ELECTROCHEMISTRY

Teacher's Guide

ElectrochemistryTeacher's GuideThis teacher's guide is designed for use withthe Electrochemistry series of programsproduced by TVOntario, the television serviceof The Ontario Educational CommunicationsAuthority. The series is available on videotapeto educational institutions and nonprofitorganizations. For ordering information, seepage 24.

The SeriesProducer: David ChamberlainWriter: William KonradNarrator: James MoriartyConsultant: John EixDesigned and produced by NortheyProductions for TVOntario, 1986.

The GuideWriter: John EixConsultant: John StrattonEditor: Deborah BurrettDesigner: Michael E. Bowness

ContentsProgram 1: The Building Blocks of

Electrochemistry ........ 1Program 2: Electrochemical Cells .... 6Program 3: Designing Electrochemical

Cells ................... 9Program 4: Commercial Electro-

chemical Cells .......... 12Program 5: Corrosion ............... 17Program 6: Electroplating ........... 20

©Copyright 1987 by the Ontario EducationalCommunications Authority.

All rights reserved.

Printed in Canada 2316186

Note to TeachersThis series of six 10-minute programs isdesigned to introduce senior chemistrystudents to the basic concepts of electro-chemistry: "redox" reactions, activity series,reduction potentials, electrochemical cells,corrosion, and electrolysis. The termsoxidation, reduction, oxidizing agent, reducingagent, anode, cathode, polarization, anion,cation, and half reaction are introduced earlyand then used to explain a number of electro-chemical processes. These processes includethe production of energy in an electro-chemical cell, corrosion, galvanization, theuse of sacrificial anodes, and electroplating.Throughout Electrochemistry the example of arobot subject to "heart failure" (weak orcorroding batteries) and "skin disease"(corrosion) is used to demonstrate the effectsof the processes being discussed. Theproblem of building a better robot links eachprogram to the next.

This guide contains a description of eachprogram. It also provides learning objectives,activities for before and after viewing eachvideotape, and suggestions for furtherreading.

The Building Blocks of ElectrochemistryObjectivesAfter viewing the program and completing this chapter,students should be able to:

1. Define and provide examples of these terms: oxidation,reduction, redox reactions, oxidizing agent, reducingagent, electrode, electrolyte, anode, cathode, anion, andcation.

2. Identify the roles that redox reactions play in everyday life.

Program DescriptionAt the beginning of this program, the students meet the"main character" of the series, a robot - a battery-operatedi ron man who at first appears to be a tireless helper. Soonenough, however, his weaknesses begin to show: hisbatteries run down and his metal parts begin to rust. It isexplained that both these developments are the result ofsimilar chemical reactions. The program then goes on toexamine redox reactions, explaining oxidation and reduction,and describing the role of the oxidizing and reducing agents.Reactions between sodium and oxygen and between sodiumand chlorine are used as examples.

The production of electrical energy by the redox reaction inan electrochemical cell is studied next. The students arei ntroduced to the terms electrode, electrolyte, anode,cathode, anion, and cation. The program ends by pointing outan apparent anomaly in the chemical activity in the cell:negatively charged anions move toward the negative elec-trode (anode), while positively charged cations move towardthe positive electrode (cathode). This problem is taken up inProgram 2.

Before ViewingSince the program assumes a familiarity with electrons andions, it is advisable to review these ideas with students be-fore viewing it. One way is to review nomenclature and oxida-tion numbers or states. Activity 1, a worksheet, provides a setof rules for assigning oxidation numbers. You could assignthis as homework the day before viewing.

The program introduces oxidation as a loss of electrons;

however, it can also be defined as an increase in the oxida-tion number. Activity 2 will help the students see that thedefinitions for oxidation and reduction essentially say thesame thing. The definition involving oxidation numbers is, ofcourse, more applicable with non-ionic compounds. Question1 of Activity 2 can be used effectively before or after viewingthe program.

After ViewingUse question 2 of Activity 2 to practice the use of theterminology introduced in the program. At this point, it isnecessary to deal with balancing redox equations. Activity 3is a worksheet that you can use to introduce students to thisbalancing task, using the number of electrons transferred.

ActivitiesActivity 1: Oxidation Numbers

An oxidation number i s defined as thecharge an atom appears to have whenelectrons involved in a chemical bond arecounted and assigned to the atomsinvolved in the bond, according to somerather arbitrary rules. The word "appears"is crucial because ions are involved onlyin ionic compounds. After you have readthese rules you may ask yourself whatthe point is. If this is your conclusion,reread the first sentence in the followingparagraph carefully: oxidation numbersare very useful to a chemist because theymake the job easier.

Oxidation numbers are useful for writ-ing formulas of chemical compounds, fordescribing what happens on a micro-scopic scale during a chemical reaction,for predicting the charge on an ion, andfor balancing chemical equations. Youmay have met oxidation numbers beforeor be aware of very similar concepts such

1

as oxidation state, electrovalence,valency, and combining power.

Rules for AssigningOxidation NumbersThere are two basic rules for countingand assigning the electrons involved in achemical bond:1. Electrons involved in a chemical bond

between two unlike atoms are countedwith the more electronegative atom.

2. Electrons involved in a chemical bondbetween two like atoms are dividedequally between the two atoms.

Examples

before bonding with the hydrogen atoms.This gives it an apparent charge of 2-and an oxidation number of 2-.

I n principle it is possible to deduce theoxidation number of an atom in any com-pound this way. The method, however, islaborious. It is more convenient to use thefollowing operational rules, which are der-ived from the two above.

Operational Rules for AssigningOxidation Numbers1. In the free elements, each atom has an

oxidation number of 0.For example. hvdrooen in

2. In simple ions (ions containing oneatom), the oxidation number is equal tothe charge on the ion.

I n these cases, the apparent charge onthe atom is the actual charge on the ion.Thus, the 3+ aluminum ion and the 2-oxygen atom in the ionic compoundaluminum oxide have oxidation numbersof 3+ and 2-, respectively.It is useful to remember that the alkalimetals form only 1+ ions and the alka-line earth metals form only 2+ ions.Therefore, the elements lithium, sodium,potassium, rubidium, cesium, andfrancium will always have an oxidationnumber of 1+; and beryllium, magnesium,calcium, strontium, barium, and radiumwill always have an oxidation numberof 2+. Similarly, the halogens as halidesalways have an oxidation number of 1-.

3. In most compounds containingoxygen, the oxidation number of eachoxygen atom is 2-.

The exception to this rule is the oxida-tion number of oxygen in peroxides,which is 1-.

4. In most compounds containing hydro-

gen, the oxidation number of eachhydrogen atom is 1+.The exception is hydrogen in a hydride.I ts oxidation number is 1-.

5. All oxidation numbers must be consis-tent with the conservation of charge.

Charge must be conserved in the sensethat all the apparent charges (i.e., the oxi-dation numbers) in a molecule or ionmust add up to the net charge on themolecule or ion. For example, in a neutralmolecule the oxidation numbers of all theatoms must add up to 0. For complexions (ions involving more than one atom),the oxidation numbers of all the atomsmust add up to the charge on thecomplex ion.Rule 5, the neutrality rule, enables us toassign oxidation numbers to atoms thatare not covered by the first four opera-tional rules. For example; it allows us todetermine the oxidation number of sulfuri n H2SO4 in the following way:• From rule 4 we know that each

hydrogen atom has an oxidationnumber of 1+, but because there aretwo hydrogen atoms, the total is 2+.

• From rule 3 we know that each oxygenatom has an oxidation number of 2-,but because there are four oxygenatoms, the total is 8-.

• From rule 5 we know that the oxidationnumbers of all atoms must add up to0. Thus,

NOTE: ON = Oxidation number

As the oxidation number of sulfur inH 2SO4 is 6+, the compound is sometimescalled sulfuric(VI) acid, with the Romannumeral Vi indicating sulfur's oxidationnumber. The oxidation number of nitrogen

2

The electron dot (or Lewis structure) forH 2 is H:H. Since both atoms are identical,each hydrogen atom gets one of the twoelectrons involved in the chemical bond.Since hydrogen's atomic number is 1, thenuclear charge is 1+. Thus the 1 electronfrom the bond will balance the nuclearcharge. The apparent charge on thehydrogen atom will be 0. Thus, the oxida-

will be 0.tion number of H iiWhat is the oxidation number of the

H and C) ntomc inThe electron dot or Lewis structure for

is H:O:HSince the hydrogen and oxygen atoms

are not alike and the oxygen is moreelectronegative, the electrons involved inthe chemical bonds are counted with theoxygen atom. Thus, the hydrogen atomappears to have no electrons and acharge of 1+. The hydrogen atom isassigned an oxidation number of 1+. Theoxygen atom - because it is more elec-tronegative - has all four of the elec-trons involved in the chemical bondsassigned to it, and it has apparentlygained two electrons that it didn't have

sodium inid phosphorus in P4 al

have oxidation numbers of 0.Na, sulfur in

Because the assignment of oxidationnumbers is quite arbitrary, do not besurprised if you get oxidation numbersthat do not agree with your ideasabout charge. Keep in mind that mostcompounds are not ionic and that theoxidation number is the apparentcharge, based on the assignment ofelectrons to an atom according to itsrelative electronegativity.

Practice

1. Use the two basic rules for assigningoxidation numbers to explain whyoxygen's oxidation number inperoxides is 1-.

2. Use the two basic rules for assigningoxidation numbers to explain whyhydrogen's oxidation number inhydrides is 1-.

i . Assign oxidation numbers to each ofthe atoms in the followina: NaCl CaO

Activity 2: Oxidation Numbersand Chemical ReactionsOxidation numbers allow chemists todistinguish oxidizing agents fromreducing agents. This distinction isimportant when storing chemicals. Forexample, it's extremely dangerous tostore oxidizing agents such as potassiummanganate(VII) - commonly known aspotassium permanganate - with fuelssuch as glycerine; accidental mixing ofthe two could result in an explosion.When the oxidation number of an elementi ncreases, the substance to which the

element originally belonged is said to beoxidized, and the change is calledoxidation. The substance in the chemicalreaction that causes the oxidation iscalled the oxidizing agent. The oxidizingagent is always reduced.When the oxidation number of an elementdecreases, the substance to which theelement originally belonged is said to bereduced, and the change is calledreduction. The substance in the chemicalreaction that causes the reduction iscalled the reducing agent. The reducingagent is always oxidized.

Consider the following chemical reaction:oxidation

+-------------------------------+2+ ON of Cu 0

decreases reduction

In this reaction, zinc i s oxidized to zinc sulfateand the change is called oxidation. Thecopper(II) sulfate causes this oxidation, and isthus the oxidizing agent. Looking at the samereaction from the other perspective, thecopper(II) sulfate is reduced to copper, and thechange is called reduction. The zinc causesthe reduction of the copper and thus is thereducing agent. Note that the oxidizing agenti s reduced and that the reducing agent isoxidized.

Practice

1. For chemical reactions i to viii below,a) assign oxidation numbers for each

element;b) record the process as shown in the

example above; identify theoxidation and reduction parts of thereaction; and

c) identify the reducing and oxidizingagents.

2. After viewing "The Building Blocks ofElectrochemistry" video program,describe the electron transfer thattakes place in unbalanced chemicalreactions i to viii below. (To help yourecall the important relationships, usethe mnemonic "LEO the lion saysGER" LEO stands for "Loss ofElectrons-Oxidation:' GER stands for"Gain of Electrons-Reduction")

Activity 3: BalancingRedox EquationsIt is essential to have a balancedchemical equation in order to understandwhat happens at the atomic level duringthe reaction. Reacting mole ratios arealso essential if the reaction is used in anindustrial process or in analytical work.Therefore, balancing chemical equationsis a necessary skill.

Since redox reactions are complex, thetrial-and-error method used to balanceordinary chemical equations can befrustrating, and often unsuccessful. Amore effective method exists for redoxreactions. Unlike the trial-and-errormethod, whose one starting principle isthe conservation of atoms, the redoxreaction method starts with the balancingof the number of electrons gained withthe number of electrons lost. This meansthat two principles are used-conserva-tion of charge and conservation of atoms.

There are two methods of applying

3

these principles. An example of the firstmethod-using oxidation numbers-isshown below.

Using Oxidation Numbers To BalanceRedox Equations

Example 1

You can balance the following equationby inspection, and you should try to do sobefore working through the six steps out-lined below for the oxidation-numbermethod.

STEP 1Analyze the equation and assign oxida-tion numbers for those elements thatchange.

STEP 2Determine the number of electrons to betransferred per atom.

gain 3e/atomreduction

oxidation reducing agent: Sl oss 4e/atom

STEP 3Determine the number of electrons to betransferred per formula-unit.

gain 6e/formula-unitgain 3e/atomreduction

oxidation reducing agent: Sloss 4e/atomloss 4e/formula-unit

STEP 4Balance electrons lost and gained bywriting appropriate coefficients for theoxidizing agent and the reducing agent.

STEP 6Add coefficients to all substances exceptthe oxidizing agent and the reducingagent to balance ionic charge and con-serve atoms.

Thus, the balanced equation is:

Example 2

Consider the situation in which a solutionof sodium sulfite is mixed with an acidicsolution of potassium manganate(VII).Through analysis of the resulting color-less solution, sulfate ions andmanganese 2+ ions are identified. Whatis the balanced chemical equation?

STEP 1Analyze the equation and assign oxida-tion numbers for those elements thatchange.

Because the reaction takes place in solu-tion, it is possible to write the followingionic equation:

total gain = 12e2*gain 6e/formula-unitgain 3e/atomreduction

4

STEP 2Determine the number of electrons to betransferred per atom.

gain 5e/atomreduction

STEP 3Determine the number of electrons to betransferred per formula-unit.

gain 5e/per formula-unitgain 5e/atomreduction oxidizing agent: KMn04

oxidation reducing agent: Na2 SO 3l oss 2e/atoml oss 2e/formula-unit

STEP 4Balance electrons lost and gained bywriting appropriate coefficients for theoxidizing agent and the reducing agent.

total gain = 10e2*gain 5e/formula-unitgain 5e/atomreduction

STEP 6Add coefficients to all substances exceptthe oxidizing agent and the reducingagent to balance ionic charge and con-serve atoms.

At this stage, if you take into accountconservation of atoms, the equationwould read:

The total charge on reactant side is2*1- + 5*2- = 12-, and the total onthe product side 2*2+ + 5*2- = 6-. Tobalance the ionic charge. add 6 hvdrooen

If you check the number of oxygen atomsrepresented on each side you will notethat the equation is balanced. Example 2represents the kind of reasoning youmust do to determine the balanced equa-tion for reactions that take place inaqueous solutions. Example 1 illustratesthe reasoning to use when all reactantsand products for a reaction are known.

Practice1. Balance the following equations and

identify for each the oxidizing andreducing agents.

Further ReadingSienko, Michell J., and Robert A. Plane.

Chemistry, fifth edition. New York:McGraw-Hill, 1976Silk, T Electrochemistry. London:

Heinemann Educational Books, 1968.Toon, E.R., and G.P Ellis. Foundations of

Chemistry. New York: Holt, Rinehart andWinston, 1973.

5

balance the Hs, add 3 water molecules tothe product side. The resultant equationi s:

. to the reactant side. Tci ons

ObjectivesStudents should be able to:1. Write redox equations for electrochemical cells.2. Describe on a microscopic and macroscopic scale what is

happening in an electrochemical cell.3. Explain polarization and how it is avoided in a cell.4. Interpret an activity series in terms of ease of reduction of

a given substance.

Program DescriptionProgram 2 begins with a review of oxidation and reduction,using the example of a dry cell to focus on the electronprocesses that occur at the anode and the cathode. This isfollowed by a detailed look at the reaction between zinc andcopper(II) sulfate. In the competition between the zinc atomand the copper ion for the zinc's electrons, copper is present-ed as the "bully," "stealing" the electrons away. In otherwords, the copper(II) is shown to be more easily reduced thanthe zinc. By introducing a second example in which mer-cury(II) ions are able to "steal" electrons from copper atoms,the program demonstrates how a "pecking order" of ions canbe established. This "pecking order," in which the ion at thetop is most easily reduced, and so on down the scale, is whatis known as the activity series.

After pointing out that the activity series is useful whenselecting possible materials for a battery, the programreturns to the reactions between zinc atoms and copper (II)ions. According to the activity series, it should be possible touse these two substances to create a battery. However, whenthe zinc and copper half reactions are set up in separatebeakers with a wire connecting the two electrodes, no currentis produced. The concept of polarization is introduced toexplain why a salt bridge is necessary to complete the circuitin an electrochemical cell.

The program closes by pointing out that in order to build abetter "heart" for a robot, it would be useful to be able toquantify the amount of electricity produced by various pairsin the activity series.

Before ViewingDo Activity 1. It provides a link between Programs 1 and 2,giving the students practice applying the terminology ofProgram 1 as well as providing data for creating an activityseries. Have the students answer discussion question 1before viewing the video.

After ViewingHave students complete Activity 1.

ActivitiesActivity 1: Redox ReactionsIn this activity you will make observationsthat will enable you to identify theproducts of reactions. Armed with yourobservations, the knowledge that atomsare conserved in a chemical reaction, andthe Handbook of Chemistry and Physics,you should be able to identify theproducts. Then you will identify theoxidizing agent and the reducing agentand balance the equation.

The activity has been organized intostations which you can work your waythrough. Order is not important. Some ofthe stations will be demonstrated by yourteacher because of the hazards involved.For all stations, record your observations,write a balanced chemical equation forthe reaction, and identify the substanceoxidized, the substance reduced, theoxidizing agent, and the reducing agent.

6

Electrochemical Cells

STATION 1:Acids and MetalsAt this station you should note the rate ofthe reactions. From these observationsyou will be able to set up a "peckingorder" of metallic elements. It is impor-tant that you record observations in sucha way that you can compare variousreactions.

Materials

1. small bottles containing small chunks(max 1 cm3) of as many of the follow-ing metals as possible: Mg, AI, Zn, Fe,Pb, Cu, AgNOTE: Since the metals you will beworking with react with air, it is impor-tant that you make sure the surfacesof the metals are clean.

3. 7 test tubes and a test tube rack orbeakers to hold test tubes

4. supply of dry test tubes

5. Bunsen burner, striker, burner pad

6. emery cloth or sandpaper for cleaningmetals if required

Procedure

CAUTION Goggles must be worn duringthis activity.CAUTION Reactions require adequateventilation, especially when the nitricacid is used.CAUTION Acids are corrosive. Wear pro-tective gloves. If acid comes in contactwith flesh or clothing, rinse with anabundance of running water.

1. Add hydrochloric acid to the level of2-3 cm in each of the seven test tubes,one for each metal provided.

2. Add a different metal simultaneouslyto each test tube.

3. Collect the gas produced by the testtubes with the most vigorous reac-tions. Use a dry test tube and collectthe gas by the downward displacementof air.CAUTION Keep the flame as far awayas possible from the test tube produc-i ng the gas.

4. Bring the mouth of the test tube inwhich you have collected the gas tothe side of the burner flame. If you donot get a characteristic pop, wave thetest tube through the air and try again.

5. When you have recorded your obser-vations, rinse test tubes and metalswith copious quantities of tap water.

6. Repeat procedures 1-5 using sulfuricacid.

7. Repeat procedures 1-5 using nitricacid.

8. When you have completed all threesets of procedures, be sure that allused metals have been thoroughlyrinsed. Put each metal in the appro-priate USED metal container.

STATION 2:Metal ReplacementsAt this station we are looking for a "peck-i ng order" among chemicals; that is, weare looking to see what each chemicalcan do and to which other chemicals. Youwill want to make observations that willenable you to compare reactions. Use atable such as the one in Observations,below.

Materials

1. small bottles containing small chunks(max 1 cm3) of as many of the follow-ing metals as possible: Mg, Zn, Pb, Cu,AgNOTE: Since the metals you will beworking with react with air, it is impor-tant that you make sure the surfacesare clean.

3. 25 test tubes or small petri dishes

4. emery cloth or sandpaper to cleanmetals if required

Procedure

1. Place a piece of zinc in a test tube orsmall petri dish. Cover with thecopper(II) nitrate solution. Let it sit for10-20 min.

2. Repeat for all possible combinationsshown in the following table. Recordyour observations in the table.

Observations

Metal-Metal Ion Redox Reactions

STATION 3Halogen Replacements(TEACHER DEMO IF NO FUMEHOOD)

At this station we are again looking for a"pecking order" among chemicals. It isimportant that you make observationsthat enable you to compare reactions.

Materials1. chlorine water (dilute bleach), bromine

water, iodine dissolved in a 1:1 wateralcohol mixture

3. TTE (t rich IorotrifIuoroethane) [orhexane] in a dropper bottle

4. prepared solutions of TTE [or hexane]with iodine, chlorine, and bromine

5. test tubes, rubber stoppers to fit andtest tube rack

Procedure

POUR DOWN THE SINK.CAUTION Chlorine and bromine are verycorrosive. If they come in contact withflesh or clothing rinse with a dilute solu-tion of sodium thiosulfate, commonlycalled "hypo," and copious quantities ofrunning water. Wear protective gloves.

5. Repeat procedures 1-4 twice, substitut-i ng first the bromine water and thenthe iodine solution for the chlorinewater. Record your observations in thetable below.

6. Do the following in subdued light. Addsodium chloride to the level of 1 cm ina test tube. Add approximately 0.5 cmof silver nitrate solution to the testtube. Swirl the contents and add a fewdrops of chlorine water, swirl again.Repeat twice, replacing the chlorinewater with the bromine water and thenthe iodine solution. Repeat for thesodium bromide and sodium iodidesolutions.

Observations

2. For stations 1 and 2:a) List the metals from most reactive

to least.b) State the evidence you used to de-

velop your list.c) Write half reactions describing

what happens to the metals inthese reactions.

d) Order the ions of the metals in anactivity series with the most easilyreduced ion at the top.

e) Describe the series you havecreated in terms of the ability of theion to act as an oxidizing agent.

3. For station 3:a) Answer parts a-e of question 2.

Note that in station 3 the element isreduced, not the ion.

b) Is it possible to link the two activityseries you have created based onthe evidence gathered in thisactivity? If so, do it; if not, explainwhat additional evidence you wouldneed.

CAUTION The chemicals used in this sta-tion are toxic and corrosive. Wear pro-tective gloves and ensure adequateventilation.

1. Pour sodium chloride solution into atest tube to the level of 1 cm. Repeatfor the sodium bromide and sodiumiodide solutions, in separate testtubes.

2. Add 1 cm of chlorine water to each testtube.

3. Use the dropper bottle to add approxi-mately 0.5 cm of TTE to each test tube.

4. Stopper the tubes and shake. Comparethe color of the TTE layer with the pre-pared TTE solutions. Dispose of thecontents of the test tube into thebottle labelled TTE Residues. DO NOT

Halogen Replacements

Discussion1. For stations 1-3:

a) Write balanced chemical reactionsthat account for observations made.

b) Identify: element oxidized, elementreduced, reducing agent, oxidizingagent.

Further ReadingParry, Robert W., et al. Chemistry:

Experimental Foundations. EnglewoodCliffs, N.J.: Prentice-Hall, 1975.

Sienko, Michell J., and Robert A. Plane.Chemistry, fifth edition. New York:McGraw-Hill, 1976.

Silk, T Electrochemistry. London:Heinemann Educational Books, 1968.

Toon, E.R., and G.R Ellis. Foundationsof Chemistry. New York: Holt, Rinehartand Winston, 1973.

8

2. 0.1 kmollm3 solutions of NaCl, NaBr,i n dropper bottles

Designing Electrochemical CellsObjectivesStudents should be able to:1. Define and illustrate the term "standard reduction

potential"2. Describe the standard hydrogen electrode.3. Describe how a table of standard reduction potentials is

obtained experimentally.4. Use a table of standard reduction potentials to calculate

the potential of an electrochemical cell.5. Given experimentally measured cell potentials, create a

table of reduction potentials.

Program DescriptionAfter a brief review, this program takes up the problem of howto design a "super cell:' It explains that in order to select thebest materials for a battery, it is necessary to be able topredict what the energy output of any pair of half reactionswould be. This can be done by determining the comparativereduction potential of each material. The program explains

reduction potential of the second substance to obtain its oxi-dation potential; and c) add the potential for the reductionreaction to the potential for the oxidation reaction. The totalis the cell's potential.

Before ViewingWhen discussing the activity series developed by thestudents in Activity 1, Program 2, establish that if valuescould be attached to the positions it would be possible toselect the best possible materials for a battery. Ask studentsto suggest methods for doing this. After briefly exploring themerits of their suggestions, show "Designing Electro-chemical Cells."

After ViewingHave students complete Activity 1. In this activity they havean opportunity to measure cell potentials and create a tableof half-reaction reduction potentials, using a method similarto the one shown in the video. Note that if you wish to obtainvalues that are close to the accepted values, voltmeters witha very high impedance must be used. Otherwise, the meterrequ-ires too much current to register a reading and themeasured values will be too low.

Because the video leaves the impression that reductionpotentials are all you need to know to design batteries, it isi mportant that students become aware of the other factorsthat must be considered in designing a viable battery.Activity 2 will help make them aware of these other factors. Italso provides a link with Program 4.

ActivitiesActivity 1:

Electrochemical Cells

I n this activity you measure the cellpotential of three electrochemical cells_

9

used to demonstrate that one must a) determine which sub-stance has the greater reduction potential; b) reverse the

used instead of the copper, the current flows in the oppositedirection. Zinc ions are therefore assigned a negative value:-0.76 V The program points out that half-cell reactions are,by convention, written in tables as reductions. When anyreduction equation is reversed to become an oxidation equa-tion, the sign of the potential is also reversed (from positiveto negative or vice-versa). The program then explains thatwith this knowledge it is possible to predict the potential(voltage) of any electrochemical cell. The examples of the

that a hydrogen half-cell _standard reduction potential (E°) ,of 000 V When this cell isconnected with a copper half-cell

i s assigned n

a

] i s

meter reading of 0.34 V is obtained. The electrons are movingfrom the hydrogen electrode to the copper electrode becausethe copper ions are more easily reduced. The copper ions aretherefore assigned a positive standard reduction value:+0.34 V Because zinc ions are more difficult to reduce thanhydrogen ions, when a zinc half cell

half-cell. You can use it as the

1 0

standard half-cell and create a table ofreduction potentials.

Materials

Procedure

CAUTION Lead solutions can beabsorbed through the skin; lead com-pounds are toxic. Wear protective gloves.

1.Fill three U-tubes with the ammoniumnitrate solution. Stuff cotton batteni nto the ends of the U-tubes so it ispossible to turn the tubes over withoutthe contents coming out. Use a differ-ent U-tube with each pair of solutionstested to minimize contamination ofthe solutions.

2.Pour approximately 200 mL of thecopper solution into one of the 250 mLbeakers. Pour 200 mL of the zinc solu-tion into another beaker.

3.Place a copper strip in the coppersolution and a zinc strip in the zincsolution.

4. Connect the strips to the voltmeterusing the connecting wires. Connectthe leads so it is possible to read thevoltage. Adjust the meter, if possible,to get the maximum deflection. (Yourteacher will explain how to do this.)Record the voltage in the table below.

Voltmeter

5.

6.

7.

beaker containing the zinc solutionwith one containing a lead strip inapproximately 200 mL of the leadnitrate solution. Use a differentU-tube for this cell. Connect theleads, read and record the voltage.Determine the electrode that acts asanode, as cathode.

8. Remove the U-tube and replace thebeaker containing the lead solutionwith one containing a silver strip inapproximately 200 mL of the silvernitrate solution. Use a differentU-tube for thi s cell. Connect theleads, read and record the voltage.Determine the electrode that acts asanode, as cathode.

9. Move either electrode up and down inthe solution. Record what happens tothe voltage.

10. (TEACHER DEMO) Reassemble thezinc-copper cell. Add the sodium sul-fide solution slowly, stirring steadily,to the beaker containing the coppersolution. Record what happens in thebeaker and what happens to thevoltage.

Observations

invert one of the U-tubes into the twobeakers as shown in the diagram, andrepeat step 4.

Look at the posts on the voltmeter toidentify which is positive and whichis negative. Use this information todeduce the direction of current flowand which strip (electrode) is actingas an anode, and which, the cathode.

Remove the U-tube and replace the

Discussion1. Explain why there was no voltage

reading in step 4.

3. U-tube and cotton batten

4. 1x11 cm strips of copper, zinc, lead,and silver

5. 2 connecting wires with alligator clips

6. 1 0-5 V high impedance voltmeter; eg.,a VTVM

7. 4 250 mL beakers8. stirring rod

2.

Cell Anode Cathode Voltage

Experi-mental

Accepted

zinc-copper

l ead-copper

silver-copper

2. What is the purpose of the U-tube?

3. Describe how the U-tube performs thefunction identified in question 2.

4. Write half reactions for the chemicalchanges taking place at each elec-trode. Write a chemical equationdescribing the total reaction takingplace in each cell.

5. Describe the direction of flow of elec-trons in each cell.

6. The following shorthand notationdescribes the zinc-copper cell:

Describe the conventions used in thisshorthand and write similar descrip-tions for the other cells studied.

7. Use the measured voltages to estab-lish a table of reduction potentials.

8. Use the table of reduction potentialsin your text or the Handbook ofChemistry and Physics to determinethe accepted potential for each cellstudied. Record the value in the tableabove.

9. Suggest plausible reasons for anydifferences.

Activity 2:Commercial ElectrochemicalCells or BatteriesYour task is to produce an attractivedisplay of information about one type ofbattery. Your efforts will be judged onaccuracy, completeness, and on howeffectively information is communicated.Keep in mind that to communicate youfirst have to gain a person's attention.Your only restriction is that the informa-tion must be displayed on a single sheetof bristol board.

The following is a list of somecellsibatteries: Volta's Pile; Leclanche cell(common D cell); lead-acid battery (carbattery); "Energizer"I"Duracell" (modifiedLeclanche); silver oxide cell; solid-electrolyte cell (lithium iodine); fuel cells;solar battery; nickel-cadmium battery(rechargeable); other rechargeablebatteries; nuclear battery; sodium/sulfur battery; biological battery; plasticbattery (using plastics that conductelectricity); mercury or Ruben Mallorycell; indium/bismuth cell; silver chloridecell; liquid sodium or lithium batteries;Danniel cell; your choice.

Further ReadingParry, Robert W., et al. Chemistry.

Experimental Foundations. EnglewoodCliffs, NA: Prentice-Hall, 1975.

Sienko, Michell J., and Robert A. Plane.Chemistry, fifth edition. New York:McGraw-Hill, 1976.

Silk, T Electrochemistry. London:Heinemann Educational Books, 1968.

Toon, E.R., and G.P Ellis. Foundationsof Chemistry. New York: Holt, Rinehartand Winston, 1973.

10. Describe what happened when theelectrodes were partially removedfrom the solutions. Suggest plausibleexplanations.

11. Describe what happened when thesolution of sodium sulfide was addedto the copper half-cell. Suggest plausible explanations.

Information to gather:1. Use half-reactions and descriptions of

the processes occurring to describehow the cell/battery operates.

2. Describe how it is constructed andwhy it was made that way.

3. Describe ways people use thecell/battery.

4. Describe the characteristics of thecell/battery that make it suitable for itsuses (i.e., current produced, voltage,time before it wears out, etc.).

11

1 2

ObjectivesStudents should be able to:1. Describe the interior of a Leclanche cell and identify the

anode, cathode, and salt bridge.2. Explain the function of the electrolyte, cardboard cover,

and steel jacket in the Leclanchd cell.3. Explain how a Leclanche cell could perforate and result in

damage to the toy or flashlight in which it is used.4. Write chemical equations to describe the reactions that

occur at the anode and cathode of the Leclanche andnickel-cadmium cells.

5. Describe what happens when a nickel-cadmium recharge-able cell is recharged.

Program DescriptionThis program begins by pointing out that while electro-chemical cells made of beakers and solutions are helpful indetermining a table of reduction potentials, they are notpractical for everyday use. The program then turns to thehistory of the development of batteries. Volta's pile andGalvani's experiments with frogs' legs are described. Thenthe Leclanche cell is examined. The components of theLeclanche cell are described and compared to the com-ponents of the electrochemical cells made of beakers. Next,the reactions at the anode and cathode are explained. Thedifficulty with a Leclanche cell is easily understood: becausethe anode is actually the zinc container of the cell andbecause the oxidation reaction at the anode gradually con-sumes the zinc, eventually the container may perforate,allowing the highly corrosive electrolyte to leak out. Tworesponses to this problem are then described: 1) the develop-ment of the rechargeable nickel-cadmium cell, and 2) theimprovement of the Leclanche cell by wrapping the zinc canwith an absorbent cardboard casing and sealing it with asteel jacket. The program ends by emphasizing the fact thateven a steel jacket cannot hold in the corrosive contentsforever-many a toy and flashlight has been ruined by oldleaking batteries.

Before ViewingIt is important that the students know the basics of electro-chemical cells before viewing this program, because theanode and cathode reactions are quite complex. It would alsobe best to wait until the students have completed theirassignment on commercial electrochemical cells, outlined inthe suggestions for Program 3. Then it would be possible toshow this program in conjunction with student presentationson some of the other cells. This program could be shown outof sequence, if desirable, as the information in Programs 5and 6 does not depend on it.

After ViewingHave students present their descriptions of other cells.I nvolve the rest of the class by having them compare the cellspresented to the Leclanche cell. If possible, have dissectedcells for the students to look at. You may also wish toencourage students to test the claims of various batterymakers. Is it worth the extra cost to purchase Brand X?Another possible challenge would be to build a betterLeclanche cell. Students could certainly investigate dis-charge times for various electrolytes, various compositions ofMn02 and carbon, shapes and placement of electrodes, etc. Ifthey could actually improve on the design, they wouldbecome instantly wealthy. What could be better incentive fora project?

A Battery Primer

The following is a brief description ofmost of the common batteries with whichstudents come in contact. See the sum-mary table for all chemical reactions,electrolytes, and uses of these differenttypes of cells. Although there are manyelectrochemical cells in mass production,little is known about the actual chemicalreactions that are occurring in them.Therefore, the chemical equations given

Commercial Electrochemical Cells

for each type should be treated only asmodels. These equations provide the bestexplanations we have at present. Itshould also be noted that salt bridges arenot used to separate the cathode andanode in commercial cells. One commonmethod of separation is simply the use ofpaper soaked in an electrolyte. Thisprevents the shorting out of the battery,but allows the migration of ions neces-sary to complete the circuit.

A. Storage Batteries

A storage battery is a device which storesenergy in the form of chemical energy.The input and output of energy is in theform of electricity. As the voltage in abattery drops during discharge, the move-ment of the electrons-the current-decreases. One of the most commonstorage batteries is the lead-acid batteryused in most cars. The voltage producedby each lead-acid cell is 2 V The popular

Alternatingplates ofPb and Pb02

12 V car battery is thus made up of 6cells.

Another common storage battery is thenickel-cadmium (NiCd) battery. Thesehave become very popular with the adventof relatively inexpensive battery chargers.They are commonly used in cameras,flashguns, "Walkmans," and smokealarms. The working voltage of each NiCdcell is 1.2 V Thus the typical AA batteriesused in a "Walkman" consist of only onecell, but the 9 V battery used in a smokealarm consists of 8 cells.

B. Primary Batteries

A primary battery is designed to deliveronly one continuous or intermittentdischarge. It cannot be recharged.Primary batteries are classified by thetype of electrolyte used.

Aqueous-electrolyte cells

These batteries use solutions of acids,bases, or salts in water as the electrolyte.

1. Dry CellsThe commonly used dry cells and alkalinecells-Leclanche cells-are batteries ofthis type. Dry cells are batteries in whichthe acqueous electrolyte is a paste ratherthan a solution. In the paste there is suf-ficient water to allow ions to migrate butnot enough to allow the electrolyte tospill. A dry cell has a working potential ofbetween 0.9 and 1.4 V As the battery dis-charges, zinc is consumed. Since thebattery container is usually the zincanode, this can often lead to perforationof the container and the resultantdestruction of, for example, the flashlight.At the cathode, manganese (IV) oxide is

Lead-acid Storage Cell

1 3

1 4

converted into other products which differunder different loads. The equation in thetable is only one of several possibilities.I t should be noted once again that,although some of these batteries havebeen around since the 1860s, the

reduces the possibility of polarization.Reducing polarization allows more com-plete use of the reactive materials in thecell and results in a longer working lifewhen the cell is in continuous use. Notethe operative phrase here is "continuoususe,' as all the ads show. It would beinteresting to find out if these cellsshowed any advantage for the extra costwhen they are used for only short periodsof time.

2. Mercury CellsAnother type of aqueous-electrolyte cellis the mercury cell, also known as theRuben Mallory cell. When this type of celldischarges, zinc is consumed andmercury(II) oxide is converted intoelemental mercury. There is a buildup ofpositive charge in the anode due to the

Mercury cell

This would happen almost instantly ifthere were not a counteracting ioniccurrent in which K1 + i ons migrate fromthe anode to the cathode and OH1 - ionsflow in the opposite direction. Theelectrolyte in a mercury cell (a porousmaterial soaked in a concentrated solu-tion of potassium hydroxide and zincoxide in water) prevents electrons from

Leclanche cell

chemical reactions are not always wellunderstood-different sources give dif-ferent equations.

Alkaline dry cells, e.g., the "Energizer"and "Duracell,"are a relatively new modi-fication of the Leclanche cell. The anodeand cathode reactions are similar tothose in a regular dry cell, but the electro-lyte used is different. (See table.) Also, thezinc anode is in the centre of the cell.This avoids the problem of perforation.The central zinc cylinder is surrounded bya manganese(IV) oxide-carbon mixturewhich acts as the cathode. Thus, thecathode in an alkaline cell is much larger.These two changes-increasing the sizeof the cathode and changing the electro-lyte-facilitate the migration of ions. This

removal ofof negative charge i n the cathode due tothe production ofbuildup persisted, the battery would bepolarized and the current would cease.

There is also a buildup

. I f such a charae

passing directly from the zinc to the mer-cury(II) oxide. Without the electrolyte, thezinc electrons would not - have to passthrough the external circuit at all. Theoperating potential of the mercury celli s 1.3 V

3. Silver Oxide CellsAnother common battery is the silveroxide cell. The electrolyte is a 50%aqueous solution of potassium hydroxide.

Silver oxide cell

The operating potential of the silver oxidecell is about 1.5 V. These are used incameras, watches, and calculators.

Solid-electrolyte Cells

The solid state primary battery enjoyingthe greatest commercial success is theone used in pacemakers. The electrolyteused is solid lithium iodide. The move-ment of ions through the solid can bepictured in the same way as the move-ment of ions through a salt bridge. Ani odide ion produced at the cathode reactswith a nearby lithium ion, leaving ani odide ion in the crystal without a positivelithium ion. As a result, the iodide ionsteals one. This process continues to the

other side of the crystal where a lithiumion, produced at the anode, ultimatelypairs up with an iodide ion that has lostits lithium ion. As a result, there is amovement of positive ions from theanode to the cathode and of negativeions from the cathode to the anode. Obvi-ously, this ion movement meets far moreresistance than do ion migrations inliquid electrolytes. As a result, thesebatteries are only capable of producingsmall currents. However, they have theadvantages of not corroding and of beingvery small; hence their popularity in medi-cal applications.

C. Fuel CellsThese cells have two intriguingadvantages. First, they are capable ofextracting approximately twice the elec-trical energy from a fuel as we cancurrently get by burning that same fuel toproduce steam to power turbines to, inturn, produce electricity. Second, wasteheat is reduced by a factor of two. How-ever, in spite of considerable research onthese devices, at present we have nocommercially practical high-energy fuelcell.

I n this type of cell, an aqueous solutionof potassium hydroxide is used as theelectrolyte. As the cell discharges, thereis a buildup of OHM - ions at the cathodeand their corresponding depletion at theanode. This creates an ionic current ofpotassium ions flowing to the cathodeand of hydroxyl ions flowing to the anode.This type of battery was used on theApollo moon flights.

15

The ionic current is a flow of Ithe cathode and +e, +hc nne%An

to

Commercial Electrochemical Cells

Type ofBattery

STORAGE(a) Lead-Acid

(b) Nickel-Cadmium

PRIMARY(i) Aqueous-electrolyte

cells(a) Leclanche (dry cell)

(b) Modified Leclanche (alkaline dry cell)

(c) Ruben-Mallory( mercury cell)

(d) Silver oxide

(ii) Solid-electrolyte cells

(iii) Fuel cells(Hydrogen/oxygen)

Chemical Reactions While Discharging

OperatingPotential Common

(V) Electrolyte Uses

Further ReadingAshcroft, Ken, and Joe Hammill.

Chemistryof the Car. Agincourt, Ont.:The Book Society of Canada, 1975.

Brady, James E., and Gerard E. Humiston.General Chemistry. Principles andStructure. New York: John Wiley &Sons, 1982.

Parry, Robert W., et al. Chemistry.Experimental Foundations. EnglewoodCliffs, N.J.: Prentice-Hall, 1975.

Sienko, Michell J., and Robert A. Plane.Chemistry, fifth edition. NewYork:McGraw-Hill, 1976.

Silk, T Electrochemistry. London:Heinemann Educational Books, 1968.

Toon, E.R., and G.P. Ellis. Foundationsof Chemistry. New York: Holt, Rinehartand Winston,1973.

Williams, Dr. lolo Wyn. ChemistryandElectricity. London: Longmans/Penguin Books, 1969.

16

CorrosionObjectivesStudents should be able to:1. Use half reactions to help describe how a drop of water sit-

ting on iron acts as an electrochemical cell.2. I dentify the conditions necessary for rusting to occur.3. List two methods of preventing rust.4. Describe the processes of galvanization and sacrificial

anode protection.5. Use a table of standard reduction potentials to explain

how galvanization and a sacrificial anode protect iron.

Program DescriptionTo this point the "Electrochemistry" series has focussed oni mproving the robot's "heart" by designing a better battery.Program 5 looks at the robot's "skin" problem-corrosion. Itexplains that corrosion is the result of the same electro-chemical processes that produce power in a battery. In fact, adrop of water sitting on an iron surface creates an electro-chemical cell. The edge of the water drop forms the cathodewhere oxygen (easily replenished by the atmosphere) isreduced. The oxygen gains electrons from the iron at thecentre of the drop, which thus forms the anode. The two halfreactions described are as follows:

principle: a substance with a greater oxidation potential thani ron will react instead of the iron with the oxidizing agent.Thus, the program explains, the same process that can des-troy the robot's "skin" can also be used to protect it.

Before ViewingDo Activity 1. This activity must be left at least overnight. Havestudents set up the test tubes and agar plates one day. On thenext they can make their observations and begin working onthe discussion questions. Develop a chart of their results onan overhead projector or on the board. Briefly discuss studentanswers, but only evaluate them in terms of the data gathered.Then show the video.

After ViewingUse the information provided in the program to take anotherl ook at the students' observations and their answers to thediscussion questions for Activity 1. Point out that althoughmuch is known about the corrosion of iron, it is not fullyunderstood. Again, this topic is rich in its potential forstudent projects and research. Billions of dollars every yearare spent trying to cope with the rusting of iron. An interestedstudent would find The Rusting of Iron: Causes and Controlby U.R. Evans an excellent starting point.

Polarization does not occur because the drop of water con-tains other substances which act as an electrolyte. Thesesubstances include dissolved sodium chloride (in the case oftears) and dissolved gases such as carbon dioxide or sulfurdioxide (in the case of rainwater). Rusting takes place whenhydroxyl ions (created at the cathode) interact with the iron(II)i ons (formed at the anode) to form iron(li) hydroxide, whichprecipitates from the solution. The iron(ll) hydroxide then oxi-dizes to form rust, a hydrate of iron(III) oxide. Thus rusting isthe result of the operation of an electrochemical cell.

The program then examines two methods of protecting ironfrom rusting: galvanization and sacrificial anode protection.Both methods, it explains, work according to the same

ActivitiesActivity 1:Some Investigations intothe Corrosion of IronEach year billions of dollars are spent asa result of the rusting of steel. In thisinvestigation you will explore some of thefactors that affect rusting and considerways of reducing it.

17

1 8

Materials

1. 9 clean iron nails. Use emery cloth toclean if nails or other metals are notshiny.

2. thin strip of zinc about 8 cm long and2 mm wide or a chunk of mossy zinc

3. a strip of magnesium ribbon, 8-10 cmlong

4. 10 cm bare copper wire5. 2 g of powdered agar6. 0.1 kmol/m3 solutions of the following

substances (the solutions should bestored in dropper bottles): iron(II)

7. 5 test tubes and test-tube stand8. litmus or pH paper9. 250 mL beaker

10. ring stand, ring, gauze, Bunsenburner, striker, and pad, or hot plate

11. 2 petri dishes12. distilled water

NOTE: If dionized water is used, itshould be boiled to remove excessacidity caused by dissolved carbondioxide.

Procedure

CAUTION Many of the solutions in thisactivity are corrosive and toxic. Wearprotective gloves.

Part I:Reactions of Iron withVarious Aqueous Reagents1. Place a clean, bright nail in each of 5

test tubes. Carefully slide each naildown the side to avoid breaking thetube.

2. Partially fill each of the test tubes withone of the following reagents so thatthe nail is just covered. Your teacherwill tell you which of the groups to use.All solutions are 1.0 kmol/m3.

3.

4.

5.

Determine if the solutions are acidic orbasic using the litmus or pH paper. Goon to Part II.Allow the nails to stand overnight inthe solutions. Observe and record anychanges that have taken place. Com-pare your results with those ofstudents using the other groups ofreagents. Record their results alongwith yours in tabular form.After the solutions have stood over-night, add one or two drops ofpotassium ferricyanide solution to

each test tube. Observe and record anychange.

6. Add one drop of potassium ferri-cyanide to about 1 mL of iron(II) sul-fate solution. Compare this result withthose obtained in step 5. What conclu-sion can be made based on the resultsof step 5?

Part II:Reactions Involving Iron inContact with Other Metals1. Prepare about 200 mL of an agar

mixture as follows: Heat about 200 mLof distilled water to a gentle boil.Remove the burner and stir in 2 g ofpowdered agar. Continue the heatingand stirring until the agar is uniformlydispersed.NOTE: To save time, the instructor mayprepare enough agar solution for theclass.

2. Add about 10 drops of potassium ferri-cyanide and 5 drops of phenol-phthalein to the agar mixture. Stirthoroughly.

3. While the agar is cooling, prepare fourclean, bright nails as shown in the dia-gram. Place one nail on one side of apetri dish. Bend another sharply with a

sulfate

pair of pliers and place it on the otherside of the same dish.

4. Twist a clean piece of bare copper wirearound a third nail. Then remove thenail and tighten the wire coil so thatwhen the nail is forced back through,the wire makes tight contact with thenail. Place this in a second dish.

5. Repeat step 4 using a zinc ormagnesium strip. If a strip of zinc isnot available, push the nail through apiece of mossy zinc in at least twoplaces. Place this nail in the secondpetri dish as shown in the diagram. Besure the nails do not touch.

6. When the agar has cooled to lukewarmand is still fluid, carefully pour it intothe petri dishes until the nails andmetals are covered to a depth of about0.5 cm.

7. Make observations during the remain-i ng class time. Store the dishes over-night in the place designated by yourteacher. Record any changesobserved next day.

Discussion1. List the reagents in Part I where no in-

dication of rusting was observed.2. List the reagents in Part I where there

was an indication of rusting.3. What regularity exists in the data?

Suggest a plausible explanation.4. Where on the nail did the largest

amount of rusting take place?Suggest a plausible explanation.

5. Why can a nail stay on a store shelffor months and not rust, but rust over-night in tap water.

6. Describe a test for iron(ll) ions. Writea chemical equation for the test.

7. Describe the color changes in Part 11.Suggest plausible explanations.

8. Write half-reactions to account forthe color changes in Part 11.

9. Use reduction potential tables topredict the color changes ifmagnesium ribbon had been usedinstead of zinc.

10. What is the name given to themethod of protection provided by thezinc. Describe how it works.

11. What is the name given to themethod of protection provided bymagnesium. Describe how it works.

12. Magnesium blocks are sometimesplaced in hot water heaters, riveted tothe hulls of ships and attached toiron pipes buried underground. Why?What ultimately happens to theseblocks? What is the name given tothis method of protection?

13. The table of reduction potentialssuggests that copper is more difficultto oxidize than iron. Based on this,copper should make an excellentprotective coating for iron. Explainwhy it is not used.

14. In 1986 New Yorkers celebrated therestoration of the Statue of Liberty.The statue is actually copper supported by an iron framework. Whathad to be restored? Explain yourreasoning.

15. Chrome plating involves coating theiron with a relatively thick layer ofnickel first, then applying a very thinl ayer, or flash, of chrome. Examinethe table of reduction potentials andsuggest a plausible reason for doingthis.

Further ReadingAshcroft, Ken, and Joe Hammill.

Chemistry of the Car. Agincourt, Ont.:The Book Society of Canada, 1975.

Brady, James E., and Gerard E. Humiston.General Chemistry. Principles andStructure. New York: John Wiley &Sons, 1982.

Evans, Ulick R. The Rusting of Iron:Causes and Control. London: EdwardArnold, 1972.

Parry, Robert W., et al. Chemistry:Experimental Foundations. EnglewoodCliffs, N.J.: Prentice-Hall, 1975.

Sienko, Michell J., and Robert A. Plane.Chemistry, fifth edition. New York:McGraw-Hill, 1976.

Silk, T Electrochemistry. London:Heinemann Educational Books, 1968.

Toon, E.R., and G.R Ellis. Foundationsof Chemistry. New York: Holt, Rinehartand Winston, 1973.

Williams, Dr. lolo Wyn. Chemistry andElectricity. London: Long mans/PenguinBooks, 1969.

1 9

20

Electroplating

ObjectivesStudents should be able to:1. Compare the process of electrolysis and the action of

electrochemical cells.2. State one use of electrolysis.3. Use reduction potentials to explain why electrolysis

reactions require a battery.4. Describe how an object could be plated with copper.5. Use half reactions to explain what is happening when

copper is electroplated on an object.6. Given reduction potentials, predict the products of an

electrolysis reaction.

Program DescriptionAfter reviewing the corrosion process, galvanization, and theuse of sacrificial anodes, this program goes on to examineanother electrochemical process that could be used to pro-tect an iron surface-electroplating. It argues that since achemical reaction can be used to produce electricity, itshould be possible to use electricity to produce a chemicalreaction. In other words, it should be possible to use elec-tricity to coat a robot with a protective metal, e.g., chromium.The program then tests this idea by exploring what wouldhappen if an electric current were passed through a solutionof copper(II) sulfate using carbon electrodes. After describ-

not occur spontaneously-electricity is needed to bring itabout.

Thus, the program explains, it would be possible to coatthe robot with copper by making the robot the cathode in theelectrolysis of copper(II) sulfate. This process, known aselectroplating, could also be used with a different solution tocoat the robot in chrome. The program ends by pointing outthat electrochemistry is both the curse and the cure of iron.

Before ViewingDo Activity 1, a series of electrolysis and electroplatingreactions. Students apply chemical tests to identify theproducts of the reactions. Assign the discussion questions butdon't take them up until after viewing the program.

After ViewingStudents should use the information provided in the programto take another look at their observations and answers to thediscussion questions. You should also discuss why it wouldbe unwise to coat iron with a less active metal like copper. Ifthere is an electroplating shop in your community, it would beinteresting to do a case study of the shop. Much is knownabout the electroplating process, but the function of many ofthe ingredients in a commercial electroplating bath is notunderstood. Different ingredients are only known to make theplate "shinier" or "stick better." Also, many of the chemicalsused are toxic and some shops still do not dispose of themproperly.

ActivitiesActivity 1:Electrolysis ofSome CompoundsI n this activity you will investigate theeffect of passing current through anumber of solutions. Because of the

the slight ionization and dissociation of water), reductionpotentials are used to predict which species at each elec-trode will react. The reduction potential is considered first.Because the copper ion has the greatest potential to bereduced, copper plates out on the negative electrode. Theprogram points out that because reduction occurs at thenegative electrode, it is called the cathode. The reaction atthe positive electrode is then described. Having shown howthe two half reactions are decided upon, the programexplains how they can be added together to give the overallequation for the electrolysis of copper(II) sulfate. Since thereaction potential turns out to be negative, the reaction will

ing the species in the solution ( from thefromionic compound copper(II) sulfate, and

amount of equipment required, each pairof students will set up one electrolysisreaction. Your teacher will inform youwhich part you will do and how the infor-mation will be shared.

Materials

1. 9 U-tubes

2. the following solutions:

to color the solution deeply, usedroplets of acid and base to adjustthe litmus color to violet, indi-cating a neutral solution.

- saturated NaCI(320 g NaCI/L) andenough bromothymol blue to colorthe solution deeply. Use dropletsof acid and base to adjust thecolor to green, indicating a neutralsolution.

- 0.1 kmol/m3 NaCl (6 g NaCI/L) andenough bromothymol blue to colorthe solution deeply. Use dropletsof acid and base to adjust thecolor to green, indicating aneutralsolution.

3. the following indicator solutionsstored in dropper bottles:

4. 18 carbon electrodes fitted in 2-holerubber stoppers that fit the mouths ofthe U-tubes

5. 9 ring stands and adjustable clamps6. 9 D.C. power supplies, 0-12 V, and 18

connecting leads (9 black, 9 red) withalligator clips

7. 16 medicine droppers

8. 150 mL beaker

9. salt bridge of 16 mm glass tubingthat will fit between arms of a U-tube,cotton batten to stopper ends ofU-tube

10. voltmeter, high impedance, 0-5 V

11. 5 spot plates12. masking tape for labelling arms of

U-tube

Procedure

CAUTION It is possible chlorine gas willbe produced when electrolyizingchlorides. Chlorine is toxic; adequateventilation is required. Perform the elec-trolysis of chlorides in a fumehood ifpossible.CAUTION Concentrated nitric aciddestroys flesh. Wear protective gloves.

Part I:Electrolysis of KI (aq)

1. Add enough potassium iodide solu-tion to fill the U-tube to within 2 cmfrom the top.

2.

3. Set up connecting wires and elec-trodes as in diagram. Mark on theU-tube the positive and negativesides.

4. Turn on the power supply and slowlyincrease the voltage until you observechanges taking place at the elec-trodes. Do not exceed 12 V

Use the adjustable clamp to mountthe U-tube on the ring stand. Seediagram.

2 1

- universal indicator- starch solution(4 g starch dis-

solved in 1 L of boiling water)- TTE, trichlorotrifluoroethane, or

hexane- concentrated nitric acid

and enough solid litmus

22

5. Record any changes that occur at theelectrodes.CAUTION Use the hand wave methodfor detecting odor.

6. After running the electrolysis for5-10 min, turn off the power supplyand remove the electrodes. Note anyodor.

7. Use the medicine dropper to draw offsome of the solution from both sidesof the U-tube. Rinse the dropperbefore switching sides. Add thedrawn off solution to the spot plate.Check the steps that follow to deter-mine how many spots of each solu-tion you will need.

8. Test both solutions with 2-3 drops ofstarch solution, iron (III) chloride solu-tion and universal indicator.

9. Use the medicine dropper to transferapproximately 2 mL of solution fromthe negative side of the U-tube into atest tube. Do the same for the posi-tive side but use a different test tube.

10. Add approximately 1 mL of TTE orhexane to the test tubes, shake andallow the TTE to settle or the hexaneto rise. Note the color changes.After everyone has noted the observa-tions, pour the TTE(hexane)lsolutionmixture in the test tubes into thebottle labelled TTE(hexane) wastes.DO NOT POUR MIXTURE DOWN THESINK.

11. When everyone has had an oppor-tunity to view your results, wash thespot plate and U-tube with copiousquantities of water.

Part I1:Electrolysis ofSaturated NaCI(aq)1. Add enough saturated sodium chloride

solution to fill the U-tube to within2 cm from the top.

2. Use the adjustable clamp to mount theU-tube on the ring stand. See diagram.

3. Set up connecting wires and elec-trodes as in the diagram. Mark on theU-tube the positive and negative sides.

4. Turn on the power supply and slowlyincrease the voltage until you observechanges taking place at the elec-trodes. Do not exceed 12 V.

5. Record any changes that occur at theelectrodes.

6. Test the gas produced at the positiveside with a glowing splint.CAUTION Use the hand wave methodto detect odor.

7. After running the electrolysis for5-10 min, turn off the power supply andremove the electrodes. Note any odor.

8. When everyone has had an opportunityto view your results, wash the U-tubewith copious quantities of water.

Part III:Electrolysis ofDilute NaCl (aq)1. Repeat Part II using 0.1 kmol/m3 NaCl

i nstead of saturated NaCl.

1. Soak cotton batten in a small amountof the 0.1 kmol/m3 NaCl solution.

2. Push the ball of cotton to the centreof the "U" of the U-tube. The ball ofcotton should effectively seal thepassage between the two arms of theU-tube.

3. Add enough of the FeCl2 solution tofill one side of the U-tube to within2 cm from the top.

4. Add enough of the FeCl 3 solution tofill the other side of the U-tube towithin 2 cm from the top.

5. Use the adjustable clamp to mountthe U-tube on the ring stand. Seediagram.

6. Set up connecting wires and elec-trodes as in the diagram. Mark on theU-tube the positive and negativesides. Make the side containing theFeCl 3 negative and the side con-taining the FeCl 2 positive.

7. Turn on the power supply and slowlyi ncrease the voltage until you observechanges taking place at the elec-trodes. Do not exceed 12 V

8. Record any changes that occur at theelectrodes.

9. After running the electrolysis for5-10 min, turn off the power supplyand connect the positive side of thepower supply to the negative side ofthe U-tube and the negative side ofthe power supply to the positive sideof the U-tube. Turn on the power sup-ply for an additional 5-10 min. Recordany changes that occur.

Part IV:Electrolysis of

10. When everyone has had an oppor-tunity to view your results, wash theU-tube with copious quantities of water.

CAUTION The gases produced in the nextsteps are toxic. Adequate ventilation isrequired.

2. In addition, after removing the elec-trodes, place one in a 150 mL beakerand add 2-4 drops of concentratednitric acid to the part of the electrodethat was immersed in the solution inthe U-tube. Record any changes.

3. Repeat step 2 for the other electrode,using a clean beaker.

1. Repeat Part I usingCAUTION The gases produced in the nextsteps are toxic. Adequate ventilation isrequired. Do these procedures in a fume-hood if possible.

2. In addition, after removing the elec-trodes, place one in a 150 mL beakerand add 2-4 drops of concentratednitric acid to the part of the electrodethat was immersed in the solution inthe U-tube. Record any changes.

3. Repeat step 2 for the other electrode,using a clean beaker.

CAUTION The gases produced in the nextsteps are toxic. Adequate ventilation isrequired. Do these procedures in a fume-hood if possible.

2. In addition, after removing the elec-trodes, place one in a 150 mL beaker,add 2-4 drops of concentrated nitricacid to the part of the electrode thatwas immersed in the solution in theU-tube. Record any changes.

3. Repeat step 2 for the other electrode.

Part IX:Electrolysis of HCI Solution.CAUTION The hydrogen and chlorine pro-duced in this activity explode when mixedi n the presence of light. Adequate ventila-tion is required and the apparatus shouldbe behind an explosion shield. Weargoggles.

1. Add enough of the HCI solution to fillthe U-tube to within 2 cm from the top.

2. Use the adjustable clamp to mount theU-tube on the ring stand. See diagram.In addition, place a piece of cardboard30 cm wide between the two arms ofthe U-tube, and extending 50 cm abovethe tube to help keep the gases frommixing.

3. Set up connecting wires and elec-trodes as in the diagram. Mark on theU-tube the positive and negative sides.

4. Turn on the power supply and slowlyi ncrease the voltage until you observechanges taking place at the elec-trodes. Do not exceed 12 V

5. Record any changes that occur at theelectrodes.

6. After running the electrolysis for5-10 min, turn off the power supply andremove the connections from thepower supply. Instead connect them tothe voltmeter. If possible adjust thevoltmeter so the full scale deflection is2-5 V Record any reading on the volt-meter. Record any changes that occuri n the U-tube.

6. When everyone has had an opportunityto view your results, wash the U-tubewith copious quantities of water.

Discussion

1. Indicate the evidence you used to iden-tify the following as products of elec-

2. In each case which side of the U-tubewas the cathode? Which was theanode?

3. For each part write half reactions toexplain the changes observed in theanode and cathode parts of the U-tube.

4. Write half reactions to account for thevoltmeter reading observed in Part IX.What is the U-tube behaving as? Whatdoes this imply about electrochemicalcells and electrolysis?

23

trolysis

1. Repeat Part II usingsaturated NaCl.

i nstead of

i nstead of KI.

Part VII:Electrolysis of

Solution.

24

Further ReadingAlyea, Hubert N., and Frederick B. Dutton.

Tested Demonstrations in Chemistry.Easton: Chemical Education Publish-i ng, 1965.

Ashcroft, Ken, and Joe Hammill.Chemistry of the Car. Agincourt, Ont.:The Book Society of Canada, 1975.

Brady, James E., and Gerard E. Humiston.Genera/ Chemistry: Principles andStructure. New York: John Wiley &Sons, 1982.

Evans, Ulick R. The Rusting of Iron:Causes and Control. London: EdwardArnold, 1972.

Parry, Robert W., et al. Chemistry:Experimental Foundations. EnglewoodCliffs, N.J.: Prentice-Hall, 1975.

Parry Robert W., et al. Teacher's Guide.Chemistry. Experimental Foundations.Englewood Cliffs, N.J.: Prentice-Hall,1975.

Sienko, Michell J., and Robert A. Plane.Chemistry, fifth edition. New York:McGraw-Hill, 1976.

Silk, T Electrochemistry. London:Heinemann Educational Books, 1968.

Toon, E.R., and G.P Ellis. Foundationsof Chemistry. New York: Holt, Rinehartand Winston, 1973.

Williams, Dr. lolo Wyn. Chemistry andElectricity. London: Longmans/PenguinBooks, 1969.

OrderingIInformationTo order this publication or videotapes of theprograms in the Electrochemistry series, or foradditional information, please contact one ofthe following:

CanadaTVOntario Sales and LicensingBox 200, Station QToronto, Ontario M4T 2T1(416) 484-2613

United States U.S. Sales Office901 Kildaire Farm RoadBuilding ACary, North Carolina27511Phone: 800-331-9566Fax: 919-380-0961E-mail: ussales@tvo.org

Videotapes BPNProgram 1: The Building Blocks of

Electrochemistry 260501Program 2: Electrochemical Cells 260502Program 3: Designing Electro-

chemical Cells 260503Program 4: Commercial Electro-

chemical Cells 260504Program 5: Corrosion 260505Program 6: Electroplating 260506