characteristics such as voltage and current. Why isthis so? What happens in different parts of a cell? To answer these questions, chemistsuse a cell that has two electrodes and their electrolytes separated. This is not a very prac-tical arrangement, but it greatly facilitates the study of cells. Each electrode is in contactwith an electrolyte. A porous boundary separates the two electrolytes, at least for a shorttime, while still permitting ions to move between the two solutions through tiny open-ings in the cotton plugs of a salt bridge (Figure 1(a)) or in the walls of a porcelain cup(Figure 1(b)).

Using this design modification, a cell can be split into two parts connected by a porousboundary. Each part, called a half-cell, consists of one electrode and one electrolyte.For example, the copper–zinc cell shown in Figure 2 has two half-cells, copper metal ina solution of copper ions and zinc metal in a solution of zinc ions.

This cell can be represented using the following abbreviated (“shorthand”) notation,called a cell notation:

Zn(s) ⏐ Zn(NO3)2(aq) ⏐⏐ Cu(NO3)2(aq) ⏐ Cu(s)

electrolyte

cotton plugs

ionsions

electrolyte electrolyte

Figure 1(a) A salt bridge is a U-shaped

tube containing an inert(unreactive) aqueouselectrolyte such as sodiumsulfate.

(b) An unglazed porcelain (porous)cup containing one electrolytesits in a container of a secondelectrolyte.

(a)

zinc nitrateelectrolyte

salt bridge

copper electrode zinc electrode

copper(ll) nitrateelectrolyte

wire

NaNO3(aq)

Cu(s) | Cu(NO3)2(aq)half-cell

Zn(NO3)2(aq) | Zn(s)half-cell

Figure 2Each electrode is in its own electrolyte, forming a half-cell. The twohalf-cells are connected by a salt bridge (containing NaNO3(aq))and by an external wire to make a complete circuit.

electrolytesions

(b)



Section 14.2



A Voltaic Cell (Demonstration)An important characteristic of consumer, commercial, andindustrial cells is a simple, efficient design that works for theintended application. For scientific research, this is not asimportant as a design that can be easily manipulated and studied.

PurposeThe purpose of this investigation is to test the design andoperation of a voltaic cell used in scientific research.

ProblemWhat is the design and operation of a voltaic cell?

DesignAn electric cell with only one electrolyte is compared with voltaiccells containing the same electrodes, but two electrolytes.


In this notation, a single line (|) indicates a phase boundary, such as the interface of anelectrode and an electrolyte in a half-cell. A double line (||) represents a physical boundary,such as a porous boundary between half-cells. A voltaic cell is an arrangement of twohalf-cells separated by a porous boundary. Voltaic cells, such as the one in Figure 2, areespecially suitable for scientific study.

Learning TipIt is a common practice to writethe cell notation with the anodeon the left-hand side and thecathode on the right-hand side.Diagrams or photos of cells canhave any arrangement.

Table 1 Evidence and Interpretations of the Silver–Copper Cell

Evidence Interpretation

The copper electrode decreases in size Oxidation of copper metal is occurring:and the intensity of the blue colour of the Cu(s) → Cu2�(aq) � 2 e�

electrolyte increases. blue

The silver electrode increases in size as Reduction of silver ions is occurring:long, silver-coloured crystals grow. Ag�(aq) � e� → Ag(s)

A blue colour slowly moves up the U-tube Copper(II) ions (cations) move toward the from the copper half-cell to the silver half-cell. cathode.

An ammeter shows that the electric current Electrons move through the wire from the flows along a wire between the copper copper electrode to the silver electrode.electrode and the silver electrode.

A voltmeter indicates that the silver electrode Electrons have a tendency to leave the cathode (positive) and the copper copper half-cell and enter the silver half-electrode is the anode (negative). cell.

Learning TipPeople often use acronyms orsimilar devices to help themremember importantinformation. “OIL RIG (oxidationis loss, reduction is gain)” is anexample. One way to help youremember important details of acell is the expressionSOAC/GERC, loosely read as“soak a jerk.” Translated, thismeans the Strongest OxidizingAgent at the Cathode GainsElectrons and is Reduced at theCathode. Another example is“An ox ate a red cat,” whichhelps to recall anode oxidation;reduction cathode.

A Theoretical Description of a Voltaic CellObservation of a voltaic cell as it operates provides evidence to explain what is hap-pening inside the cell. For example, during your study of the silver–copper cell inInvestigation 14.2, you observed the evidence listed in Table 1. Note that the table alsoincludes a theoretical interpretation of each point, which is also shown in Figure 3.

According to the electron transfer theory and the concept of relative strengths of oxi-dizing and reducing agents, silver ions are the strongest oxidizing agents in the cell; theyundergo a reduction half-reaction at the cathode. The strongest oxidizing agent in thecell always undergoes a reduction at the cathode. Copper atoms, which are the strongestreducing agents in the cell, give up electrons in an oxidation half-reaction and enter the


624 Chapter 14 NEL

solution at the anode. The strongest reducing agentin the cell always undergoes an oxidation at theanode. Therefore, the cathode is the electrode wherereduction occurs and the anode is the electrodewhere oxidation occurs.

• The strongest oxidizing agent present in the cellalways undergoes a reduction at the cathode.

• The strongest reducing agent present in the cellalways undergoes an oxidation at the anode.

Electrons released by the oxidation of copper atomsat the anode travel through the connecting wire tothe silver cathode. We can explain the direction ofelectron flow in terms of competition for electrons.According to the Table of Relative Strengths ofOxidizing and Reducing Agents (redox table) inAppendix I, silver ions are stronger oxidizing agentsthan copper(II) ions. Silver ions win the tug of war forthe electrons available from the conducting wire.

To write the net equation for the silver–coppervoltaic cell, we first identify the strongest oxidizing

and reducing agents present in the mixture. (This is the same procedure you followedwhen predicting redox reactions in Section 13.2.) We then follow the same procedure forpredicting half-reactions in which the two materials are in contact with each other.

OA SOA

Cu(s) ⏐ Cu2�(aq) ⏐⏐ Ag�(aq) ⏐ Ag(s)SRA RA

reduction at the cathode: 2 [Ag�(aq) � e� → Ag(s)]oxidation at the anode: Cu(s) → Cu2�(aq) � 2 e�

net: Cu(s) � 2 Ag�(aq) → Cu2�(aq) � 2 Ag(s)

Experimental evidence shows that all parts of the cell remain electrically neutral. Youcan see from the half-reaction equations that silver ions are removed from the solutionwhen they gain electrons to form silver solid. Because the solution started as electricallyneutral, the removal of positively charged ions would mean the solution would becomenegatively charged. However, based on the evidence, this does not happen. Cations, whichin this case are Na�(aq) ions, move from the salt bridge into the solution in the cathodecompartment to maintain an electrically neutral solution (refer to Figure 3). Similarly,as positively charged copper(II) ions form at the anode, anions, which in this case areNO3

�(aq) ions, move from the salt bridge into the anode compartment to maintainelectrical neutrality.

For a general voltaic cell, we can summarize the components and processes for a gen-eral voltaic cell using the cell notation and cell diagram shown in Figure 4.

anode (—) | electrolyte || electrolyte | cathode (+)

electrons

(reduction)(oxidation)cationsanions

Figure 4In an operating voltaic cell, theelectrical circuit is completed by theelectron flow in the external part (wires) of the cell and the ion flow inthe internal part (solutions) of the cell.

cathode (+) anode (–)

Cu(s)Ag(s)

anode half-cell

Cu(s) Cu2+(aq) + 2 e–

cathode half-cell

Ag+(aq) + e– Ag(s)

(reduction)

Na+(aq)

Na+(aq)

Ag+(aq)

Ag+(aq)

Cu2+(aq)NO3–(aq)

NO3–(aq)

NO3–(aq)

NO3–(aq)

NO3–(aq)

(oxidation)

Cu2+(aq)

e–e–

Figure 3A theoretical interpretation of thesilver–copper cell

cathode (+)anode (—)

anions

cations

e—e—



Voltaic Cells with Inert ElectrodesFor cells containing metals and metal ions, the electrodes are usu-ally the metals, and half-reactions take place on the surface of themetals. What happens if an oxidizing or a reducing agent otherthan these is used? For example, an acidic dichromate solution isa strong oxidizing agent that reacts spontaneously with coppermetal. To construct this cell, you can use a copper half-cell, asshown in Figure 5, but an electrode is also required for the dichro-mate half-cell. You cannot use solid sodium dichromate as anelectrode because solid ionic compounds do not conduct elec-tricity and solid sodium dichromate would also dissolve in thesolution. You need a solid conductor that will not react in the cellor interfere with the desired cell reaction. In other words, you need an unreactive orinert electrode. Inert electrodes provide a location to connect a wire and a surface onwhich a half-reaction can occur. A carbon (graphite) rod or platinum metal foil are twoinert electrodes that are commonly used as a cathode or anode.

Section 14.2

anions

e–

cations

Cu(s)anode (–)

C(s)cathode (+)

Cu2+(aq) Cr2O72–(aq)

H+(aq)

Figure 5The copper electrode decreases inmass and the blue colour of theelectrolyte increases, which indicatesoxidation at the anode. The carbonelectrode remains unchanged, butthe orange colour of the dichromatesolution becomes less intense andchanges to greenish-yellow,evidence that reduction is occurringin this half cell.

(a) Write equations for the half-reactions and the overall reaction that occur in the following cell:

C(s) | Fe2�(aq), Fe3�(aq) || Cr2O72�(aq), H�(aq) | C(s)

(b) Draw a diagram of the cell, labelling electrodes, electrolytes, the direction of electronflow, and the direction of ion movement.

Solution(a)

(b)


OA SOA OA

SRA

C(s) ⏐ Fe2�(aq), Fe3�(aq) ⏐⏐ Cr2O72�(aq), H�(aq) ⏐ C(s)

Cr2O72�(aq) � 14 H�(aq) � 6 e� → 2 Cr 3�(aq) + 7 H2O(l)

Cr2O72�(aq) � 14 H�(aq) � 6 Fe2�(aq) → 6 Fe 3�(aq) + 2 Cr 3�(aq) + 7 H2O(l)

6 [ Fe2�(aq) → Fe3�(aq) + e� ]cathode:anode:

net:

e—

C(s)anode(—)

cations

anions

C(s)cathode

(—)

Fe3�(aq)Fe2�(aq), Cr2O72—(aq), H�(aq)

Faraday and the Term “Ion”The term “ion,” from the ancientGreek meaning “going” was firstused by Michael Faraday in 1834.

“Finally, I require a term toexpress those bodies whichcan pass to the electrodes … I propose to distinguish thesebodies by calling those anionswhich go to the anode of thedecomposing body; and thosepassing to the cathode, cations;and when I have occasion tospeak of these together, I shallcall them ions.”

DID YOU KNOW ??


626 Chapter 14 NEL

Learning TipA variety of names can be usedfor cells based uponspontaneous redox reactions:electric, voltaic, galvanic, andelectrochemical. In this book,electric cell is used forconsumer cells and voltaic cellis used for scientific researchcells. Electrochemical cell is ageneral term that includes bothvoltaic (spontaneous) andelectrolytic (nonspontaneous)cells.

• A voltaic cell consists of two half-cells separated by a porous boundary with solidelectrodes connected by an external circuit.

• The cathode is the positive electrode. The strongest oxidizing agent is reduced atthe cathode.

• The anode is the negative electrode. The strongest reducing agent is oxidized atthe anode.

• Electrons travel in the external circuit from the anode to the cathode.

• Internally, anions move toward the anode and cations move toward the cathodeas the cell operates. The solutions remain electrically neutral.

• Cell notation: anode ⏐ electrolyte ⏐⏐ electrolyte ⏐ cathode where a single vertical line represents a phase boundary and a double vertical linerepresents a porous boundary.

Voltaic CellsSUMMARY

Practice1. Write an empirical description of each of the following terms: voltaic cell, half-cell,

porous boundary, salt bridge, electrolyte, external circuit, and inert electrode.

2. Write a theoretical definition of a cathode and an anode.

3. Indicate whether the following processes occur at the cathode or at the anode of avoltaic cell. (a) reduction half-reaction(b) oxidation half-reaction(c) reaction of the strongest reducing agent(d) reaction of the strongest oxidizing agent

4. When is an inert electrode used? Give two common examples.

5. What are the characteristics of the solution in a salt bridge? Provide an example.

6. For each of the following cells, use the given cell notation to identify the strongestoxidizing and reducing agents. Write chemical equations to represent the cathode,anode, and net cell reactions. Draw a diagram of each cell, labelling the electrodes,electrolytes, direction of electron flow, and direction of ion movement. (a) Zn(s) ⏐ Zn2�(aq) ⏐⏐ Ag�(aq) ⏐Ag(s)(b) Al(s) ⏐ Al3�(aq) ⏐⏐ NO3

�(aq), H�(aq) | Pt(s)

7. Ions move through a porous boundary between the two half-cells of a voltaic cell. (a) In what direction do the cations and anions move?(b) Why do the ions move? Take your answer and convert it into another “why”

question. Now answer this question.(c) For the copper–silver cell shown in Figure 3 (page 624), we know that Na�(aq)

and NO3�(aq) in the salt bridge move into the cathode and anode compartments,

respectively. Explain the evidence that the blue colour moves up the salt bridgetowards the cathode (refer to Table 1, page 623)?

8. Draw and label a diagram for a voltaic cell constructed from some (not all) of thefollowing materials:

strip of cadmium metal voltmeterstrip of nickel metal connecting wiressolid cadmium sulfate glass U-tubesolid nickel(II) sulfate cottonsolid potassium sulfate various beakersdistilled water porous porcelain cup

9. Redesign the voltaic cell in question 8 by changing at least one electrode and oneelectrolyte. The net reaction should remain the same for the redesigned cell.

Electron Sources and SinksChemists sometimes refer to theanode as the electron source andthe cathode as the electron sink.Look at some half-reactionequations to see why these termsapply.

DID YOU KNOW ??



Section 14.2

Standard Cells and Cell PotentialsThe investigations and activities you have completed show that the design and compositionof a cell affect its operation. To make comparisons and scientific study easier, chemists specifythe composition of a cell and the conditions under which the cell operates. A standard cellis a voltaic cell in which each half-cell contains all entities shown in the half-reaction equa-tion at SATP conditions, with a concentration of 1.0 mol/L for the aqueous solutions. If a metalis not part of a half-cell, then an inert electrode is used to construct the standard cell. Forexample, for a standard zinc–dichromate cell, the cell description is

Zn(s) ⏐ Zn2�(aq) ⏐⏐ Cr2O72�(aq), H�(aq), Cr3�(aq) ⏐ C(s) at SATP

1.0 mol/L 1.0 mol/L 1.0 mol/L 1.0 mol/L

The standard cell potential, E °cell , is the maximum electric potential difference(voltage) of the cell operating under standard conditions; E°cell represents the energy dif-ference (per unit of charge) between the cathode and the anode. The degree sign (°) indi-cates that standard 1.0 mol/L and SATP conditions apply. Based on the idea of competitionfor electrons, a standard reduction potential, E°r, represents the ability of a standardhalf-cell to attract electrons, thus undergoing a reduction. The half-cell with the greaterattraction for electrons—that is, the one with the more positive reduction potential—gains electrons from the half-cell with the lower reduction potential. The standard cellpotential is the difference between the reduction potentials of the two standard half-cells.

E °cell � E°r � E°rcathode anode

It is impossible to determine experimentally the reduction potential of a single half-cellbecause electron transfer requires both an oxidizing agent and a reducing agent. Notethat a voltmeter can only measure a potential difference, E °cell. To assign values for standard reduction potentials, we measure the “reducing” strength of all possible half-cells relative to an accepted, standard half-cell. The half-cell used for this purpose is thestandard hydrogen half-cell. A half-cell such as this, that is chosen as a reference and arbitrarily assigned an electrode potential of exactly zero volts, is called a reference half-cell.

Standard Hydrogen Half-CellThe standard hydrogen half-cell (Figure 6) consists of an inert platinum electrodeimmersed in a 1.00 mol/L solution of hydrogen ions, with hydrogen gas at a pressure of100 kPa bubbling over the electrode. The pressure and temperature of the cell are keptat SATP conditions. Standard reduction potentials for all other half-cells are measuredrelative to that of the standard hydrogen half-cell. The reduction potential of the hydrogenion reduction half-reaction is defined to be exactly zero volts.

2 H�(aq) � 2 e�0 H2(g) E°r � 0.00 V

As a result, we can assign a numerical value to the reduction potential associated with everyother reaction. A reduction potential that has a positive value means that the oxidizingagent is a stronger oxidizing agent than hydrogen ions. A negative reduction potential meansthat the oxidizing agent is a weaker oxidizing agent than hydrogen ions. The choice of thestandard hydrogen half-cell as a reference is the accepted convention. If a different half-cellhad been chosen as the reference, individual reduction potentials would be different, but theirrelative values would remain the same.

Learning TipThink of the standard cellpotential E°cell as representingthe difference in ability of twohalf-cells to gain electrons. Thispotential difference can only bemeasured accurately if nocurrent is allowed to flow. Agood-quality voltmeter has alarge internal resistance toprevent current flow.

connecting wire

1.00 mol/L H+(aq)at 25 °C

Pt(s)

Pt(s) | H2(g) | H+(aq) E r = 0.00 V°

H2(g)at SATP

Figure 6The standard hydrogen half-cell isused internationally as the referencehalf-cell in electrochemical research.Notice that the second vertical linein the cell notation designates thephase boundary between the gasand the liquid.


628 Chapter 14 NEL

Measuring Standard Reduction PotentialsWe can measure the standard reduction potential of a half-cell by constructing a stan-dard cell using a hydrogen reference half-cell and the half-cell whose reduction poten-tial you want to measure. There are two things you need to know: the voltage and thedirection of the current. The magnitude of the voltage determines the numerical valueof the half-cell potential and the direction of the current determines the sign of the half-cell potential. The cell potential is measured with a voltmeter, which also shows thedirection the electrons tend to flow. If E°cell is positive, then the positive terminal on thevoltmeter is connected to the cathode and the oxidizing agent at the cathode is strongerthan hydrogen ions.

The cell shown in Figure 7 can be represented as follows:

Pt(s) ⏐ H2(g) ⏐ H�(aq) ⏐⏐ Cu2�(aq) ⏐ Cu(s) E°cell � �0.34 Vanode cathode

Learning TipA voltmeter has two terminals,positive (red) and negative(black). Connect these to theelectrodes of any cell so thatthe voltmeter gives a positivereading. Whatever electrode isconnected to the positiveterminal will be the cathode,and the other electrode will bethe anode.

cathode: Cu2�(aq) � 2 e� → Cu(s)

anode: H2(g) → 2 H�(aq) � 2 e�

net: Cu2�(aq) � H2(g) → Cu(s) � 2 H�(aq)

E°cell � E°r � E°rcathode anode

� 0.34 V � 0.00 V� �0.34 V

+0.34

0.34 V

0.00

E° (V)Cu2+(aq) + 2 e— 0 Cu(s)

2 H+(aq) + 2 e— 0 H2(g)

Figure 8Copper(II) ions are strongeroxidizing agents than hydrogenions. The cell potential provides aquantitative measurement of howmuch stronger.

The voltmeter shows that the copper electrode is the cathodeand is 0.34 V higher in potential than the platinum anode. Ifyou were to replace the voltmeter with a connecting wire sothat the current is allowed to flow, the blue colour of thecopper(II) ion disappears and the pH of the hydrogen half-celldecreases as more hydrogen ions are produced and the solu-tion becomes more acidic. Based on this evidence, copper(II)ions are being reduced to copper metal and hydrogen mole-cules are being oxidized to hydrogen ions. Since this redoxreaction is spontaneous, copper(II) ions are stronger oxidizing

agents than are hydrogen ions. The standard cell potential, E°cell � 0.34 V, is the differencebetween the reduction potentials of these two half-cells (Figure 8).

Suppose a standard aluminium half-cell was set up with a standard hydrogen half-cell (Figure 9). We can represent the cell as

Al(s) ⏐ Al3�(aq) ⏐⏐ H�(aq) ⏐ H2(g) ⏐ Pt(s) E °cell � �1.66 Vanode cathode

According to the voltmeter, the platinum electrode is the cathode and the aluminium elec-trode is the anode. This indicates that hydrogen ions are stronger oxidizing agents thanaluminium ions by 1.66 V. Since the reduction potential of hydrogen ions is defined as0.00 V, the reduction potential of the aluminium ions must be 1.66 V below that ofhydrogen, or �1.66 V (Figure 10).

1.00 mol/LCu2�(aq) at 25 °C

1.00 mol/LH�(aq) at 25 °C

H2(g)25 °C100 kPa

0.34 V

Cu(s)

Pt(s)

Figure 7A copper–hydrogen standard cell



2 H�(aq) � 2 e�0 H2(g) E°r � 0.00 V

Al3�(aq) � 3 e�0 Al(s) E°r � �1.66 V

The standard cell potential, E°cell � 1.66 V, is the difference between the reduction poten-tials of these two half-cells. To obtain the net or overall cell reaction, add the reductionand oxidation half-reactions, but remember to balance and cancel the electrons.

cathode: 3 [2 H�(aq) � 2 e� → H2(g)]

anode: 2 [Al(s) → Al3�(aq) � 3 e�]

net: 6 H�(aq) � 2 Al(s) → 3 H2(g) � 2 Al3�(aq)


� 0.00 V � (�1.66 V)

� �1.66 V

Notice that the half-reaction equations were multiplied by appropriate factors to balancethe electrons, but the reduction potentials were not altered by the factors used to balancethe electrons. Electric potential represents energy per coulomb of charge (1 V � 1 J/C).Multiplying the aluminium half-reaction by a factor of 2 doubles both the energy andthe charge transferred, so that the ratio of energy (J) to charge (C), that is the voltage,is unaffected.

In both of these examples, the strongest oxidizing agent is reduced at the cathode andthe strongest reducing agent is oxidized at the anode. The measured cell potential is thedifference between the reduction potentials at the cathode and at the anode.

A positive cell potential (E°cell > 0) indicates that the net reaction is spontaneous—a requirement for all voltaic cells.

Section 14.2

1.66 V

Pt(s)

1.00 mol/LH+(aq) at 25 °C

1.00 mol/LAl3+(aq) at 25 °C

Al(s)H2(g)25 °C100 kPa

Figure 9 An aluminium–hydrogen standard cell

Learning TipAltering the coefficients in ahalf-reaction equation does notaffect the reduction potential.

0.00

—1.66

1.66 V

E° (V)

Al3+(aq) + 3 e— 0 Al(s)

2 H+(aq) + 2 e— 0 H2(g)

Figure 10On a redox table, hydrogen ions arestronger oxidizing agents thanaluminium ions. The cell potentialtells us the hydrogen ions are 1.66 Vabove aluminium ions.


Figure 11Measurements of standard cellpotentials show that the reductionpotential of Cu2�(aq) is �0.34 Vgreater than that of H�(aq), which is�1.66 V greater than that ofAl3�(aq). If you set up a standardcell using copper and aluminium,what would be the cell potential,E°cell. (Answer: �2.00 V)

+0.34

0.34 V

0.00

—1.66

1.66 V

E° (V)

Al3+(aq) + 3 e— 0 Al(s)

2 H+(aq) + 2 e— 0 H2(g)

Cu2+(aq) + 2 e— 0 Cu(s)

You can analyze a standard cell knowing the contents of both half-cells using one ormore of the following steps:

• Determine which electrode is the cathode. The cathode is the electrode where thestrongest oxidizing agent present in the cell reacts, i.e., the oxidizing agent that isclosest to the top on the left side of the redox table. If required, copy thereduction half-reaction for the strongest oxidizing agent and its reductionpotential.

• Determine which electrode is the anode. The anode is the electrode where thestrongest reducing agent present in the cell reacts, i.e., the reducing agent that isclosest to the bottom on the right side of the redox table. If required, copy theoxidation half-reaction (reverse the half-reaction by reading from right to left)for the strongest reducing agent and its reduction potential.

• Determine the overall cell reaction. Balance the electrons for the two half-reaction equations (but do not change the E°r) and add the half-reactionequations.

• Determine the standard cell potential, E °cell, using the equation:


Rules for Analyzing Standard CellsSUMMARY

630 Chapter 14 NEL

Figure 11 shows the combined results from the copper–hydrogen and aluminium–hydrogen standard cells. This process of measuring standard cell potentials can quicklybe extended to more and more oxidizing agents. Notice that this process, although startedwith the hydrogen reference cell, does not require that it be used for all cell measurements.For example, knowing that the reduction potential of copper(II) ions is 0.34 V, we cannow set up many cells that include a standard copper half-cell. A more extensive list ofreduction potentials is found in the redox table in Appendix I. You can predict the reac-tion that occurs spontaneously in any voltaic cell operating under standard conditionsusing the redox table in Appendix I. The standard cell potential is predicted as follows:


This order of subtraction is necessary to find out whether the reaction is spontaneousunder standard conditions. If E°cell is positive, the reaction is spontaneous.

WEB Activity

Simulation—Voltaic Cells Under Standard ConditionsIn this computer simulation, you will construct voltaic cells using different half-cells, and thenanalyze their components and processes. A series of questions will prompt you as you formgeneralizations about the spontaneity of various reactions.




Section 14.2

Learning TipTo ensure a correctinterpretation, always write thecathode half-reaction of theSOA first. This will help you toremember to subtract thereduction potentials in thecorrect order.

A standard lead–dichromate cell is constructed. Write the cell notation, label the electrodes, and calculate the standard cell potential.

Solution

Pb(s) ⏐ Pb2�(aq) ⏐⏐ Cr2O72�(aq), H�(aq), Cr3�(aq) ⏐ C(s)

anode cathode

E °cell � 1.23 V � (�0.13 V)

� �1.36 V


A standard scandium–copper cell is constructed and the cell potential measured. Thevoltmeter indicates that copper electrode is positive.

Sc(s) ⏐ Sc3�(aq) ⏐⏐ Cu2�(aq) ⏐Cu(s) E°cell � �2.36 V

Write and label the half-reaction and net equations, and calculate the standard reduction potential of the scandium ion.

Solution

cathode: 3 [Cu2�(aq) � 2 e� → Cu(s)] E °r � �0.34 V

anode: 2 [Sc(s) → Sc3�(aq) � 3 e�] E °r � ?

net: 3 Cu2�(aq) � 2 Sc(s) → 3 Cu(s) � 2 Sc3�(aq) E °cell � �2.36 V

2.36 V � 0.34 V � E °r

E °r � �2.02 V


Practice10. Standard cells are very important in the scientific study of voltaic cells.

(a) Describe the contents and conditions of a standard cell. (b) Define the standard cell potential in words and in symbols.

11. For each of the following cells, write the equations for the reactions occurring at thecathode and at the anode, and an equation for the overall or net cell reaction.Calculate the standard cell potential. (Use the redox table in Appendix I.)(a) Cr(s) ⏐ Cr2�(aq) ⏐⏐ Sn2�(aq) ⏐ Sn(s)(b) Co(s) ⏐ Co2�(aq) ⏐⏐ SO4

2�(aq), H�(aq), H2SO3(aq) ⏐ C(s)(c) Pt(s) ⏐ H2(g) ⏐ OH�(aq) ⏐⏐ OH�(aq) ⏐ O2(g) ⏐ Pt(s)



Testing Voltaic CellsTesting is a procedure that is common to both technology andscience. In technology, testing is necessary to determine how aproduct or process works using criteria such as efficiency,reliability, and cost. In science, testing is a key part in theadvancement of knowledge. Scientific concepts are developedand then tested to determine their validity and limitations. Newideas that fail the test then need to be restricted, revised, orreplaced.

In your Evaluation, pay particular attention to sources of erroror uncertainty, and to limitations of the evidence collected.

PurposeThe purpose of this investigation is to test the predictions of cellpotentials and the identity of the electrodes of various cells.

ProblemIn cells constructed from various combinations of copper,aluminium, silver, and zinc half-cells, what are the standard cellpotentials, and which is the anode and cathode in each case?



632 Chapter 14 NEL

Figure 12(a) The water behind the gates in a lock has a certain potential energy, �E, relative to the

bottom of the closed outlet.(b) When the outlet is opened, water spontaneously flows to the lower level on the other side of

the gates. Potential energy, �E, is converted to kinetic energy of the flowing water. Thewater flowing through the outlet is analogous to electron flow.

(c) The flow of water ceases when the levels on both sides of the gates become equal. Thegates open, and the ship can then exit to the next lock.

(a) (b) (c)

Cell Potentials Under Nonstandard ConditionsThe electric potential difference or voltage of a cell decreases slowly as the cell operates.Simultaneously, observable changes such as colour changes and precipitate formation occur.If the cell is left for a very long time, the voltage would eventually become zero and no fur-ther changes would be observed in the cell. When people refer to a “dead”cell or battery, thisis often what they mean.

The electric potential difference of a cell is a measure of the tendency for electrons toflow. Ideally, during a measurement of the cell potential, a voltmeter should not allowany electrons to flow. If electrons flow, oxidation and reduction reactions occur which, inturn, change the concentrations from the standard 1.0 mol/L value. The value that ismeasured by a voltmeter represents an electric potential or stored energy, just as the waterbehind a lock in a canal has gravitational potential energy (Figure 12(a)). Connectingthe electrodes of a cell in a circuit allows the electrons to flow from the anode to thecathode. This is analogous to opening the outlet and allowing the water to flow frombehind the gates to a lower point in front of the gates (Figure 12(b)). In both cases, storedpotential energy is converted to kinetic energy of electrons or water. If the water availablebehind a lock is allowed to flow out, then eventually no more water will flow. The level(potential energy) of the water on the two sides of the gate is equalized. An equilibriumis reached with no potential energy difference (Figure 12(c)). A similar situation occurswith an operating cell. If electrons are allowed to flow, eventually an equilibrium will bereached when the flow ceases. The rate of the forward reaction, which predominates ini-tially, decreases as the rate of the reverse reaction increases, until the two rates become equal.This is the equilibrium condition and no net flow of electrons will occur. At this time, theelectric potential difference as measured by a voltmeter becomes zero.

Discrepancies between Measured and Predicted Cell PotentialsNonstandard conditions for concentrations, temperature, and pressure will cause dif-ferences between the cell potentials predicted from a standard redox table and onesmeasured in a laboratory. However, these differences are generally small if the condi-tions are relatively close to standard values. Other more important reasons for discrep-ancies include the purity of the substances, the presence of oxide coatings on metals,and the type and size of the porous boundary.


BIOLOGY CONNECTION

Reduction PotentialsElectron transfer in biological systems is governed by the sameconcepts as electron transfer insimpler chemical systems. In bothsystems, reduction potentials areused to rank important half-reactions. Cytochromes areproteins that are part of theelectron transfer process incellular respiration; ferredoxins areproteins that are an important partof the electron transfer inphotosynthesis. These proteinsattach to a metal ion such as Fe3�,greatly influencing its ability totransfer electrons.

(free) Fe3�(aq) � e� 0 Fe2�(aq)E°r � �0.77 V

(in cytochrome c) Fe3�(aq) � e� 0 Fe2�(aq)E°r � �0.25 V

(in ferredoxin)Fe3�(aq) � e� 0 Fe2�(aq)E°r � �0.43 V



Section 14.2

Practice12. For each of the following standard cells, refer to the redox table in Appendix I, to

represent the cell using the standard cell notation (listing the anode first). Identify thecathode and anode and calculate the standard cell potential without writing half-reaction equations.(a) copper–lead standard cell(b) zinc–nickel standard cell(c) iron(III)–hydrogen standard cell

13. One experimental design for determining the position of a half-cell reaction that isnot included in a redox table is shown below. Use the following standard cell, refer tothe standard reduction potential of gold in Appendix I, and calculate the reductionpotential for the indium(III) ion.

In(s) ⏐ In3�(aq) ⏐⏐ Au3�(aq) ⏐ Au(s) E °cell � �1.84 Vanode cathode

14. Any standard half-cell could have been chosen as the reference half-cell, the zeropoint of the reduction potential scale. What would be the standard reductionpotentials for copper and zinc half-cells, assuming that the standard lithium cell werechosen as the reference half-cell with its reduction potential defined as 0.00 V? Justify your answer.

15. List some reasons for differences that might be observed between cell potentials predicted from a table of reduction potentials and cell potentials measured in a laboratory.

16. A zinc–iron cell is constructed and allowed to operate until the measured potentialdifference becomes zero. What interpretation can be made about the chemicalsystem at this point?


Developing a Redox TableWe can use standard cells and their measured cell potentials todevelop a redox table.

PurposeThe purpose of this exercise is to use the concepts and rules ofstandard cells to develop a redox table.

ProblemWhat is the table of relative strengths of oxidizing and reducingagents based on measured cell potentials?

DesignSeveral cells are investigated; each cell has at least one half-cellin common with one of the other cells. The cell potentials aremeasured, and the positive and negative electrodes of each cellare identified.

EvidenceNegative Positive electrode electrode

Pd(s) ⏐ Pd2�(aq) ⏐⏐ Cr2O72�(aq), H�(aq) ⏐ C(s) E°cell � �0.28 V

Ti(s) ⏐ Ti2�(aq) ⏐⏐ Tl�(aq) ⏐ Tl(s) E°cell � �1.29 V

Tl(s) ⏐ Tl�(aq) ⏐⏐ Pd2�(aq) ⏐ Pd(s) E°cell � �1.29 V

LAB EXERCISE 14.A Report Checklist


634 Chapter 14 NEL

CorrosionHuman history is often divided into different “ages” such as the Copper, Bronze, andSteel ages. These descriptions are based on when these metals were first widely refinedand used for tools and weapons. The process of refining a metal is electrochemical innature and requires energy to recover the pure metal from its naturally occurring com-pounds (ores). Corrosion is also an electrochemical process in which a metal reacts withsubstances in the environment, returning the metal to an ore-like state. Because we livein an oxidizing (oxygen) environment, oxidation (corrosion) of some metals occurspontaneously. In fact, we need to produce metals such as iron continually to replacethe metals lost to corrosion. Preventing corrosion and dealing with the effects of corrosionare major economic and technological problems for our society (Figure 13).

As a metal is oxidized, metal atoms lose electrons to form positive ions. A redox tableof relative strengths of oxidizing and reducing agents provides the evidence that metalsvary greatly in their ability to be oxidized. Some metals, such as gold and silver, are“noble” because they are relatively weak reducing agents. On the other hand, Group 1 and2 metals are very strong reducing agents and are, therefore, easily oxidized. In general,any metal appearing below the various oxygen half-reactions in a redox table will beoxidized in our environment. Iron (including steel) and aluminium are such metals,and are extensively used as structural materials. Why is the corrosion or rusting of ironsuch a major problem, but the corrosion of aluminium, which is a much strongerreducing agent, is not? The answer lies primarily in the nature of the oxide that formson the surface of the metal. A freshly cleaned surface of aluminium rapidly oxidizes inair to form aluminium oxide.

4 Al(s) � 3 O2(g) → 2 Al2O3(s)

The aluminium oxide adheres tightly to the surface of the metal. This prevents furthercorrosion by effectively sealing any exposed surfaces.

Unfortunately, the iron compounds that form on the surface of exposed iron do notadhere very well. They flake off, exposing new iron to be corroded. In addition, the corrosion of iron is a complex process that is significantly affected by the presence ofsubstances other than oxygen.

Rusting of Iron Studies of the corrosion of iron have shown that the presence of both oxygen and wateris required and the iron is converted into iron hydroxides and oxides. The first step of themechanism is thought to be the oxidation of iron at a wet exposed surface (Figure 14).

Fe(s) → Fe2�(aq) � 2 e�

Iron(II) ions diffuse through the water on the iron surface while the electrons easilytravel through the iron metal, which is an electrical conductor. The electrons are pickedup by oxygen molecules dissolved in water on the surface at a point away from the orig-inal oxidation site (Figure 14).

�12� O2(g) � H2O(l) � 2 e� → 2 OH�(aq)

The combination of iron(II) ions and hydroxide ions forms a low-solubility precip-itate of iron(II) hydroxide, which is further oxidized by oxygen and water to formiron(III) hydroxide, a yellow-brown solid. The familiar red-brown rust is formed by thedehydration of iron(III) hydroxide to form a mixture of iron(III) hydroxide and hydratediron(III) oxide. The amount of the hydroxide and the oxide varies, so rust is referred toas a hydrated oxide of indeterminate formula, Fe2O3• x H2O(s).

Figure 13Large ships have steel hulls. Therusting of steel involves theoxidation of iron in the steel and is a constant headache for shippingcompanies.

DID YOU KNOW ??Rates of CorrosionA tin can (tin on steel) will corrodecompletely in about 100 a; analuminium can in about 400 a; and aglass bottle in about 100 ka. (a, orannum, is the SI unit for year.)

DID YOU KNOW ??Hydrated OxideIron(III) hydroxide can be convertedto iron(III) oxide trihydrate as shownbelow:

2 Fe(OH)3(s) → Fe2O3•3 H2O(s)

It is difficult to determine how muchof the iron(III) exists in rust as thehydroxide or hydrated oxide.



This simplified mechanism for the rusting of iron can be used to explain why certainconditions promote rusting. If the iron is kept in a dry environment (low humidity) orif air has been removed from the water, little or no corrosion occurs (Figure 15).Eliminating either water or the oxygen in the water makes the reduction of aqueousoxygen impossible. Iron cannot be oxidized unless a suitable oxidizing agent is present.If oxidizing agents other than oxygen are present, such as certain metal ions, nonmetals,or hydrogen ions, the iron can still be corroded through spontaneous redox reactions.This helps to explain the corrosion of iron in acidic environments; for example, whyacid rain corrodes iron faster than natural rain does.

In general, electrolytes accelerate rusting. Ships rust more rapidly in seawater thanin fresh water and cars rust more rapidly in places where salt is used on roads. Chlorideions from salt are known to inhibit the adherence of protective oxide coatings onmany metals, thus exposing more metal to be corroded. Electrolytes like sodium chlo-ride conduct electricity and improve charge transfer, accelerating the rusting process.Plumbers know that you cannot use steel (iron) straps or nails to hold copper pipesin place because corrosion of the iron will be accelerated. Any moisture that is presentsets up an electric cell similar in principle to Volta’s original discovery of electricity fromdissimilar metals (Section 14.1). As the cell operates, the iron corrodes to form rust.In general, the rusting of iron requires the presence of oxygen and water and is acceleratedby the presence of acidic solutions, electrolytes, mechanical stresses, and contact with lessactive metals.

Section 14.2

e—

OH—

Fe2+

rustairO2(g)

cathode

ironobject

anode

waterH2O(I)

O2(g) + H2O(l) + 2 e— 2 OH—(aq) Fe(s)

e—

Fe2+(aq) + 2 e—12

Figure 14The corrosion of iron is a smallelectrochemical cell with ironoxidation at one location (theanode) and oxygen reduction atanother location (the cathode).

Figure 15Rusting of exposed iron is almostnegligible when the relative humidityis less than 50%. This iron pillar inDelhi, India, has existed for about1500 years because of the very dryand unpolluted environment.

mini Investigation Home Corrosion Experiment

Soft drinks are acidic and contain electrolytes. Would differenttypes of soft drink corrode iron at different rates?

Materials: soft drinks (cola, lemon-lime flavoured soda), 2 identical steel nails, 2 plastic glasses(a) Predict which drink will cause the nails to corrode faster.• Test your prediction. Place a clean steel nail in a plastic cup

filled with the cola (Figure 16). Put the other nail in theother cup filled with the other soft drink.

• After at least one day, examine the nails for evidence of anychanges.

(b) Record and explain your results.• Dispose of all materials as directed by your teacher.

Figure 16What happens to a steelnail in clear soda?


636 Chapter 14 NEL

Corrosion PreventionCorrosion is such an important problem in our society that many technologies have beendeveloped, and continue to be developed, to minimize this problem. There are also manycareer opportunities as engineers or technologists. Methods used for preventing orminimizing the corrosion of iron can be divided into two categories: barrier methods thatemploy protective coatings, and the method of cathodic protection. In some critical situations,such as a large fuel tank, both methods may be used.

Protective CoatingsPaint and other similar coatings are a simple method of corrosion prevention. Thismethod works well as long as the surface is completely covered and the coating remainsintact. Unfortunately, a scratch or chip in the surface can easily expose a small surfaceof iron and corrosion begins.

Both tin and zinc are used as metallic coatings. Tin, as in the familiar tin can, adhereswell to the iron and provides a strong, shiny coating. The outer surface of the tin coatinghas a thin, strongly adhering layer of tin oxide that protects the tin as long as the foodstored in the can is not too acidic. If a crack or break occurs in the tin layer, moisture cancollect in the crack and an electric cell with tin and iron electrodes is established. Sinceiron is more easily oxidized than tin, iron becomes the anode in this cell. The electronsreleased by the oxidation of iron flow to the tin and corrosion is accelerated. Evidenceof this is the typical iron rust on tin cans that have been crushed and left outside.

A spontaneous electric cell also arises when a zinc coating on an iron object is broken.However, in this case, the zinc is more easily oxidized than the iron. The zinc is preferentiallyoxidized, preventing corrosion of the iron. Zinc plating (galvanizing) of steel or iron pro-vides double protection: a protective layer and preferential corrosion of the zinc.

Cathodic Protection According to the redox theory of a cell, oxidation is the loss of electrons and occurs atthe anode of a cell. Therefore, an effective method of preventing corrosion of iron iscathodic protection in which the iron is forced to become the cathode by supplyingthe iron with electrons, using either an impressed current or a sacrificial anode.

For a battery or DC generator connected in a circuit, electrons flow out of the negativeterminal and into the positive terminal. If the negative terminal is connected to the ironobject and the positive terminal is connected to an inert carbon electrode, an electriccurrent is forced to flow to the iron through an electrolyte, such as ground water, from thecarbon electrode. The iron is forced to become the cathode and is prevented from corroding.An impressed current is an electric current forced to flow toward an iron object by anexternal potential difference. This method of corrosion prevention requires a constantelectric power supply (typically 8 mV) and is used as cathodic protection for pipelines andculverts.

A simpler method of cathodic protection is the use of a sacrificial anode. A sacrificialanode is a metal more easily oxidized than iron and connected to the iron object to be pro-tected. The practice of zinc plating (galvanizing) iron objects is a common example of thismethod. Sacrificial zinc anodes are also connected to the exposed underwater metal sur-faces of ships and boats to prevent the corrosion of the iron in the steel. Blocks of mag-nesium can also be used as sacrificial anodes (Figure 17). In all cases, the more activemetal (appearing below iron in a half-reaction table) is slowly consumed or sacrificed atthe anode, forcing the iron object to be the cathode of the cell.

copper wire

undergroundsteel tank

magnesium

electron flow

Figure 17Corrosion of iron involves theoxidation of iron at the anode of acell. If the iron is attached directlyor connected electrically to a metalthat is more easily oxidized (asacrificial anode), then aspontaneous cell develops in whichiron is the cathode. The electrolyteof the cell is the moisture in theground.

CAREER CONNECTION

Materials EngineeringTechnologistHow does materials engineeringrelate to corrosion? Research theentrance requirements, jobprospects, and typical salaries fortechnologists in this field.

Corrosion PreventionYou might have seen shiny,reddish copper roofs turn green-grey, after they have been installedfor a while. This is a natural patina,that actually helps prevent furthercorrosion. Scientists are workingon developing an artificial patina,that protects the metal much morequickly. Listen to this audio clip tofind out more about this process.

EXTENSION +





Section 14.2

Figure 18The metal in this fence is galvanizedsteel: steel coated with zinc. As thezinc surface layer corrodes, it formsa protective coating that preventsthe steel from corroding.

Practice17. What are the minimum requirements for the corrosion of iron?

18. List some factors that accelerate or promote the corrosion of iron.

19. Write the balanced net ionic equation for the corrosion of iron to iron(II) ions in thepresence of oxygen and water.

20. Although the corrosion of iron is a serious problem, other metals are also corroded inair or other environments. For each of the following situations, use your knowledge ofwriting and balancing redox equations to write and label the half-reaction and netionic equations: (a) Zinc is an active metal that oxidizes when exposed to air and water.(b) A lead pipe corrodes if it is used to transport acidic solutions that also contain

dissolved oxygen.(c) In dry air, minute quantities of hydrogen sulfide gas can slowly react with silver

objects to produce hydrogen gas and silver sulfide, recognized by the dark tarnishon the surface of the silver.

21. You may have noticed that when a car body rusts, the rust appears around the breakor chip in the paint, but the damage may extend under the painted surface for somedistance.(a) What is the evidence for damage extending well beyond the break in the paint?(b) Suggest an explanation why the damage may extend far from the break in the

paint.

22. Would a basic solution prevent or slow down the corrosion of iron? Provide yourreasoning.

23. Why is a zinc coating on iron better than a tin coating?

24. What are the two methods of cathodic protection and how are they similar?

WEB Activity

Case Study—Galvanizing SteelA galvanized (zinc-coated) chain-link fence can last a long time without deteriorating (Figure 18). In this computer simulation, you will learn about the steps involved in thegalvanizing process. After you observe the computer simulations, describe the method andpurpose of each of the cleaning and galvanizing steps in the production of galvanized steel.Evaluate the corrosion resistance of a galvanized pipe compared with a painted pipe.


Section 14.2 Questions1. In the context of a voltaic cell, write a definition of each of

the following terms: half-cell, cathode, anode, cation, anion,and inert electrode.

2. State the function of a porous boundary in a voltaic celland describe two common examples.

3. Distinguish between the external circuit and the internalcircuit of a cell.

4. Define a standard cell.

5. (a) For a given cell, how is the cell potential predicted?(b) What are the restrictions on this prediction?

6. How does the cell potential indicate spontaneity of thereaction?

7. Why are the reactions in voltaic cells always spontaneous?What does this imply about the cell potential?

8. Define the hydrogen reference cell, including contents andconditions. Include the half-cell notation.

9. Why is a reference half-cell necessary?

10. (a) What is the cell potential of a standard cobalt–zinc cell?(b) What is the theoretical interpretation of this cell

potential? (c) List some factors that may account for the differences

you would see between the experimental and predictedvalues of the cell potential for the cobalt–zinc cell.


638 Chapter 14 NEL

11. For each of the following cells• identify the strongest oxidizing and reducing agents• write chemical equations to represent the cathode,

anode, and overall (net) cell reactions (include the half-cell and cell potentials)

• draw a diagram of each cell, labelling the electrodes,polarity (signs) of electrodes, electrolytes, direction ofelectron flow, and direction of ion movement

(a) Zn(s) ⏐ Zn2�(aq) ⏐⏐ Cu2�(aq) ⏐ Cu(s)(b) Sn(s) ⏐ Sn2�(aq) ⏐⏐ Cr2O7

2�(aq), H�(aq) ⏐ C(s)

12. You can determine a possible identity of an unknown half-cell from the cell potential involving a known half-cell. Usethe following evidence and the redox table in Appendix I todetermine the reduction potential and possible identity ofthe unknown X2�(aq) | X(s).

2 Ag�(aq) � X(s) → 2 Ag(s) � X2�(aq) E °cell � �1.08 V

13. An important goal of technology is to provide solutions topractical problems in society. Illustrate this goal byidentifying a common problem and the technologicalsolution.

Extension

14. A zinc wire is connected to and buried with a pipelinewhen it is built (Figure 19).

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) — Cu(s) | Cu2+(aq) || Ag+(aq) | Ag(s)

Ag(s)Zn(s)

Cu(s)Cu(s)

Figure 20Two standard cells in series

(a) Why is this done? Include a brief description of theprinciples involved.

(b) Is this the only type of corrosion protection used withmajor pipelines?

(c) Discuss the environmental and safety issues associatedwith protecting and also not protecting pipelines.

15. Complete the Prediction for the following investigation.Include your reasoning.

ProblemWhat is the total electric potential difference of two cellsconnected in series?

DesignCopper–silver and copper–zinc standard cells areconnected as shown in Figure 20. The total electricpotential difference of the two cells is measured with avoltmeter connected to the silver and zinc electrodes.


Figure 19When this pipeline was beingconstructed, a zinc wire was attachedto and buried with the pipe.



14.314.3Electrolytic CellsElectric cells and batteries used by consumers contain reactants chosen to react spon-taneously to convert their chemical energy into electrical energy. These cells or batteriescan be used to power a portable music player, start a car, or plate silver metal on jewellery.

A scientific research cell, or voltaic cell, produces electricity spontaneously becauseeach half-cell contains both oxidized and reduced entities. For example,

Zn(s) ⏐ Zn2�(aq) ⏐⏐ Pb2�(aq) ⏐ Pb(s)

The cell potential, E°cell, is always greater than zero. In a redox table, the strongest oxi-dizing agent present in the cell will always be above the strongest reducing agent present.

If a cell does not contain all oxidized and reduced species shown in the half-reactionequation, it is possible that the reactants (electrodes and electrolyte) present will notreact spontaneously. For example, if lead electrodes are placed in a solution of zinc sul-fate and the electrodes are connected with a wire, there is no evidence of any reaction.

Pb(s) ⏐ ZnSO4(aq) ⏐ Pb(s)

The strongest oxidizing agent present in this cell is Zn2�(aq) and the strongest reducingagent present is Pb(s). A quick check in the redox table shows that the oxidizing agent,Zn2�(aq), is well below the reducing agent, Pb(s), and the E°cell is negative (Figure 1).

Pb2+(aq) + 2 e— 0 Pb(s)

Ni2+(aq) + 2 e— 0 Ni(s)

Fe2+(aq) + 2 e— 0 Fe(s)

Zn2+(aq) + 2 e— 0 Zn(s)

Figure 1According to the redox spontaneityrule, if the strongest oxidizing agentpresent is below the strongestreducing agent present, nospontaneous reaction will occur(E°cell � 0).

Learning Tip

Standard CellsRecall that a standard cellcontains all entities listed in theequation for the half-reaction.In addition, the concentration ofaqueous entities is 1.0 mol/Land the conditions are SATP forall substances, including gasesat 100 kPa.

Figure 2Cominco in Trail, British Columbia,operates the world’s largest zinc andlead smelter, producing almost 300 kt of zinc annually.

We can calculate the E°cell for the only reaction that could occur.

Zn2�(aq) � 2 e� → Zn(s) E°r � �0.76 V

Pb(s) → Pb2�(aq) � 2 e� E°r � �0.13 V

Zn2�(aq) � Pb(s) → Zn(s) � Pb2�(aq)


� �0.76 V � (�0.13 V)

� �0.63 V

Since the E°cell for the reaction is negative, we conclude that the lead will not be oxidizedspontaneously in the zinc sulfate solution. Note that the reverse reaction would be spon-taneous, but could not occur because neither Pb2�(aq) nor Zn(s) is present initially.

Strictly speaking, the zinc sulfate cell is not a standard cell even if the concentrationof the zinc sulfate were 1.0 mol/L because the cell does not contain all entities listed inthe half-reaction equations. Therefore, the cell potential that is calculated is not accurate,but will be close enough for our purposes. This cell would not produce electricity becausethe reaction is nonspontaneous. Why would anyone be interested in a cell like this?Certainly not to use in a battery. However, by supplying electrical energy to a non-spontaneous cell, we can force the reaction to occur. This is especially useful for pro-ducing substances, particularly elements. For example, the zinc sulfate cell discussedabove is similar to the cell used in the industrial production of zinc metal (Figure 2).


640 Chapter 14 NEL

The term electrochemical cell is often used in chemistry to refer to either a cell with aspontaneous reaction, such as electric and voltaic cells, or a cell with a nonspontaneousreaction, which we call an electrolytic cell (Figure 3). An electrolytic cell consists of acombination of two electrodes, an electrolyte, and an external battery or power source.It uses a process called electrolysis, which is the process of supplying electrical energyto force a nonspontaneous redox reaction to occur. The external power supply acts as an“electron pump;” the electric energy is used to do work on the electrons to cause anelectron transfer inside the electrolytic cell. In an electrolytic cell, the chemical reactionis the reverse of that of a spontaneous cell. However, most of the scientific principlesyou have already studied also apply to electrolytic cells (Table 1).

electric/voltaic cell

reactants products � electrical energyelectrolytic cell

Table 1 Comparing Electrochemical Cells: Voltaic and Electrolytic

Voltaic cell Electrolytic cell

spontaneity spontaneous reaction nonspontaneous reaction

standard cell positive negativepotential, E°cell

cathode • strongest oxidizing agent present • strongest oxidizing agent presentundergoes a reduction undergoes a reduction

• positive electrode • negative electrode

anode • strongest reducing agent present • strongest reducing agent presentundergoes an oxidation undergoes an oxidation

• negative electrode • positive electrode

direction of anode → cathode anode → cathodeelectronmovement

direction of ion anions → anode anions → anodemovement cations → cathode cations → cathode

cathode anode+–

power supply

e–

Figure 3 Electrons are pulled from the anodeand pushed to the cathode by thebattery or power supply.

Secondary Cells: Electric and ElectrolyticA secondary cell is a rechargeable cell such as a nickel–cadmium (Ni–Cd) cell. A secondary cell can be used to illustrate the difference between an electric cell and anelectrolytic cell. As the cell discharges, electrical energy is spontaneously produced andthe cell functions as an electric cell. When the cell is recharged, the electrical energyforces the products to react and re-form the original reactants. During recharging, the secondary cell is functioning as an electrolytic cell.

watertower

lake

pump

houses

Figure 4A town’s water supply

Learning Tip“Positive” and “negative” arelabels that are used in manysituations, mostly decided bygeneral agreement(convention). For example,“positive” and “negative” can beused to label or describeattitudes, directions, axes on agraph, charges, and electrodes.By convention, a voltaic cell hasa cathode labelled positive andan anode labelled negative. Inan electrolytic cell, it iscustomary to reverse theselabels. It is best to think ofpositive and negative forelectrodes as labels, notcharges.

Practice1. Describe the type of agent reacting and the process occurring at the cathode and

anode of an electrolytic cell.

2. Describe the differences between the cathode and anode of an electrolytic cell and avoltaic cell.

3. Describe the direction of movement of electrons and ions within an electrolytic cell.

4. In a town’s water supply (Figure 4), water is pumped from a lake into a water tower,which supplies water to the town's houses. Describe how this is an analogy for asecondary cell.



Section 14.3

The Potassium Iodide Electrolytic Cell: A SynthesisIn the potassium iodide electrolytic cell (Figure 5), litmus paper does not change colourin the initial solution and turns blue only near the electrode from which gas bubbles. Atthe other electrode, a yellow-brown colour and a dark precipitate forms. The yellow-brown substance produces a purplish-red colour in a halogen test. This chemical evidenceagrees with the interpretation supplied by the following half-reaction equations. Accordingto the redox table of relative strengths of oxidizing and reducing agents, water is thestronger oxidizing agent present and iodide ions are the stronger reducing agents presentin a potassium iodide solution.

OA SOA

K�(aq) I�(aq) H2O(l)SRA RA

cathode: 2 H2O(l) � 2 e� → H2(g) � 2 OH�(aq)gas bubbles blue litmus

anode: 2 I�(aq) → I2(s) � 2 e�

purplish-red in hexane

Evidence from the study of this and many other aqueous electrolytic cells suggests thatthe generalizations for voltaic cells also apply to electrolytic cells. From a theoretical per-spective, the strongest oxidizing agent present in a particular mixture has the greatestattraction for electrons and gains electrons at the cathode. Notice that it does not matterwhere the electrons originate, from a power supply or directly from another electrode. Thestrongest reducing agent present in the mixture has the least attraction for electrons andloses electrons at the anode. In other words, the theoretical definitions of cathode andanode are the same for both voltaic and electrolytic cells (Table 1).

Observation of a potassium iodide cell indicates that the transfer of electrons is notspontaneous. When a voltage is supplied to the cell, electrons flow from the negative ter-minal of the battery toward the cathode of the electrolytic cell and are consumed (gained)by water molecules, which have the more positive reduction potential. Simultaneously,electrons flow from iodide ions on the surface of the anode to the positive terminal of thebattery. This explanation agrees with previous redox concepts and with the observations,so we can accept the explanation. Predictions of cathode, anode, and overall cell reactionsfor electrolytic cells follow the same steps outlined for voltaic cells in Section 14.2.



A Potassium Iodide Electrolytic CellElectrolytic cells were discovered before the science wasunderstood. However, as with all successful technologicalinventions, the important criteria was that it worked, not why itworked. Eventually, chemists understood the science and wereable to explain why electrolytic cells work.

In the Evaluation, suggest changes to the Design, Materials,and Procedure that would improve the Evidence.

PurposeThe purpose of this investigation is to use diagnostic tests todetermine the reaction products of an electrolytic cell.

ProblemWhat are the products of the reaction during the operation of anaqueous potassium iodide electrolytic cell?

DesignInert electrodes are placed in a 0.50 mol/L solution of potassiumiodide, and a battery or power supply provides a direct current ofelectricity to the cell. The litmus and halogen diagnostic tests areconducted to test the solution near each electrode before andafter the reaction.


Figure 5A power supply provides the energyfor the chemical reactions at the twoelectrodes.

e—e—

power supply

K+(aq) I—(aq)H2O(l)

carbonelectrode


642 Chapter 14 NEL

What are the cell reactions and the cell potential of the aqueous potassium iodideelectrolytic cell?

First, we identify the major entities in the solution and use the redox table in Appendix I toidentify the strongest oxidizing and reducing agents, as we did in Chapter 13.

OA SOA

K�(aq), I�(aq), H2O(l)SRA RA

Now we can write the half-reaction equations and calculate the cell potential. Thepotassium iodide cell is not a standard cell because the products of the reactions are notpresent initially. Therefore, the reduction potentials given in the table of half-reactions arenot strictly applicable, but we will use them to approximate the cell potential.

cathode: 2 H2O(l) � 2 e� → H2(g) � 2 OH�(aq) E°r � �0.83 V

anode: 2 I�(aq) → I2(s) � 2 e� E°r � �0.54 V

net: 2 H2O(l) � 2 I�(aq) → H2(g) � 2 OH�(aq) � I2(s)


� �0.83 V � (�0.54 V)

� �1.37 V

SAMPLE problem 14.1

A negative sign for a cell potential indicates that the chemical process is nonsponta-neous. The more negative the cell potential, the more energy is required. In SampleProblem 14.1, electrons must be supplied with a minimum of �1.37 V from an externalbattery or other power supply to force the cell reactions. In practice, however, a greatervoltage is required, for example, to make the reaction occur at a reasonable rate.

The procedure for analyzing electrolytic cells is essentially the same as for voltaic cells.

• Use the redox table (Appendix I) to identify the strongest oxidizing and reducing agents present. (Do not forget to consider water for aqueouselectrolytes.)

• Write equations for the reduction (cathode) and oxidation (anode) half-reactions. Include the reduction potentials if required.

• Balance the electrons and write the net cell reaction including the cell potential.


• If required, state the minimum electric potential (voltage) to force the reaction tooccur. (The minimum voltage is the absolute value of E°cell.)

• If a diagram is requested, use the general outline shown in Figure 6 and addspecific labels for chemical entities.

SUMMARYProcedure for Analyzing Electrolytic Cells

Figure 6A generic electrolytic cell

+ — power supply

anode(+)

cathode(—)

e—

cations

electrolyte

anions

e—

Learning TipAt first glance, the cell notation,C(s) ⏐ KI(aq) ⏐ C(s), does notshow that water is involved inthe reaction. Of course, thesubscript (aq) means “dissolvedin water.” Don’t forget toconsider the presence of waterbecause it is often a reactant inaqueous electrolytic cells.



Section 14.3

An electrolytic cell containing cobalt(II) chloride solution and lead electrodes is assembled.The notation for the cell is as follows:

Pb(s) ⏐ Co2�(aq), Cl�(aq) ⏐ Pb(s)

(a) Predict the reactions at the cathode and anode, and in the overall cell.(b) Draw and label a cell diagram for this electrolytic cell, including the power supply.(c) What minimum voltage must be applied to make this cell work?

Solution

(a) SRA SOA

Pb(s) ⏐ Co2�(aq), Cl�(aq) ⏐ Pb(s)

cathode: Co2�(aq) � 2 e� → Co(s)

anode: Pb(s) → Pb2�(aq) � 2 e�

net: Co2�(aq) � Pb(s) → Co(s) � Pb2�(aq)

(b)

(c) E°cell � E°r � E°r

cathode anode

� �0.28 V � (�0.13 V)

� �0.15 V

According to the redox table, a minimum voltage of �0.15 V is required.


+ —

+ —

Pb(s)anode

Pb(s)cathode

Co2+(aq)

Cl—(aq)

power supply

e—e—


644 Chapter 14 NEL

DID YOU KNOW ??HydrometallurgyHydrometallurgy is the process ofextracting metals or theircompounds from ores usingaqueous solutions. Sherritt Metals ofFort Saskatchewan, Alberta is aworld leader in high pressurehydrometallurgy of nickel and cobaltores. Sherritt uses an ammonialeaching of nickel and cobaltsulfides followed by a chemicalreduction using hydrogen as areducing agent. Otherhydrometallurgical operations useelectrolysis instead of a chemicalreduction to obtain or purify metals.

Practice5. Predict the cathode, anode, and net cell reactions for each of the following electrolytic

cells. Calculate the minimum potential difference that must be applied to force thecell reaction to occur. (a) C(s) ⏐ Ni2�(aq), I�(aq) ⏐ C(s) (b) Pt(s) ⏐ Na�(aq), OH�(aq) ⏐ Pt(s)

6. What is the minimum electric potential difference of an external power supply thatproduces chemical changes in the following electrolytic cells? (a) C(s) ⏐ Cr3�(aq), Br�(aq) ⏐ C(s) (b) Cu(s) ⏐ Cu2�(aq), SO4

2�(aq) ⏐ Cu(s)



Electrolysis (Demonstration)Scientific knowledge progresses by the experimental testing ofideas. The more rigorous the test, the more certain the knowledgeor the better the chance of making new discoveries.

PurposeThe purpose of this demonstration is to test the method ofpredicting the products of electrolytic cells.

ProblemWhat are the products of electrolytic cells containing one of thefollowing aqueous solutions: copper(II) sulfate, sodium sulfate,and sodium chloride?

DesignThe electrolysis of the aqueous copper(II) sulfate is carried out ina U-tube, and the electrolysis of aqueous sodium sulfate andsodium chloride is carried out in a Hoffman apparatus so that anygases produced can be collected. Diagnostic tests with necessarycontrol tests (before electrolysis) are conducted to determine thepresence of the predicted products.


An electrolytic cell is set up with a power supply connected to two nickel electrodesimmersed in an aqueous solution containing cadmium nitrate and zinc nitrate. Predict theequations for the initial reaction at each electrode and the net cell reaction. Calculate theminimum voltage that must be applied to make the reaction occur.

Solution

SRA SOA

Ni(s), H2O(l), Cd2�(aq), NO3�(aq), Zn2�(aq)

cathode: Cd2�(aq) � 2 e� → Cd(s) E°r � �0.40 V

anode: Ni(s) → Ni2�(aq) � 2 e� E°r � �0.26 V

net: Cd2�(aq) � Ni(s) → Cd(s) � Ni2�(aq)


� �0.40 V � (�0.26 V)

� �0.14 V

According to the redox table, a minimum voltage of �0.14 V is required.




Section 14.3

Evaluation of Predicted Reactions—The Chloride AnomalyAs you know from Investigation 14.5, some redox reactions predicted using the strongestoxidizing and reducing agents from a redox table do not always occur in an electrolyticcell. Like other chemical processes, a half-cell reaction at an electrode has an activationenergy (Unit 6) that varies for different half-reactions and conditions. Therefore, theactual reduction potential required for a particular half-reaction and the reported half-reaction reduction potential may be quite different. This difference is known as the half-cell overvoltage. It is generally much greater for the production of oxygen than forthe production of chlorine, for example, in the electrolysis of aqueous chloride com-pounds. The explanation for overvoltage is well beyond the level of this book. As anempirical rule, you should recognize that chlorine gas is produced instead of oxygen gas insituations where chloride and water are the only reducing agents present.

Learning TipThere are exceptions to all rulesand generalizations. You onlyneed to remember the chlorideanomaly. This occurs during theelectrolysis of solutionscontaining the chloride ion.Since water is the strongestreducing agent present, watershould react at the anode.However, the chloride ionsreact preferentially to watermolecules.

• An electrolytic cell is based upon a reaction that is nonspontaneous; the E°cell forthe reaction is negative. An applied voltage of at least the absolute value of E°cell isrequired to force the reactions to occur.

• The strongest oxidizing agent undergoes reduction at the cathode (negativeelectrode).

• The strongest reducing agent undergoes oxidation at the anode (positive electrode).

• Electrons are forced by a power supply to travel from the anode to the cathodethrough the external circuit.

• Internally, anions move toward the anode and cations move toward the cathode.

Electrolytic CellsSUMMARY

Practice7. List the main similarities between a voltaic cell and an electrolytic cell.

8. What is the key difference between voltaic and electrolytic cells?

9. Why is the procedure for analyzing voltaic and electrolytic cells so similar?

10. Explain why a power supply is necessary for an electrolytic cell.

11. Which of the following cells would produce a spontaneous reaction? Justify eachanswer, using the cell potential. (a) C(s) ⏐ Cr(NO3)2(aq) ⏐ C(s) (b) Cu(s) ⏐ FeCl3(aq) ⏐ Cu(s)

12. For each of the following electrolytic cells, write equations for the cathode and anodehalf-reactions and the net reaction. Determine the minimum potential difference thatmust be applied to make the cell operate. (a) C(s) ⏐ K2SO4(aq) ⏐ Cd(s) (b) Pt(s) ⏐ SnBr2(aq) ⏐ Pt(s)

13. Draw a diagram of an electrolytic cell containing a zinc iodide solution and inertcarbon electrodes. • Label the power supply and electrodes, including signs, the electrolyte, and the

directions of electron and ion movements.• Write half-reaction and net equations.• Calculate the cell potential, using standard values.

14. State the chloride anomaly, and include how it is recognized.

15. Describe a specific consumer product that you use sometimes as an electric cell andsometimes as an electrolytic cell. What practical problem does this technology solve?


646 Chapter 14 NEL

Science and Technology of ElectrolysisVolta’s invention of the electric cell in 1800 immediately resulted in many discoveries inchemistry. One discovery was that electric cells could be used as an electric power sourcefor electrolytic cells. Many natural substances, such as soda (sodium carbonate) andpotash (potassium carbonate) that were thought to be elements, were shown to be com-posed of the previously unknown elements sodium and potassium. Industrial applicationsof electrolytic cells include the production of elements, the refining of metals, and theplating of metals onto the surface of an object. The study of electrolysis in industry revealsthe strong relationship between science and technology.

Production of ElementsMost elements occur naturally combined with other elements in compounds. Forexample, ionic compounds of sodium, potassium, lithium, magnesium, calcium, andaluminium are abundant, but these reactive metals are not found uncombined in nature.The explanation is that the reduction potentials for these metals are very negative.Consequently, the metals are easily oxidized by practically all other substances. Evenwater has a more positive reduction potential than any of these metal ions. If the metalsdid exist naturally, a spontaneous reaction would convert them into their ions.

SOA Zn2�(aq) � 2 e� 0 Zn(s)

2 H2O(l) � 2 e� 0 H2(g) � 2 OH�(aq)

Mg2�(aq) � 2 e� 0 Mg(s)

Na�(aq) � e� 0 Na(s)

Many metals can be produced by electrolysis of solutions of their ionic compounds,but two difficulties arise. First, many naturally occurring ionic compounds have a lowsolubility in water and second, water is a stronger oxidizing agent than many activemetal cations. To overcome these difficulties, we can use a technological design in whichwater is not present. Fortunately, ionic compounds can be melted. These molten ioniccompounds are good electrical conductors and can function as the electrolyte in a cell.

The production of active metals (strong reducing agents) from their minerals typicallyinvolves the electrolysis of molten compounds of the metal, a technology first used in thescientific work of Humphry Davy (Figure 7). Strontium metal was one of many activemetals discovered by Davy using the electrolysis of molten salts. Strontium chloride wasfirst melted in an electrolytic cell with inert electrodes. In this cell, there are only two kindsof ions present, Sr2�(l) and Cl�(l). You may recall from the previous chapter that metalcations generally tend to undergo a reduction and nonmetal anions tend to undergo anoxidation. In this cell, there are no other competing substances. Therefore, the stron-tium ions will consume (gain) electrons at the cathode to form strontium metal:

Sr2�(l) � 2 e� → Sr(s) (reduction at the cathode)

At the anode, chloride ions will give up (lose) electrons to form chlorine gas:

2 Cl�(l) → Cl2(g) � 2 e� (oxidation at the anode)

The electrons are balanced. Adding the two equations gives the overall reaction in the cell.

Sr2�(l) � 2 Cl�(l) → Sr(s) � Cl2(g)

This reaction would not be possible in an aqueous solution because water is a stronger oxi-dizing agent (has a more positive reduction potential) than aqueous strontium ions.

Figure 7In his youth, Sir Humphry Davy(1778–1829) worked as an assistantto a physician who was interested inthe therapeutic properties of gases.Davy studied nitrous oxide (laughinggas) by conducting experiments onhimself. He was eventually fired fromhis job, supposedly because of hisliking for explosive chemicalreactions. Davy’s main fame camefrom his experiments with electricity.He constructed a voltaic pile withover 250 metal plates. He used thispowerful cell to decompose stable compounds and discovered the elements sodium, potassium, barium,strontium, calcium, and magnesium.Given his habit of tasting, inhaling,and exploding new chemicals, it isnot surprising that he was an invalidin his early thirties and died inmiddle age, probably of chemicalpoisoning.



Section 14.3

In molten-salt electrolysis, metal cations move to the cathode and are reduced to metals,and nonmetal anions move to the anode and are oxidized to nonmetals.

Production of AluminiumAluminium is the third most abundant element on Earth. It was discovered in France inthe early 1800s. At the time, aluminium was more expensive than gold. The wonderfulproperties of aluminium—shiny, light, strong, and corrosion resistant—made it ideal forjewellery and cutlery, so there was a high demand for the metal, especially among the aris-tocracy. However, the supply of aluminium was limited because the technology for pro-ducing aluminium was not yet practical or economically viable for mass production.

Initial efforts to produce aluminium by electrolysis were unproductive because itscommon ore, Al2O3(s), has a high melting point, 2072 °C. No material could be foundto hold the molten compound. In 1886, two scientists working independently andknowing nothing of each other’s work made the same discovery. Charles Hall in theUnited States and Paul Héroult in France discovered that Al2O3(s) dissolves in a moltenmineral called cryolite, Na3AlF6. In this design, the cryolite acts as an inert solvent for theelectrolysis of aluminium oxide and forms a molten conducting mixture with a meltingpoint around 1000 °C. Aluminium (m.p. 660 °C) can be produced electrolytically fromthis molten mixture (Figure 8). This discovery had an immediate effect on the supply andcost of aluminium. Around 1855, aluminium was sold for $45,000 per kilogram; a fewyears after the Hall–Héroult invention, the cost was about 90 cents.

Learning TipNo reduction potentials can belisted for the electrolysis of amolten salt. The redox table inAppendix I lists only electricpotentials for half-reactions in1.0 mol/L aqueous solutions atSATP.

Lithium is the least dense of all metals and is a very strong reducing agent; both qualitiesmake it an excellent anode for batteries. Lithium can be produced by electrolysis of moltenlithium chloride at a temperature greater than 605 °C, the melting point of lithium chloride.Write the equations for the cathode and anode half-reactions, and the net cell reaction.

Solution

cathode: 2 [Li�(l) � e� → Li(s)]

anode: 2 Cl�(l) → Cl2(g) � 2 e�

net: 2 Li�(l) � 2 Cl�(l) → 2 Li(s) � Cl2(g)


Aluminium Production inCanadaThe production of aluminium isimportant to Canada’s economy,although Canada does not havelarge deposits of aluminium ore.Hydroelectric power is used toproduce aluminium metal fromconcentrated, imported bauxite inan electrolytic cell. Recyclingaluminium from soft drink andbeer cans requires only 5% of theenergy required to producealuminium by electrolysis.

DID YOU KNOW ??

C(s) cathode(lining of cell) C(s) anode

Al2O3 in Na3AlF6(l)electrolyte

Al(l)

alumina, Al2O3(s) in hopper

Figure 8The Hall–Héroult cell for theproduction of aluminium. Thecathode is the carbon lining of thesteel cell. At the cathode, aluminiumions are reduced to produce liquidaluminium, which collects at thebottom of the cell and is periodicallydrained away. At the carbon anode,oxide ions are oxidized to produceoxygen gas. The oxygen produced atthe anode reacts with the carbonelectrodes, producing carbondioxide, so these electrodes must bereplaced frequently.


648 Chapter 14 NEL

Aluminium oxide is obtained from bauxite, an aluminium ore. Once the ore is puri-fied, the aluminium oxide is dissolved in molten cryolite and it dissociates into indi-vidual ions. The reactions occurring at the electrodes in a Hall–Héroult cell aresummarized below:

cathode: 4[Al3�(cryolite) � 3 e� → Al(l)]

anode: 3[2 O2�(cryolite) → O2(g) � 4 e�]

4 Al3�(cryolite) � 6 O2�(cryolite) → 4 Al(l) � 3 O2(g)

The overall cell reaction is a decomposition of aluminium oxide.

2 Al2O3(s) → 4 Al(s) � 3 O2(g)

The Chlor–Alkali ProcessThe most important nonmetal produced by electrolysis is chlorine. More than 95% ofthe world production of chlorine and almost 100% of the world production of sodiumhydroxide is done using the chlor–alkali process, which is a reaction that you studiedin Investigation 14.5. The chlor–alkali process is the electrolysis of aqueous sodiumchloride (brine) to produce chlorine, hydrogen, and sodium hydroxide:

2 NaCl(aq) � 2 H2O(l) → Cl2(g) � H2(g) � 2 NaOH(aq)(anode (cathode products)product)

You may recall that the anode reaction in this electrolysis was unexpected in Investigation14.5 because water is a slightly stronger reducing agent than chloride ions. This is the chlo-ride anomaly and is an important exception to the rules for predicting half-reactions.

Dow Chemical in Fort Saskatchewan, Alberta, has an ideal location for its chlor–alkaliplant (Figure 9). The sodium chloride is extracted from large underground deposits usinghot water and pumped to electrolytic cells at the surface. Dow uses diaphragm electrolyticcells, but newer membrane electrolytic cells are becoming more common in the chlor–alkaliindustry. These two technologies are similar, but the membrane cell has the advantage ofproducing sodium hydroxide with very little sodium chloride contamination. This isaccomplished by using an ion-exchange membrane that allows sodium ions, but notchloride ions, to move from the anode to the cathode compartments (Figure 10).

Hydrogen gas is used to make ammonia, hydrogen peroxide, and margarine, and tocrack petroleum. It may also be used on site as a fuel to produce electricity. Chlorine isused as a disinfectant for drinking water and to manufacture bleach (sodium hypochlo-rite), plastics, pesticides, and solvents. Sodium hydroxide is used on a large scale inindustry to make cellophane, pulp and paper, aluminium, and detergents.

Figure 9Large salt deposits sit hundreds ofmetres below Fort Saskatchewan,Alberta. This means that FortSakatchewan is a good location forDow’s chlor–alkali plant becausethe main raw material is available onsite.

Figure 10 Membrane cells are a relativelyrecent technology. Most companiesthroughout the world are switchingto these cells from the olderdiaphragm and mercury cells.

+ —

NaOH(aq)

H2O(l)

NaCl(aq)24%

H2(g)Cl2(g)

NaCl(aq)26%

cathode

power supply

membrane

anode

Aluminium Energy Source Laptops and even some cars of thefuture may use aluminium as anenergy source. It takes a greatdeal of electrical energy toproduce aluminium and some ofthis energy can be recovered. TwoUniversity of British Columbiaresearchers have devised a way toproduce hydrogen from aluminiumusing only salt and water. Thehydrogen that is produced can beused to power a fuel cell forportable electronics or even a car.


EXTENSION +



Section 14.3

Practice16. Describe two difficulties associated with the electrolysis of aqueous ionic compounds

in the production of active metals. What is a technological solution that overcomesboth difficulties?

17. Scandium is a metal with a low density and a melting point that is higher than that ofaluminium. These properties are of interest to engineers who design space vehicles.Scandium metal is produced by electrolysis of molten scandium chloride. List all ionspresent in the electrolysis cell, and write the equations for the reactions that occur atthe cathode and anode, and the net cell reaction.

18. Write the cathode, anode, and cell reaction equations for the chlor–alkali process.

19. The following statements summarize the steps in the chemical technology ofobtaining magnesium from seawater. Write a balanced equation to represent eachreaction.(a) Slaked lime (solid calcium hydroxide) is added to seawater (ignore all solutes

except MgCl2(aq)) in a double displacement reaction to precipitate magnesiumhydroxide.

(b) Hydrochloric acid is added to the magnesium hydroxide precipitate.(c) After the magnesium chloride product is separated and dried, it is melted in

preparation for electrolysis. List all ions present in the electrolysis, and write theequations for the reactions that occur at the cathode and anode, and the net cellreaction.

(d) An alternative process produces magnesium from dolomite, a mineral containingCaCO3 and MgCO3. Suggest some technological advantages and disadvantagesof the dolomite process compared with the seawater process.

20. What products in your home may have originated from substances produced in thechlor–alkali process?

21. Why should we recycle metals such as aluminium? State several arguments that youmight use in a debate.

The Pidgeon ProcessLloyd Pidgeon (1903–1999) wasborn in Markham, Ontario, studiedundergraduate chemistry at theUniversity of Manitoba, andobtained his Ph.D. from McGillUniversity in 1929. Later, whileworking at the National ResearchCouncil in Ottawa, Pidgeondeveloped the first process forproducing high-qualitymagnesium metal from dolomite(calcium magnesium carbonate).This led to the formation ofDominion Magnesium Ltd. usingdolomite mined in the OttawaValley. The Pidgeon Process is stillused in many countries to producemagnesium.

DID YOU KNOW ??

Refining of MetalsIn the production of metals, the initial product is usually an impure metal. Impuritiesare often other metals that come from various compounds in the original ore. To purifyor refine a metal, a variety of methods are used. However, a common method, known aselectrorefining, uses an electrolytic cell to obtain high-grade metals at the cathode froman impure metal at the anode.

A good example is the electrorefining of copper. The presence of impurities in copperlowers its electrical conductivity, not a desirable property considering that one of themost common uses of copper is in electrical wiring. The initial smelting process producescopper that is about 99% pure, containing some silver, gold, platinum, iron, and zinc. Thesevaluable impurities can be recovered and sold to help pay for the process. As shown inFigure 11, a slab of impure copper is the anode of an electrolytic cell that containscopper(II) sulfate dissolved in sulfuric acid. The cathode is a thin sheet of very purecopper. As the cell operates, copper and some of the other metals in the anode are oxi-dized, but only copper is reduced at the cathode. An understanding of oxidation, reduc-tion, and reduction potentials allows precise control over what is oxidized and what isreduced; after electrorefining, the copper is about 99.98% pure. The half-reactions are:

cathode: reduction of copper Cu2�(aq) � 2 e� → Cu(s)

anode: oxidation of copper Cu(s) → Cu2�(aq) � 2 e�

oxidation of zinc Zn(s) → Zn2�(aq) � 2 e�

oxidation of iron Fe(s) → Fe2�(aq) � 2 e�

Cu2+(aq)impurecopper

purecopper

cathode anode

e—

electrolyte(CuSO4(aq),H2SO4(aq))

sludge(Ag, Au, Pt)

e—

— +power supply

Figure 11Only copper and metals more easilyoxidized than copper, such as ironand zinc, are oxidized to ions anddissolve at the anode. Only copper isreduced at the cathode. Otherimpurities in the anode, such assilver, gold, and platinum, do notreact; these fall to the bottom of thecell as a sludge called anode mud.


650 Chapter 14 NEL

Another method of purifying metals is to reduce metal cations from a molten oraqueous electrolyte at the cathode of an electrolytic cell, much like the production of ele-ments discussed previously. This method, which uses a molten salt, is known as elec-trowinning. It is the only way to obtain some active metals, such as those in Group 1. Manyother metals, such as zinc, can be produced by electrowinning an aqueous solution. Forexample, Cominco’s operation at Trail, British Columbia, uses the electrolysis of anacidic zinc sulfate solution with a specially treated lead anode to deposit very pure zincmetal at the cathode.

cathode: 2[Zn2�(aq) � 2 e� → Zn(s)]

anode: 2 H2O(l) → O2(g) � 4 H�(aq) � 4 e�

net: 2 Zn2�(aq) � 2 H2O(l) → 2 Zn(s) � O2(g) � 4 H�(aq)

ElectroplatingSeveral metals, such as silver, gold, zinc, and chromium, are valuable because of theirresistance to corrosion. However, products made from these metals in their pure formare either too expensive or they lack suitable mechanical properties, such as strengthand hardness. To achieve the best compromise among price, mechanical properties,appearance, and corrosion resistance, utensils and jewellery may be made of a relativelyinexpensive, yet strong alloy such as steel, and then coated (plated) with another metalor alloy to improve appearance or corrosion resistance. Plating of a metal at the cathodeof an electrolytic cell is called electroplating and is a common technology that is used tocover the surface of an object with a thin layer of metal.

The process for plating metals is obtained by systematic trial and error, involving thecareful manipulation of one possible variable at a time. In this situation, a scientific per-spective helps identify variables, but cannot usually provide successful predictions.

The development and use of electric cells preceded scientific understanding of theprocesses involved. Today, we still have examples of technological processes that are notfully understood, such as chromium plating (Figure 12) and silver plating. For example,there is no satisfactory explanation for why silver deposited during electrolysis of a silvernitrate solution does not adhere well to any surface, whereas silver plated from silvercyanide solution does.

Electroplating is only one method used to cover the surface of an object with a metal.One other method is vapour deposition, in which metal vapour is condensed on thesurface. Another method is dipping, in which an object is dipped into a molten metal thatsolidifies on the surface. Zinc-plated nails for exterior use are made by dipping.

Figure 12Chromium is best plated from asolution of chromic acid. A thin layerof chromium metal is very shiny and,like aluminium, protects itself fromcorrosion by forming a tough oxidelayer.

• In molten-salt electrolysis, metal cations are reduced to metal atoms at thecathode and nonmetal anions are oxidized at the anode.

• Electrorefining is a process used to obtain high-grade metals at the cathode froman impure metal at the anode.

• Electroplating is a process in which a metal is deposited on the surface of anobject placed at the cathode of an electrolytic cell.

Applications of Electrolytic CellsSUMMARY

CAREER CONNECTION

Chemical TechnologistChemical technologists may workclosely with scientists and engineers studying electroplatingprocesses. What education andtraining is required for this job?Outline some current job opportunities in this area,including typical salaries.




Section 14.3

Section 14.3 Questions1. Define an electrolytic cell.

2. List the key similarities and differences in the typicallaboratory construction of a voltaic cell and an electrolyticcell.

3. The rules for predicting the chemical reactions, and ion andelectron movements are essentially the same for bothvoltaic and electrolytic cells. What is different about the cellreaction equations and characteristics of each type of cell?

4. Two tin electrodes are placed in an aqueous solutioncontaining potassium nitrate and magnesium iodide. (a) If a power supply is connected to force any reactions to

occur, what would be the reactions at the cathode,anode, and in the overall cell?

(b) Draw and label a cell diagram, including electrodes,electrolyte, power supply, and the direction ofmovement of electrons and ions.

(c) Would a 1.5 V cell be suitable as a power supply? Justifyyour answer.

5. For the electrolysis of aqueous solutions, describe thecommon exception to the rules for predicting products ofan electrolytic cell. Clearly identify the circumstances whenthis exception is used.

6. List three uses of electrolytic cells in industry.

7. Why were many metals discovered only after the inventionof the electric cell?

8. How does the occurrence of metals in nature relate to theredox table?

9. Draw and label a simple cell for the electrolysis of moltenpotassium iodide (m.p. 682 °C). Label electrodes and powersupply, directions of electron and ion flow, and write half-reaction and net equations.

10. When refining metals in an electrolytic cell, why must themetal product form at the cathode?

11. High-purity copper metal is produced using electrorefining. (a) At which electrode is the impure copper placed? Why?(b) What is the minimum electric potential difference

required for this cell?(c) Why is it unlikely that your answer to (b) would be

used? Discuss briefly.

12. How can you predict which metals might be refined froman aqueous solution?

13. Describe the relationship of science and technology in thearea of electrolysis. Include several examples in yourdescription.

14. “German silver” is an alloy containing copper, zinc, andnickel. A piece of German silver is used as the anode in anelectrolytic cell containing aqueous sodium sulfate. Theother electrode is platinum metal.(a) As the applied voltage is slowly increased, in what

order will the half-reactions occur at the anode? Writean equation for each half-reaction.

(b) Describe what happens at the cathode.

(c) German silver does not contain any silver metal. Why isit called German silver? Why is it a very useful alloy?

15. Design a cell to electroplate zinc onto an iron spoon. Inyour cell diagram, include: • ions in the solution• substances used for the electrodes• anode and cathode labels• power supply, showing signs and connections• directions of ion and electron flow

16. Suppose you work for a mining company and you are givena job to design a process that will recover nickel metal froma waste solution containing nickel(II) ions. (a) Propose an experimental design involving electrolysis

that could be tested in the laboratory on a small scale.(b) What are some possible complications or factors that

need to be considered? List these as questions.

Extension

17. The one-dollar coin, or theloonie (Figure 13) replacedthe one-dollar bill, whichtypically wore out in a fewmonths. Sherritt Gordon ofFort Saskatchewan, Albertadeveloped a unique processfor plating the loonie coin. (a) Research the production

and composition of theloonie.

(b) What is the golden“aureate” finish on theloonie? Describe thematerials and process forproducing this finish.

(c) Why did the coin end upwith a loon stamped on it?

18. Aluminium cans are widely used to contain beverages. Writea short report about the production of aluminium cans,including how the can is made, how the top is attached tothe can, how the construction of the can has changed sincethe first model, and the advantages of using recycledaluminium instead of new aluminium.

19. “Cold fusion” made front-page news when it wasannounced in 1989 as a new practical source of energy.(a) Briefly describe cold fusion.(b) How is cold fusion related to electrochemistry?(c) Outline some theoretical and empirical arguments in

the controversy about the existence of cold fusion.


Figure 13The Royal Canadian Mintin Winnipeg produces theCanadian loonie forgeneral circulation.




652 Chapter 14 NEL

14.414.4 Cell StoichiometryIn the production of elements, the refining of metals, and electroplating, the quantity ofelectricity that passes through a cell determines the masses of substances that react or areproduced at the electrodes. As you know from oxidation and reduction half-reactions,a specific number of electrons are lost or gained. For example, when zinc is plated ontoa steel pipe to galvanize it, two moles of electrons must be gained by one mole of zincions to deposit one mole of zinc atoms as metal.

Zn2�(aq) � 2 e� → Zn(s)

As in all stoichiometry, this relationship establishes a mole ratio of electrons to someother substance in the half-reaction equation. Unfortunately, there is no meter or instru-ment for measuring directly (or counting directly) the number of electrons. The numberof electrons (as moles of electrons) is determined indirectly. In the past, you have meas-ured mass and then converted to a chemical amount of a substance; a similar proce-dure is necessary for determining the amount of electrons.

Before we can look at the amount of electrons, we need to see how the charge is deter-mined. Charge, Q, in coulombs, is determined from the electric current, I, in amperes(coulombs per second), and the time, t, in seconds, according to the following definition:

Q � It

One coulomb (C) is the quantity of charge transferred by a current of one ampere(A) during a time of one second

The technology of the Hall–Héroult cell for producing aluminium has improvedconsiderably since the first industrial factory. Modern electrolytic cells may use up to 300 kA of current. What is the charge that passes through one of these cells in a 24 hperiod?

By definition, a current in amperes (A) is the number of coulombs per second, 1 A � 1 C/s.You always need to convert the time into seconds before time can be used in the calculation of charge.

t � 24 h� � �36

10h�0 s� � 8.6 � 104 s

Now the charge in coulombs can be calculated as follows:

Q � It

� 300 � 103 �Cs�� 8.6 � 104 s�

� 2.6 � 1010 C

Therefore, a current of 300 kA for 24 h transfers 2.6 � 1010 C of charge. This is a hugequantity of charge. For comparison, the charge passing through a 100 W light bulb in 24 his about 7.2 � 104 C.

SAMPLE problem 14.2DID YOU KNOW ??Science versus TechnologyJack Kilby (1923–2005) was a juniorengineer when he invented themicrochip that is the foundation ofall modern electronics. He wasawarded a Nobel Prize for Physics in2000. On the topic of his prize, Kilbycommented:

“Those big prizes are for theadvancement of understanding.They are for scientists who aremotivated by pure knowledge.But I'm an engineer. I’mmotivated by a need to solveproblems, to make somethingwork. For guys like me, the prizeis seeing a successful solution.”



Faraday’s LawThe relationship between electricity and electrochemical changes was first investigatedby Michael Faraday (Figure 1) in the 1830s. Faraday continued Humphry Davy’s workin electrochemistry, coining the terms electrolysis, electrolyte, electrode, anode, cathode,cation, and anion. His quantitative study of electrolysis identified the factors that deter-mine the mass of an element produced or consumed at an electrode. He discovered thatthis mass was directly proportional to the time the cell operated, as long as the currentwas constant (Faraday’s law). Furthermore, he found that 9.65 � 10 4 C of charge istransferred for every mole of electrons that flows in the cell. In modern terms, this value isthe molar charge of electrons, also called the Faraday constant, F.

F � 9.65 � 104 �mo

C

l e��

This constant can be used as a conversion factor in converting electric charge to anamount in moles of electrons—in the same way that molar mass is used to convert massto a chemical amount.

ne� �

Since Q � It, the amount of electrons can now be written as

ne� �It�F

Q�F

Section 14.4

Practice1. Calculate the charge transferred by a current of 1.5 A flowing for 30 s.

2. In an electrolytic cell, 87.6 C of charge is transferred in 22.5 s. Determine the electriccurrent.

3. Calculate the charge transferred by a current of 250 mA in a time of 28.5 s.

4. How long, in minutes, does it take a current of 1.60 A to transfer a charge of 375 C?

What amount of electrons is transferred in a cell that operates for 1.25 h at a current of0.150 A?

Recall from the calculation of electric charge that the time must always be in secondsbecause the ampere is defined as coulombs per second (1 A � 1 C/s).

t � 1.25 h� � � 4.50 � 103 s

Now you can calculate the amount, in moles, of electrons using the Faraday constant as aconversion factor:

ne� � (0.150 �C�s�� 4.50 � 103 s�) �

� 6.99 � 10�3 mol

The amount of electrons transferred is 6.99 mmol.

1 mol��9.65 � 104 C�

3600 s�

1 h�

SAMPLE problem 14.3

Figure 1As a young man, Michael Faraday(1791–1867) taught himselfchemistry and convinced the famousEnglish chemist Humphry Davy tohire him as his assistant. Faradayproved himself more than worthy ofthis trust. He eventually madeimportant contributions in the studyof gases, low temperatures, thediscovery of benzene, quantitativeaspects of electrolysis, electricmotors, generators, andtransformers. A deeply religiousman, Faraday had strong convictionsabout the appropriate uses ofscience and technology. He refusedto help Britain produce a poison gasfor use against the Russians in theCrimean War (1854–56).


654 Chapter 14 NEL

The same method can be used to calculate electric current or time if the other vari-ables are known, as shown in Communication Example 2.

Learning TipNote the similarity between various definitions of amountsusing molar quantities:

n substance � �Mm

�

ngas � �Vv

�

ne� � �QF

�

Practice5. An electroplating cell operates for 35 min with a current of 1.9 A. Calculate the

amount, in moles, of electrons transferred.

6. A cell transferred 0.146 mol of electrons with a constant current of 1.24 A. How long,in hours, did this take?

7. Calculate the current required to transfer 0.015 mol of electrons in 20 min.

Convert a current of 1.74 A for 10.0 min into an amount of electrons.

Solution

t � 10.0 min� � �16m0

isn�

� � 600 s

ne� � (1.74 �C�s�� 600 s�) �

� 0.0108 mol

According to Faraday’s law, the amount of electrons transferred is 0.0108 mol or10.8 mmol.

1 mol��9.65 � 104 C�


How long, in minutes, will it take a current of 3.50 A to transfer 0.100 mol of electrons?

Solution

t � (0.100 mol� � 9.65 � 104 �m

C�ol��) � �

3.510s

C��

� 2.76 � 103 s� � �16m0

is�n

�

� 46.0 min

According to Faraday’s law, it would require 46.0 min to transfer 0.100 mol e� using acurrent of 3.50 A.

COMMUNICATION example 2Learning TipAn alternative solution is to usethe mathematical formula andsolve for time, t.

ne� �

t �

You can either substitute valuesfirst and then solve for time, orrearrange the mathematicalformula and then substitute thevalues. Either way is acceptable.

ne�F�

I

It�F



Section 14.4

Half-Cell CalculationsSince the mass of an element produced at an electrode depends on the amount of trans-ferred electrons, a half-reaction equation showing the number of electrons involved isnecessary to do stoichiometric calculations. This applies to all electrochemical cells,whether voltaic or electrolytic. Separate calculations are carried out for each electrode,although the same charge and, therefore, the same amount of electrons passes through eachelectrode in a cell or a group of cells in series. As the following examples show, conceptsof stoichiometry used in other calculations also apply to half-cell calculations. The onlynew part of the stoichiometry is the calculation of the amount of electrons based onthe Faraday constant.

What is the mass of copper deposited at the cathode of a copper electrorefining cell(Figure 2) operated at 12.0 A for 40.0 min?

First identify and write the appropriate half-cell equation. Because copper is beingdeposited at the cathode, copper( II) ions must be gaining electrons to form copper metal.Write the equation for this reduction and list all information given, including constantssuch as molar mass and Faraday.

Cu2�(aq) � 2 e� → Cu(s)

40.0 min m12.0 A 63.55 g/mol9.65 � 104 C/mol

Notice that we have all of the information necessary to calculate the amount of electrons.Don’t forget to make sure the time is converted to units of seconds, if necessary.

ne� � (12.0 �C�s� � 40.0 min� � �

16m0

is�n��) �

� 0.298 mol

The procedure that is common to all stoichiometry is the use of the mole ratio from abalanced equation. The mole ratio is what allows us to convert from a chemical amount inmoles of one substance to another. In the reduction half-reaction given, notice that 1 molof copper metal is formed when 2 mol of electrons are transferred.

nCu � 0.298 mol � �12

�

� 0.149 mol

The final step is to convert to the quantity requested in the question, in this case, themass of copper metal.

mCu � 0.149 mol� �

� 9.48 g

or

mCu � 40.0 min� � �16m0

is�n�

� � 12.0 �C1es�

�

� � � �

� 9.48 g

According to Faraday’s law and the stoichiometric method, the mass of copper metaldeposited is 9.48 g.

63.55 g Cu��1 mol Cu

1 mol Cu�2 mol e�

1 mol e�

��9.65 � 104 C e�

63.55 g�

mol�

1 mol��9.65 � 104 C�

SAMPLE problem 14.4

Figure 2These pure copper cathodes fromthe electrolytic refining cells areready for melting and processinginto copper wire and other copperproducts.


656 Chapter 14 NEL

Learning TipNote the similarity of theprocedure for stoichiometrycalculations of half-cells to allother stoichiometry calculationsyou have done in the past.Essentially, the only differenceis a new relationship (formulabased on Faraday’s law) to convert to and from the amountof electrons.

Step 1: Write the balanced equation for the half-cell reaction of the substance pro-duced or consumed. List the measurements and conversion factors for the givenand required entities.

Step 2: Convert the given measurements to an amount in moles by using the appropriateconversion factor (M, c, F).

Step 3: Calculate the amount of the required substance by using the mole ratio from thehalf-reaction equation.

Step 4: Convert the calculated amount to the final quantity by using the appropriate conversion factor (M, c, F).

Procedure for Half-Cell StoichiometrySUMMARY

Silver is deposited on objects (Figure 3) in a silver electroplating cell. If 0.175 g of silver isto be deposited from a silver cyanide solution in a time of 10.0 min, predict the currentrequired.

Solution

Ag�(aq) � e� → Ag(s)

10.0 min 0.175 g

I 107.87 g/mol

9.65 � 104 C/mol

nAg � 0.175 g� �

� 1.62 � 10�3 mol

ne� � 1.62 � 10�3 mol �

� 1.62 � 10�3 mol

I �

� 0.261 C/s

or

I � 0.175 g Ag� � � � � �10

1min��

� 0.261 C/s

According to the stoichiometry and Faraday’s law, the current required to plate 0.175 gof silver in 10.0 min is 0.261 A.

1 min��60 s

9.65 � 104 C e�

��1 mol e�

1 mol e�

��1 mol Ag

1 mol Ag��107.87 g Ag�

1.62 � 10�3 mol� � 9.65 � 104 �mCol��

��

10.0 min� � �6m0in�s

�

1�1

1 mol�107.87 g�


Figure 3A silver electroplating cell usessilver cyanide to silver plate objects,such as this tray.



Section 14.4

Section 14.4 Questions1. A battery delivers 0.300 A for 15.0 min. What amount of

electrons, in moles, is transferred?

2. A current of 55 kA passes through a chlor–alkali cell. Whatmass of chlorine is formed during 8.0 h ?

3. A family wishes to plate an antique teapot with 10.00 g ofsilver. If the current to be used is 1.80 A, what length oftime, in minutes, is required?

4. A typical Hall–Héroult cell produces 425 kg of moltenaluminium in 24.0 h. Calculate the current used.

5. Magnesium metal is produced in an electrolytic cellcontaining molten magnesium chloride. A current of2.0 � 105 A is passed through the cell for 18.0 h. (a) Determine the mass of magnesium produced.(b) What mass of chlorine is produced at the same time?

6. Cobalt metal is plated from 250.0 mL of cobalt(II) sulfatesolution. What is the minimum concentration of cobalt(II)sulfate required for this cell to operate for 2.05 h with a current of 1.14 A?

7. A 25.72 g piece of copper metal is the anode in a cell inwhich a current of 0.876 A flows for 75.0 min. Determine thefinal mass of the copper electrode.

8. A student reconstructs Volta’s electric battery using sheetsof copper and zinc, and a current of 0.500 A is produced for10.0 min. Calculate the mass of zinc oxidized to aqueouszinc ions.

9. Electroplating is a common technological process forcoating objects with a metal to enhance the appearance ofthe object or its resistance to corrosion. (a) A car bumper is plated with chromium using

chromium(III) ions in solution. If a current of 54 A flowsin the cell for 45 min 30 s, determine the mass ofchromium deposited on the bumper.

(b) For corrosion resistance, a steel bolt is plated withnickel from a solution of nickel(II) sulfate. If 0.250 g ofnickel produces a plating of the required thickness anda current of 0.540 A is used, predict how long inminutes the process will take.

10. During the electrolysis of molten aluminium chloride in anelectrolytic cell, 5.40 g of aluminium is produced at thecathode. Predict the mass of chlorine produced at theanode.

11. Chromium metal can be plated onto an object from anacidic solution of dichromate ions. What average current isrequired to plate 17.8 g of chromium metal in a time of2.20 h? (You will need to construct your own equation forthe half-reaction.)

12. The purpose of this experiment is to test the method ofstoichiometry in cells. Complete the Prediction, Analysis,and Evaluation (Part 2 only) of the investigation report.

ProblemWhat is the mass of tin electroplated at the cathode of atin-plating cell by a current of 3.46 A for 6.00 min?

DesignA steel can is placed in an electroplating cell as thecathode. An electric current of 3.46 A flows through thecell, which contains a 3.25 mol/L solution of tin(II) chloride, for 6.00 min.

Evidenceinitial mass of can � 117.34 gfinal mass of can � 118.05 g

13. Using a specific example of an electrolytic cell, describehow Faraday’s law is useful in designing and controllingthe process.

Extension

14. A rapidly developing technology is the production of lessexpensive, more durable, and more energy-denseelectrochemical cells, that is, cells with a high energy-to-mass ratio. (a) A car battery has a rating of 125 A•h (ampere-hours).

What does this tell you about the electrical capacity ofthis battery?

(b) Why is this a useful way to rate batteries?(c) What mass of lead is oxidized as this battery

discharges?(d) If an aluminium–oxygen fuel cell has the same rating as

the car battery in (a), what mass of aluminium metalwould be oxidized?

(e) Comment on the implications of your answers to (c)and (d).

WEB Activity

Simulation—Electrolytic Cell StoichiometryIn this activity, you will simulate the operation of an electrolytic cell. You will follow a series ofsteps to calculate the approximate cost of producing a given mass of aluminium. Thesecalculations are similar to ones that scientists and engineers do when designing andmonitoring an industrial cell.



Chapter 14 INVESTIGATIONS

658 Chapter 14 NEL

Designing an Electric Cell

In this cell, an aluminium soft-drink can is one of the elec-trodes (Figure 1). The other electrode is a solid conductor,such as a piece of copper wire or pipe, an iron nail, or graphitefrom a pencil. The electrolyte may be a salt solution, or anacidic or basic solution. Although many characteristics of a cellare important, only one characteristic, voltage, is investigatedhere. Check with your teacher if you want to evaluate otherdesigns and materials.

When evaluating the Purpose (Part 3), include your opinionabout the reliability, cost, and simplicity of your final elec-tric cell.

PurposeThe purpose of this investigation is to make an electric cell.

ProblemWhat combination of electrodes and electrolyte gives thelargest voltage for an aluminium-can cell?

DesignPart 1Using the same electrolyte and aluminium can as the con-trolled variables, two or three different materials are used asthe second electrode. The voltage of each cell is measured asthe responding variable.

Part 2Using the same two electrodes as the controlled variables,two or three possible electrolytes are tested. The voltage ofeach cell is measured as the responding variable.

Part 3Additional combinations are tested, based on the analysis ofthe initial trials.



Figure 1An aluminium-can cell is an efficient design since one of theelectrodes also serves as the container.

voltmeter

cathode

electrolyte

aluminiumcan anode

+ –

Be careful when handling acidic and basic solutionsused for electrolytes, as they are corrosive. Wear eyeprotection and work near a source of water. Someelectrolytes may be toxic or irritant; follow all safetyprecautions. Avoid eye and skin contact.



Chapter 14

A Voltaic Cell (Demonstration)

An important characteristic of consumer, commercial, andindustrial cells is a simple, efficient design that works for theintended application. For scientific research, this is not asimportant as a design that can be easily manipulated andstudied.

PurposeThe purpose of this investigation is to test the design andoperation of a voltaic cell used in scientific research.

ProblemWhat is the design and operation of a voltaic cell?

DesignAn electric cell with only one electrolyte is compared with voltaic cells containing the same electrodes, but two electrolytes.

Procedure 1. Construct the three cells shown in Figure 2.

2. For each design, use a voltmeter to determine whichelectrode is positive and which is negative (seeAppendix C.3), and measure the electric potentialdifference of each cell.

3. With the voltmeter connected, remove and thenreplace the various parts of the cell.

4. For each cell, connect the two electrodes with a wire.Record any evidence of a reaction after severalminutes, and after one or two days. Measure theelectric potential difference after several days.



no porous boundary:

salt bridge:

Cu(s)⏐Cu(NO3)2(aq)⏐⏐AgNO3(aq)⏐Ag(s)

Cu(s)⏐Cu(NO3)2(aq)⏐⏐AgNO3(aq)⏐Ag(s)Cu(s)⏐NaNO3(aq)⏐Ag(s)

porous cup:

Figure 2In Investigation 14.2 , you compare three different cell designs.

Solutions used are toxic and irritant.Avoid contact with skin and eyes.


660 Chapter 14 NEL

Testing Voltaic Cells

Testing is a procedure that is common to both technologyand science. In technology, testing is necessary to determinehow a product or process works using criteria such as effi-ciency, reliability, and cost. In science, testing is a key part inthe advancement of knowledge. Scientific concepts are devel-oped and then tested to determine their validity and limita-tions. New ideas that fail the test then need to be restricted,revised, or replaced.

In your Evaluation, pay particular attention to sources of error or uncertainty, and to limitations of the evidencecollected.

PurposeThe purpose of this investigation is to test the predictions ofcell potentials and the identity of the electrodes of variouscells.

ProblemIn cells constructed from various combinations of copper,aluminium, silver, and zinc half-cells, what are the standardcell potentials, and which is the anode and cathode in eachcase?

Materialslab aproneye protectionvoltmeter and connecting wiresU-tube with cotton plugs, porous cups, or filter paperfour 100 mL beakers (or well plate)distilled watersteel woolCu(s), Al(s), Ag(s), and Zn(s) strips1.0 mol/L each of:

CuSO4(aq)Al(NO3)3(aq)AgNO3(aq)NaNO3(aq) ZnSO4(aq)



The materials used are toxic and irritant.Avoid skin and eye contact.



Chapter 14

A Potassium Iodide Electrolytic Cell

Electrolytic cells were discovered before the science was under-stood. However, as with all successful technological inven-tions, the important criteria was that it worked, not why itworked. Eventually, chemists understood the science and wereable to explain why electrolytic cells work.

In the Evaluation, suggest changes to the Design, Materials,and Procedure that would improve the Evidence.

PurposeThe purpose of this investigation is to use diagnostic tests todetermine the reaction products of an electrolytic cell.

ProblemWhat are the products of the reaction during the operationof an aqueous potassium iodide electrolytic cell?

Design Inert electrodes are placed in a 0.50 mol/L solution of potas-sium iodide, and a battery or power supply provides a directcurrent of electricity to the cell. The litmus and halogen diagnostic tests (Appendix C.4) are conducted to test thesolution near each electrode before and after the reaction.

Materialslab aproneye protectionpetri dishtwo carbon electrodestwo connecting wires3 V to 9 V battery or power supplyred and blue litmus paperring stand and two utility clampssmall test tube with stopperdropper bottle of hexane0.50 mol/L KI(aq)

Procedure 1. Set up the KI(aq) cell as shown in Figure 3, (or with

one ring stand), but with a single wire connecting theelectrodes (i.e., no power supply).

2. Observe the cell and test the solution with litmuspaper and hexane.

3. Use two connecting wires to hook up the powersupply to the electrodes.

4. Turn on the power supply.

5. Record all observations at each electrode.

6. Perform both diagnostic tests at each electrode.

7. Dispose of the solutions by putting the hexanemixture in a labelled waste container and washing thepotassium iodide solution down the sink drain.



Hexane is highly flammable. Do not use near an openflame and avoid inhaling the fumes. Use only in awell-ventilated area. Figure 3

Apparatus for the electrolysis of potassium iodide

e–e–

power supply

voltagesetting

petri dish

carbonelectrode

carbonelectrode

K+(aq) I_(aq)

H2O(l)


662 Chapter 14 NEL

Electrolysis (Demonstration)

Scientific knowledge progresses by the experimental testingof ideas. The more rigorous the test, the more certain the knowledge or the better the chance of making new discoveries.

PurposeThe purpose of this demonstration is to test the method ofpredicting the products of electrolytic cells.

ProblemsWhat are the products of electrolytic cells containing one ofthe following solutions:

• aqueous copper(II) sulfate?

• aqueous sodium sulfate?

• aqueous sodium chloride?

DesignThe electrolysis of the aqueous copper(II) sulfate is carried outin a U-tube, and the electrolysis of aqueous sodium sulfate andsodium chloride is carried out in a Hoffman apparatus(Figure 4) so that any gases produced can be collected.Diagnostic tests with necessary control tests (before elec-trolysis) are conducted to determine the presence of the pre-dicted products.



Figure 4Hoffman apparatus

Copper(II) sulfate is toxic and irritant. Avoidskin and eye contact. If you spill copper(II)sulfate solution on your skin, wash theaffected area with lots of cool water. Duringelectrolysis, corrosive substances areproduced; avoid skin and eye contact.


Chapter 14 SUMMARY


Chapter 14

Outcomes

Knowledge

• define anode, cathode, anion, cation, salt bridge/porous cup,electrolyte, voltaic cell, and electrolytic cell (14.2, 14.3)

• predict and write the half-reaction equation that occurs ateach electrode in an electrochemical cell (14.2, 14.3, 14.4)

• explain that the values of standard reduction potential are allrelative to E°r � 0.00 V set for the hydrogen electrode atstandard conditions (14.2)

• calculate the standard cell potential for electrochemical cells(14.2, 14.3)

• predict the spontaneity of redox reactions based on standardcell potentials (14.2, 14.3)

• identify the similarities and differences between a voltaic celland an electrolytic cell (14.3)

• recognize that predicted reactions do not always occur (14.3)

• calculate mass, amounts, current, and time in single voltaicand electrolytic cells by applying Faraday’s law andstoichiometry (14.4)

STS

• state examples of science leading technology andtechnology leading science (14.1, 14.3)

• recognize the values and limitations of technologicalproducts and processes (14.1, 14.3)

• describe the interactions of science and technology (14.1,14.2, 14.3)

Skills

• initiating and planning: design an experiment, including alabelled diagram, to test predictions for reactions occurringin electrochemical cells (14.2, 14.3); describe procedures forsafe handling, storage, and disposal of materials used in thelaboratory (14.1, 14.2); and develop a plan to build a cell(battery) using a trial-and-error procedure (14.1)

• performing and recording: construct and observeelectrochemical cells (14.1, 14.2, 14.3); investigate the issue ofdisposal of batteries (14.1); and compile and displayinformation about electrochemical cells (all sections)

• analyzing and interpreting: identify the products ofelectrochemical cells (all sections); compare predictions withobservations (all sections); identify the limitations ofevidence collected (14.1, 14.2, 14.3); explain discrepanciesbetween predicted and measured cell potentials (14.2, 14.3);assess the practicalities of different cell designs (14.1); andevaluate experimental designs for cells (14.1, 14.2, 14.3)

• communication and teamwork: work collaboratively inaddressing problems and use appropriate SI notation andsignificant digits (all sections); and select and integrateinformation from various sources about technologicalapplications of cells (all sections)

Key Terms

14.1electrode

electrolyte

electric potential difference

electric current

fuel cell

14.2porous boundary

half-cell

voltaic cell

cathode

anode

inert electrode

standard cell

standard cell potential

standard reduction potential

reference half-cell

corrosion

cathodic protection

14.3electrolytic cell

electrolysis

14.4Faraday’s law

Faraday constant

Key Equations

• E °cell � E °r � E °r (14.2, 14.3)cathode anode

• ne� � (14.4)It�F

Go To

The following components are available on the Nelson Web site. Follow the links for Nelson Chemistry Alberta 20–30.

• an interactive Self Quiz for Chapter 14

• additional Diploma Exam-style Review questions

• Illustrated Glossary

• additional IB-related material

There is more information on the Web site wherever you seethe Go icon in this chapter.


MAKE a summary

1. Draw a diagram of a simple, general voltaic cell. Showand label all components, including names and signswhere appropriate. Show the directions of electron andion movement. State the process occurring at eachelectrode and how the cell potential is determined. Listseveral specific technological examples.

2. Repeat question 1 for a simple, general electrolytic cell.

3. Refer back to your answers to the Starting Points questions at the beginning of this chapter. How hasyour thinking changed?


Chapter 14 REVIEW

664 Chapter 14 NEL



Part 11. Which one of the following is the best example of a new

technology that led to scientific discoveries?A. electric cellB. Faraday constantC. reference half-cellD. hydrogen fuel cell

2. Which of the following statement applies more to sciencethan to technology?A. It involves the development of devices/process that

have a practical purpose.B. It seeks explanations about the natural world.C. It applies mainly to the human-designed processes.D. It applies trial-and-error strategies.

3. The main advantage of fuel cells is that theyA. are reliableB. are portableC. are inexpensive to buildD. run continuously as reactants are added

4. In a voltaic cell, the cathode is labelled as the ��

i��

electrode where the ��

i��i��

half reaction of the strongest

��i�i�i��

agent occurs.

The above statement is completed by the information inwhich row?

5. In numerical order, the statements that are true for voltaiccells are ___ ___ ___ ___.

6. In numerical order, the statements that are true forelectrolytic cells are ___ ___ ___ ___.

7. The cell potential of a standard lead–chlorine cell ispredicted to be

�_________ V

8. The main products formed in the electrolysis of an aqueousnickel(II) chloride solution are

at the anode at the cathode

A. O2(g), H�(aq) Ni(s)

B. Cl2(g) Ni(s)

C. Cl2(g) H2(g), OH�(aq)

D. Ni2�(aq) Cl�(aq)

9. If chemists had decided to use the standard zinc half-cellas the reference half-cell, the new standard reductionpotential for Hg2�(aq) would beA. �1.61 VB. �0.09 VC. �0.09 VD. �1.61 V

10. What is the standard reduction potential for Ce4�(aq) usingthe following net reaction and standard cell potential?

Ce4�(aq) � Fe2�(aq) → Ce3�(aq) � Fe3�(aq)E o

cell � �0.67 VA. �1.44 VB. �0.10 VC. �0.10 VD. �1.44 V

11. The industrial production of alkali metals involvesA. electrolysis of molten saltsB. electrolysis of aqueous saltsC. reduction with strong reducing agentsD. oxidation with strong oxidizing agents

12. Which of the following is not an effective method ofcorrosion prevention?A. galvanizingB. sacrificial anodeC. sacrificial cathodeD. impressed current

13. In an electroplating operation, the mass of copper metalproduced from a copper(II) ion solution by a 0.325 Acurrent flowing for 3.00 h is predicted to be

_________ g.

14. The following electrolytic cell is set up:

C(s) | Co(NO3)2(aq) | C(s)The minimum voltage that must be applied to operate thiscell is

�_________ V


Row i ii iii

A. negative oxidation oxidizing

B. negative oxidation reducing

C. positive reduction oxidizing

D. positive reduction reducing

Use the following information to answer questions 5 and 6.1. net reaction is spontaneous2. net reaction is nonspontaneous3. anions migrate toward the anode4. cell potential is positive5. cell potential is negative6. electrons flow from the cathode to the anode7. ions complete the internal circuit in the cell

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Part 215. Explain the function of the following components of voltaic

cells.(a) electrolyte(b) salt bridge(c) electrode(d) connecting wire

16. For each of the following cells, write the equations for thereactions occurring at the cathode and at the anode, andthe equation for the net cell reaction.(a) Ni(s) ⏐ Ni2�(aq) ⏐⏐ Cu2�(aq) ⏐ Cu(s)(b) Zn(s) ⏐ Zn2�(aq) ⏐⏐ Cr2O7

2–(aq), H�(aq) | C(s)(c) Pt(s) ⏐ H2(g) ⏐ OH–(aq) ⏐⏐ Ag�(aq) ⏐ Ag(s)

17. Draw and label a diagram of a voltaic cell constructed fromsome (not all) of the following materials:

strip of silver metal voltmeterstrip of lead metal connecting wiresaqueous silver nitrate glass U-tubeaqueous lead(II) nitrate cottonaqueous sodium nitrate various beakersdistilled water porous porcelain cup

18. Provide the following information for the cell constructed inquestion 17.(a) anode half-reaction(b) cathode half-reaction(c) net cell reaction(d) standard cell potential

19. Cathodic protection is widely used in industry forpreventing the corrosion of steel tanks and pipes. (a) Describe two forms of cathodic protection including

the chemical principles involved.(b) Assess the importance of corrosion protection using at

least two perspectives.

20. A student constructs an electrolytic cell to nickel plate acarbon rod. Draw and label a diagram of an electrolytic cellconstructed using the following materials:

strip of nickel metal power sourcecarbon rod connecting wiresaqueous nickel(II) sulfate large beakervoltmeter

21. Provide the following information for the cell constructed inquestion 20.(a) anode half-reaction(b) cathode half-reaction(c) net cell reaction(d) minimum voltage required

Chapter 14

22. As part of a class project, a student needs to construct acell that produces a voltage sufficient to operate a 1.25 Vclock.(a) Outline the design of a suitable cell including a fully

labelled cell diagram.(b) Justify your design using appropriate chemical terms,

equations, and calculations.(c) Suggest some possible discrepancies or limitations of

this cell for the chosen task.

23. In the Hall–Héroult process, an electric current is passedthrough the electrolyte at low voltage, but very highcurrent. Calculate the mass of aluminium metal that wouldbe produced from aluminium ions by a 150 kA currentflowing in a Hall–Héroult cell for 1.00 h (Figure 1).

24. A teacher constructs a demonstration voltaic cell, Zn(s) ⏐ Zn2+(aq) ⏐⏐ Cu2+(aq) ⏐ Cu(s), and allows it tooperate for 24.0 h. The mass of the cathode was measuredbefore and after the demonstration. Use the evidencecollected to calculate the average current produced by thiscell.

final mass of the cathode: 24.68 ginitial mass of the cathode: 21.12 g

25. Briefly explain the process of corrosion. Your response should include• several examples of metal corrosion of manufactured

materials• any environmental, health, or safety issues associated

with your examples• brief descriptions of some technological solutions to

the problem of corrosion• at least one example of corrosion that is desirable

Figure 1Molten aluminium from an electrolytic cell

DE

DE

DE


Unit 7 REVIEW

666 Unit 7 NEL



Part 11. A redox reaction involves a transfer of electrons

A. from the oxidizing agent to the reducing agentB. from the reducing agent to the oxidizing agentC. through a porous barrierD. between metals only

2. A general reaction type that is not a redox reaction isA. neutralizationB. disproportionationC. combustionD. formation

3. When solutions 1. sulfuric acid2. lithium hydroxide3. gold(III) fluoride4. chromium(II) nitrate

are ranked in order of strength of oxidizing agents, theorder, from strongest to weakest oxidizing agent, is __, __, __, and __.

4. During the process of photosynthesis,

6 CO2(g) � 6 H2O(g) → C6H12O6(aq) � 6 O2(g)

A. carbon in carbon dioxide is oxidizedB. hydrogen in water is reducedC. oxygen in carbon dioxide and/or water is oxidizedD. oxygen in glucose is oxidized

5. Which of the following reaction equations describes aredox reaction?A. C2H4(g) � 3 O2(g) → 2 CO2(g) � 2 H2O(g)B. H�(aq) � OH�(aq) → H2O(l)C. Ag�(aq) � Cl�(aq) → AgCl(s)D. HMnO4(aq) → H�(aq) � MnO4

�(aq)

6. The metal molybdenum, Mo(s), reacts to form MoO2(s).The half-reaction equation that explains the change inoxidation state of molybdenum can be written asA. Mo(s) � 2 e� → Mo2�(s)B. Mo(s) → Mo2�(s) � 2 e�

C. Mo4�(s) � 4 e� → Mo(s)D. Mo(s) → Mo4�(s) � 4 e�

7. A high school laboratory’s waste container is used todispose of aqueous solutions of sodium nitrate, potassiumsulfate, hydrochloric acid, and tin(II) chloride. The mostlikely net redox reaction predicted to occur inside thewaste container is represented by the equation:

A. 2 H�(aq) � 2 K�(aq) → H2(g) � K(s)

B. Sn2�(aq) � 2 NO3�(aq) � 4 H�(aq) →

N2O4(g) � 2 H2O(l) � Sn4�(aq)

C. SO42�(aq) � 4 H�(aq) � 2 Cl�(aq) →

H2SO3(aq) � H2O(l) � Cl2(g)

D. 2 Cl�(aq) � Sn2�(aq) → 2 Cl2(g) � Sn(s)

8. In the corrosion of steel objects in the natural environment,the most likely reducing agent isA. Fe2�(aq)B. Fe3�(aq)C. Fe(s)D. O2(g)

9. The metals1. Fe(s) 3. Zn(s)2. Sn(s) 4. Mg(s)

listed in order from least to most likely to corrode undersimilar atmospheric conditions are __, __, __, and __.

10. The reduction half-reaction that is generally involved in thecorrosion of iron isA. Fe2�(aq) � 2 e� → Fe(s)B. Fe(s) → Fe3�(aq) � 3 e�

C. 2 H2O(l) → O2(g) � 4 H�(aq) � 4 e�

D. O2(g) � 2 H2O(l) � 4 e� → 4 OH�(aq)

11. The net cell potential, under standard conditions, for theiron–oxygen cell in an aqueous environment is

�______ V.

12. In a titration experiment, 10.0 mL samples of 0.650 mol/Lchromium(II) ion solution reacted with an average volumeof 12.4 mL of acidic potassium dichromate solution. Theamount concentration of the potassium dichromatesolution is

________ mmol/L.

13. All voltaic and electrolytic cells requireA. one electrode and two electrolytesB. two electrodes and one or two electrolytesC. an external power supplyD. a voltmeter

14. Standard reduction potentials for half-cells are based onthe strengths ofA. oxidizing agents relative to hydrogen ionsB. oxidizing agents relative to hydrogen gasC. reducing agents relative to hydrogen ionsD. reducing agents relative to a standard acidic solution


Steel is the most widely used alloy in the world. However, anestimated 20% of the iron and steel produced annually isused to replace that lost by corrosion through exposure to airand water. Therefore, corrosion prevention is of considerableimportance. The empirical and theoretical chemistry ofcorrosion, and technological research and developmenttogether provide solutions to this important practical problem.


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Unit 7

15. In a voltaic cell, the reduction potentials of two standardhalf-cells are �0.35 V and �1.13 V. The predicted cellpotential of the cell constructed from these two half-cells isA. �0.35 VB. �0.78 VC. �1.13 VD. �1.48 V

16. If the electrodes of a standard copper–silver cell areconnected with a wire, thenA. silver is plated at the anodeB. a voltmeter would show a reading of 1.14 VC. the solution at the anode becomes darker blueD. electrons flow from the silver to the copper electrodes

17. The electrolysis of brine, NaCl(aq), is an importantindustrial process. The major products formed at eachelectrode are

Cathode AnodeA. H2(g), OH�(aq) O2(g), H�(aq)B. H2(g), OH�(aq) Cl2(g)C. Na(s) O2(g), H�(aq)D. Cl2(g) OH�(aq)

18. Electrons are transferred through theA. solution from the anode to the cathodeB. solution from the cathode to the anodeC. external wire from the anode to the cathodeD. external wire from the cathode to the anode

19. The product(s) at the anode will beA. O2(g), H�(aq)B. K(s)C. H2(g), OH�(aq)D. O2(g), H2O(l)

20. The mass of the gas produced at the anode is

________ g.

Part 2

21. Define oxidation and reduction in three different contexts:empirical (historical), theoretical (in terms of electrons),and theoretical (in terms of oxidation numbers).

22. Using a general reaction equation (A � B →C � D), label the agents and processes for any redoxreaction.

23. Define disproportionation and provide one simpleexample.

24. From the information in this unit, list two or three examplesof situations in which technology preceded scientificexplanations.

25. Name two common reactions that occur in living andnonliving systems. For each, identify the oxidizing agent,reducing agent, and the direction of electron transfer.

26. Distinguish, in as many ways as possible, between anodeand cathode. Does your answer apply equally to voltaic andelectrolytic cells? Explain briefly.

27. Briefly describe two technological solutions to theproblem of batteries “going dead.”

28. Explain why corrosion often occurs in places where twodifferent metals (such as copper and iron) are joinedtogether.

29. Electrochemical cells are very important technologicaldevices in our society. Compare the main differencesbetween voltaic and electrolytic cells in terms of theirpurpose and the chemical reactions that occur in them.

30. Predict whether a spontaneous redox reaction will occur inthe following situations: (a) A copper penny is dropped into hydrochloric acid.(b) A nickel is dropped into nitric acid.(c) A silver earring is dropped into sulfuric acid.

31. While working on the development of a newelectrochemical cell, a research chemist places selectedPeriod 4 transition metal strips into aqueous solutions oftheir ionic compounds. She observes that the followingcombinations of metal and cations react spontaneously:

V(s) � Mn2�(aq) → V2�(aq) � Mn(s)

V2�(aq) � Ti(s) → V(s) � Ti2�(aq)

Co2�(aq) � Mn(s) → Co(s) � Mn2�(aq)(a) Use this information to develop a table of oxidizing

and reducing agents for these metals and their ions.(b) Identify the strongest oxidizing and the strongest

reducing agent in your table.

32. Make a list of everything that must be balanced in a netionic equation representing a redox reaction.

33. Write and label balanced half-reaction equations for eachof the following redox reactions.

(a) 2 Fe3�(aq) � Ni(s) → 2 Fe2�(aq) � Ni2�(aq)

(b) Br2(aq) � 2 I�(aq) → 2 Br�(aq) � I2(s)

(c) Pd2�(aq) � Sn2�(aq) → Pd(s) � Sn4�(aq)(d) Label each reactant in (a), (b), and (c) as an oxidizing

or a reducing agent.

34. Use your knowledge of electrochemistry and somebrainstorming to describe at least three methods fordetermining or approximating the position of the berylliumhalf-reaction in a table of half-reactions.


An aqueous solution of potassium hydroxide undergoeselectrolysis using 5.9 A of current for a total time of 22 min.

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35. Potassium metal spontaneously reacts with water(Figure 1).

(a) Write the half-reaction and net ionic reactionequations for this reaction.

(b) Describe diagnostic tests (procedure, evidence,analysis) that could be done to test for the predictedproducts.

36. An acidic solution of potassium dichromate is added to asodium iodide solution. Predict the net ionic equation andshow all of your work. State two diagnostic tests that couldbe used to test your reaction prediction.

37. What is the oxidation number of(a) I in I2(s)? (d) H in NH3?(b) I in CaI2(s)? (e) H in AlH3?(c) I in HIO(aq)? (f) O in CH3OH?

38. Predict the balanced redox reaction equation byconstructing and labelling oxidation and reduction half-reaction equations. (a) MnO4

2�(aq) → Mn2� (aq) � MnO4�(aq) (acidic)

(b) ClO�(aq) → ClO2�(aq) � Cl2(g) (basic)

39. Use the oxidation number method to balance the reactionequations for the following redox reactions in acidicsolutions: (a) MnO4

�(aq) � H2C2O4(aq) →Mn2�(aq) � CO2(g) � H2O(l)

(b) KIO3(aq) � KI(aq) � HCl(aq) →KCl(aq) � I2(s) � H2O(l)

40. Chromium steel alloys are analyzed using a series of redoxreactions.

Step 1: The alloy is initially reacted with perchloric acid,which converts the chromium metal into dichromate ions,while the perchloric acid is reduced to chlorine gas.

Step 2: The dichromate ions are then reduced tochromium(III) ions by adding an excess of iron(II) solution.

Step 3: The unreacted iron(II) is then titrated with a solutionof cerium(IV) ions, which reduces them to cerium(III) ions.

Write a balanced redox equation for each step of this procedure.

41. Predict the mass of gold formed when a gold(III) nitratesolution reacts completely with 125 mL of 0.352 mol/Lsulfurous acid (aqueous hydrogen sulfite).

42. Complete the Materials (including precautions) andAnalysis of the following investigation report.

PurposeThe purpose of this exercise is to use redox stoichiometry tostandardize a potassium dichromate solution.

ProblemWhat is the amount concentration of a potassium dichro-mate solution?

EvidenceVolume of iron(II) ammonium sulfate solution � 10.0 mL

Concentration of acidic Fe(NH4)2(SO4)2(aq) � 0.0625 mol/L

43. Electrolysis is used in the industrial production of severalimportant elements and compounds. (a) Define electrolysis.(b) In an electrolytic cell, what type of half-reaction

occurs at the anode? at the cathode?(c) Compare the electrolysis of molten compounds with

the electrolysis of aqueous solutions. State somesimilarities and differences.

(d) Briefly describe three important industrialapplications of electrolysis.

44. In which of the following mixtures must an external voltagebe applied to inert electrodes to observe evidence of aredox reaction? (a) a solution of cadmium nitrate(b) a solution of iron(III) iodide(c) solutions of iron(III) chloride and tin(II) sulfate in

separate half-cells connected by a salt bridge(d) solutions of potassium iodide and zinc nitrate in

separate half-cells

45. Determine the minimum potential difference that must beapplied to the following electrolytic cells to cause achemical reaction. (You do not need to write the half-cellreaction equations.) (a) iron(II) sulfate electrolyte with inert electrodes(b) hydrochloric acid electrolyte with silver electrodes(c) tin(II) chloride electrolyte with tin electrodes

46. Write the equations for reactions at the cathode andanode, and the net cell reaction. Determine the minimumpotential difference that must be applied to each of thefollowing electrolytic cells to cause a reaction. (a) C(s) ⏐ NaBr(aq) ⏐ C(s) (b) Pt(s) ⏐ KOH(aq) ⏐ Pt(s) (c) C(s) ⏐ CoCl2(aq) ⏐ C(s)

Figure 1Potassium reacts vigorouslywith water.

Trial 1 2 3



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Unit 7

47. Volta’s invention of the electric battery in 1800 led to aflurry of scientific research using this new technology. Afew weeks after he heard about Volta’s battery, WilliamNicholson, an English chemist, built his own battery andpassed a current through slightly acidified water. With thecurrent flowing, bubbles of colourless gases formed ateach electrode. This was the first demonstration that anelectric current could bring about a chemical reaction. (a) Write equations for the cathode, anode, and net

reactions that occurred in Nicholson’s demonstration.(b) Determine the minimum potential difference needed

for the reaction.

48. Potassium hydroxide is obtained commercially by theelectrolysis of aqueous potassium chloride. (a) Sketch a diagram of a cell that could be used to

electrolyze an aqueous solution of potassium chloride.Label electrodes, electrolyte, power supply, and thedirections of the electron and ion flow.

(b) Predict the cathode, anode, and net reactions, andcalculate the minimum potential difference for theelectrolysis of aqueous potassium chloride.

49. One technological process for refining zinc metal involvesthe electrolysis of a zinc sulfate solution. (a) Write equations for the cathode, anode, and net

reactions, and calculate the minimum potentialdifference for the electrolysis of a zinc sulfate solution.

(b) Calculate the time required to produce 1.00 kg of zincusing a 5.00 kA current.

50. Determine the current required to produce 1.00 kg ofaluminium per hour in a single Hall–Héroult cell for theproduction of aluminium.

51. The electrolysis of copper(II) sulfate using carbonelectrodes is demonstrated to a chemistry class. In thedemonstration, a 1.50 A current passes through 75.0 mL of0.125 mol/L copper(II) sulfate solution. How long, inminutes, would it take to plate all of the copper from thesolution?

52. Given that the typical current used in the chlor–alkali plantis 55 kA, predict the rate at which chlorine gas isproduced (in moles per hour).

53. Battery technology is a very active area of research. Oneproposal that shows some promise is a vanadium redoxflow cell, also known as the All Vanadium Redox Battery.Describe the general construction of this battery,including electrodes, electrolytes, porous boundary, andexternal tanks. What redox reactions occur at theelectrodes within this cell? List some unique aspects of thistechnology and some advantages and proposed uses.

54. Rechargeable nickel–metal hydride (NiMH) batteries havetwice the energy density of Ni–Cd batteries and a similaroperating voltage. The cells in NiMH batteries useNiO(OH)(s) as one electrode, a hydrogen-absorbing alloy asthe other, and an alkaline electrolyte. In the following

reduction half-equations, M indicates a hydrogen-absorbing alloy and Hab indicates absorbed hydrogen.

NiO(OH)(s) � H2O(l) � e� →Ni(OH)2(s) � OH�(aq) E °r � �0.49 V

M(s) � H2O(l) � e� → MHab � OH�(aq) E °r � �0.71V(a) Write balanced equations for the anode, cathode, and

net reactions occurring during the cell’s operation.(b) Determine the cell potential.(c) Research and list some of the technological,

economic, and environmental considerations involvedin evaluating the NiMH battery.

55. Electroplating finishes are often layered. For example,chromium plating does not work well on a zinc base, so alayer of copper is applied to the zinc, then a layer of nickelis added, and then the top chromium layer is plated on. (a) Propose a general design of an experiment to place a

final chromium layer onto a galvanized metal. Includea labelled diagram and general plan.

(b) In any electroplating process a particular thickness ofmetal is desired. Outline the experimental variablesand the type of calculations necessary to plan aparticular thickness of metal plating.

56. Review the focusing questions on page 552. Using theknowledge you have gained from this unit, briefly outline aresponse to each of these questions.

Extension

57. The Alberta government is considering making carbonmonoxide detectors mandatory for new residential homesas is done in Ontario. What are the sources and the healtheffects of carbon monoxide? Briefly describe how amodern CO detector works including the reactionequations and catalyst involved. Should CO detectors berequired in all homes? Justify your opinion.

58. Moli Energy of Maple Ridge, British Columbia, was the firstcompany in the world to develop a commercial,rechargeable lithium ion cell, called a Molicel (Figure 2).Research the characteristics and advantages of Molicelscompared with other secondary cells.



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safety header

positiveelectrode

negativeelectrode

separator

Figure 2The Molicel is ahigh-energy,rechargeablelithium–ion cellin a unique,jellyroll design.

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