topic redox reactions, chemical cells and...

Redox Reactions, Chemical Cells and Electrolysis

Unit 18 Chemical cells in daily life

Unit 19 Simple chemical cells

Unit 20 Oxidation and reduction

Unit 21 Oxidation and reduction in chemical cells

Unit 22 Electrolysis

Topic 5

KeyC o ncepts

Chemical cells in daily life• Different typesof chemical cells• Choosing a chemical cell for a

particularuse• Environmental impactofusing

chemical cells

Simple chemical cells• Reactions in simple chemical cells• Electrochemical series ofmetals• Salt bridge in simple chemical cell• The Daniell cell

Oxidation and reduction• Definitions of oxidation and reduction

(redox reaction)• Common oxidizing and reducing

agents• Oxidizing and reducingpowerof

various species• Oxidation number

Electrolysis• Factors affecting theorderofdischarge

of ionsduring the electrolysis ofaqueous solutions

• Industrial usesof electrolysis• Environmental impactof the

electroplating industry

Redox Reactions, Chemical Cells and

Electrolysis Oxidation and reduction in chemical cells

• Redox reactions in various chemicalcells

• Redox reactions in everyday life

Topic 5� Redox Reactions, Chemical Cells and Electrolysis �Unit 18 Chemical cells in daily life

18.1 – 18.12

Summary

1 A chemical cell is a device in which chemical energy is converted into electricalenergy. It consists of twodifferentmetals and an electrolyte.

Chemical energychemical cell

electrical energy

2 There are two classes of chemical cells: primary cells and secondary cells. Primarycells arenot rechargeable. Secondary cells are rechargeable.

3 The following table summarizes the characteristics of six commonchemical cells.

Zinc-carbon Alkaline manganese

Silver oxide

Lithiumion

Nickel metal hydride

Lead-acid accumulator

Type primary primary primary secondary secondary secondary

Materialofnegativeelectrode

zinc zinc zinc

lithiumatoms lying

betweengraphitesheets

hydrogenabsorbing

alloyslead

Materialof positiveelectrode

carbon manganese(IV)oxide

silveroxide

lithiummetal oxide

nickel(II)hydroxide

leadplatescoatedwith

lead(IV)oxide

Material ofelectrolyte

ammoniumchloride

potassiumhydroxide

potassiumhydroxide

lithium saltdissolved in

organicsolvent

potassiumhydroxide sulphuric acid

Maximumvoltage (V) 1.5 1.5 1.5 3.7 1.2 2

Advantage(s)

• low cost

• long shelflife

• highefficiencyundermoderateandcontinuousdrains

• long shelflife

• good low-temperatureperformance

• flatdischargecurve

• lightweight

• small size

• highenergydensity

• low

self-discharge

rate

• goodperformanceinhigh-draineddevices

• good low-temperatureperformance

• low cost

• spillresistant(sealedtype)

Disadvantage(s)

• poorperformanceinhigh-draineddevices

• outputcapacitydecreases asthe cell is discharged

• poor low-temperatureperformance

• expensivefornoveltyusage

• silveris veryexpensive

• veryhighinitialcost

• relativelylowdischargecurrent

• safetyconcerns*

• high initialcost

• relativelyhigh rateof self-discharge

• limited low-temperatureperformance

• heavy

• lead andleadcompoundsare toxic

*A lithium ion cell isnot very robust and cannot takehigh charge anddischarge currents. Tryingto force a rapid charge or loading the cell with excess discharge current may overheat the packand lead to explosion.

18.1 Chemical cells indaily life

18.2 Electricity fromchemical reactions

18.3 Different typesof chemical cells

18.4 Terms related to chemical cells

18.5 Zinc-carbon cell

18.6 Alkaline manganese cell

18.7 Silver oxide cell

18.8 Lithium ion cell

18.9 Nickel metalhydride (NiMH) cell

18.10 Lead-acid accumulator

18.11 Choosing a chemical cell for aparticularuse

18.12 Environmental impactofusing chemical cells

Unit 18 Chemical cells in daily life

Topic 5� Redox Reactions, Chemical Cells and Electrolysis �Unit 18 Chemical cells in daily life

Example

The diagram below shows a traditional ‘button cell’ connected to an external circuit.The cell ismade frommercury(II) oxide, zincpowder and sodiumhydroxidepaste.

electrode X

zinc powder

sodium hydroxide paste

mercury(II) oxide

electrode Y

load

a) Stateandexplainthedirectionofelectronflowintheexternalcircuitwhenthecelldischarges. (2marks)

b)Write an equation for theoverall cell reactionduringdischarge. (1mark)

c) What is the functionof the sodiumhydroxidepaste in the cell? (1mark)

d)Why should this kindof button cell bebanned in themarket? (1mark)

e) Explain whether the cell can work if mercury(II) oxide is replaced by magnesiumoxide. (1mark)

f) Explain the change of the maximum voltage supplied by the cell if zinc powder isreplacedby copperpowder. (2marks)

Answer

a) FromelectrodeX to electrodeY (1)

because zinc atomsgiveup electronsduringdischarge. (1)

b)Zn+HgO ZnO+Hg (1)

c) Provide amedium for ion transfer. (1)

d)Mercury is toxic. (1)

e) No.

Anyoneof the following:

• Magnesium loses electronsmore readily than zincdoes. (1)

• Magnesium is a stronger reducing agent than zinc. (1)

f) Themaximumvoltagewill decrease. (1)

Copper and mercury are closer in the electrochemical series than zinc andmercury. (1)

ExamtipsExamtipsExamtipsExamtips ♦ Do NOT confuse the terms: anode, cathode, negative electrode andpositive electrode.

– The negative electrode of a chemical cell may also be called theanode. Oxidation occurs here.

– The positive electrode of a chemical cell may also be called thecathode. Reduction occurs here.

♦ The maximum voltage of a zinc-carbon cell is 1.5V, regardless of itssize.

4 Aspects to considerwhen choosing a chemical cell for aparticularuse:

• price;

• voltage;

• capacityof the cell;

• whether the cell is rechargeable;

• whether the cell is able to supply a steady current;

• whether a largeor small current is required;

• whether the cell is used continuouslyor intermittently;

• whether the cell is required tooperate at low temperatures;

• shelf life; and

• riskof leakage.

5 Disposalofchemicalcellsmaycauseenvironmentalproblems.Heavymetalcomponentsin chemical cells canbe toxic toplants, animals andhumans.

Topic 5� Redox Reactions, Chemical Cells and Electrolysis �Unit 19 Simple chemical cells

➤Questionsoftenshowunfamiliarchemicalcellsandaskaboutreactionsthatoccur at the electrodes.

e.g.

– Thediagrambelowshowsa sodium-sulphurcell thatoperatesatahightemperatureof370°C,whichisabovethemeltingpointsofsodiumandsulphur.

load

inertelectrode B

sulphur

inertelectrode A

porousdevice

sodium

At electrodeA:Na(l) Na+(l) + e–

At electrodeB: S(l) + 2e– S2–(l)

Electrons flow from electrode A to electrode B in the external circuitbecause sodiumhas a higher tendency to lose electrons than sulphur.

– The diagram below shows an ancient chemical cell made by immersingan iron rod in vinegar inside a copper can.

iron rod

copper canvinegar

multimeter

Theironroddissolvesgradually,andcolourlessgasbubblesaregivenoffon the innerwall of the copper can.

At the iron rod: Fe(s) Fe2+(aq)+2e–

On the innerwall of the copper can: 2H+(aq)+2e– H2(g)

RemarksRemarks*

19.1 Reactions in simple chemical cells

19.2 Simple chemical cellsmade fromdifferentmetal couples

19.3 The electrochemical series ofmetals

19.4 Improving simple chemical cells

19.5 The roleof a salt bridge

19.6 The Daniell cell

Simple chemical cellsUnit 19

Topic 510 Redox Reactions, Chemical Cells and Electrolysis 11Unit 19 Simple chemical cells

19.1 – 19.6

Summary

1 For a simple chemical cell made from two different metals, the metal which formsionsmorereadily is thenegativeelectrode.Themetalwhichforms ions lessreadilyis the positive electrode. Electrons flow from the negative electrode to the positiveelectrode in the external circuit.

positive electrode

electrolyte

negative electrode

electron flow

digital multimeter(as voltmeter)

2 Metals canbearranged inorderof their tendencies to form ions.This is called theelectrochemical series ofmetals.

3 Theorderofmetalsintheelectrochemicalseriesisthesameasthatinthereactivityseries (except calcium).

4 To increase the efficiency of a simple chemical cell, the cell can be separated intotwohalf-cells connectedby a salt bridge.

5 A salt bridge serves two important functions:

a) It completes the circuit by allowing ions to move from one half-cell to theother.

b)It provides ions that can move into the half-cells to prevent the build-up ofexcess positively or negatively charged ions in the solutions which would causethe reaction to stop.

6 IntheDaniellcell,electronsflowfromthezincstriptothecoppercontainerintheexternal circuit.

Zn(s) Zn2+(aq) + 2e–

Cu2+(aq) + 2e– Cu(s)

ExamtipsExamtipsExamtipsExamtips ♦ Questionsmaygiveinformationaboutthedisplacementreactionbetweenametalandtheionsofanothermetal,andaskwhatwillhappenwhenthemetals are used as electrodes in a lemon cell.

e.g.

XandYaremetals.MetalX reactswithY2+(aq) ions according to thefollowing equation:

X(s) +Y2+(aq) X2+(aq)+Y(s)

ItcanbededucedthatmetalXismorereactivethanmetalYandloseselectronsmorereadily.WhenmetalsXandYareusedaselectrodes ina lemon cell,

– electrons flow from X toY in the external circuit;

– X2+(aq) ions are found in the lemon juice;

– X acts as the anode (i.e. negative electrode);

– Y acts as the cathode (i.e. positive electrode).

metal Y

lemon

electron flow


metal X

♦ Questions maygive simple chemical cells as shown below:

V

metal X Cu

NaCl(aq)

electronflow

Cell 1

V

metal Y Cu

electronflow

Cell 2

Topic 512 Redox Reactions, Chemical Cells and Electrolysis 13Unit 19 Simple chemical cells

Example

The diagram below shows a set-up with metal strips inserted in fresh potatoes. Themultimeter reading in the set-up is +0.75V.


coppermagnesium

potato A potato B

zinccopper

a) State, with explanation, the direction of electron flow across the connecting wirebetweenmagnesium strip and zinc strip. (2marks)

b)What will happen to the multimeter reading when the zinc strip and copper stripinpotatoB is interchanged?Explainyour answer. (2marks)

c) Explainwhy freshpotatoes shouldbeused in the set-up. (1mark)

Answer

a) Electrons flow from themagnesium strip to the zinc strip (1)

becausemagnesium loses electronsmore readily than zinc. (1)

b)Themultimeter readingwill increase. (1)

The twopotato cells become connected in series. (1)

c) Freshpotatoes containwater so that ionsmovemore easily. (1)

➤The above diagram shows two competing chemical cells. In this situation,themore powerful cellwill drive theweaker cell into electrolysis.

➤Besides the lemon cell, the potato cell is another popular home madecell.

Two potato cells connected in series can be used to power an LCD clock.Notice that the positive electrode of one potato cell is connected to thenegative electrode of the other potato cell (i.e. the cells are connected inseries).

zincstrip

zincstrip

copperstrip

copperstrip

LCD clock

RemarksRemarks* Students need to deduce the relative tendency of metals X and Y to

form ionsanddecidewhatwill happen if theCu inCell 1 is replacedbymetal Y.

Among metal X, metal Y and Cu, metal Y forms ions most readily.Metal Y and metal X are farther apart in the electrochemical seriesthanCu andmetal X.

If theCu inCell 1 is replacedbymetal Y,

– electrons will flow frommetal Y tometal X in the external circuit;

– a greater magnitude of voltage will be recorded.

Topic 51� Redox Reactions, Chemical Cells and Electrolysis 1�Unit 20 Oxidation and reduction

20.1 Oxidation and reduction in termsof gain and lossofoxygen

20.2 Oxidizing agent and reducing agent

20.3 Oxidation and reduction in terms of gain and loss ofhydrogen

20.4 Oxidation and reduction in termsof electron transfer

20.5 The reducingpowerofmetals

20.6 Oxidation numbers

20.7 Using oxidationnumbers to identify redox reactions

20.8 Advantagesanddisadvantagesofusingtheconceptofoxidationnumber

20.9 The Stock systemofnaming compounds

20.10 Chemical changesof commonoxidizing and reducing agents

20.11 Balancing redox equationsusing ionichalf-equations

20.12 Balancing redox equationsusingoxidationnumbermethod

20.13 Theelectrochemicalseriesandtherelativeoxidizing/reducingpowerof commonoxidizing / reducing agents

20.14 Chlorine as anoxidizing agent

20.15 Nitric acidofdifferent concentrations as oxidizing agents

20.16 Concentrated sulphuric acid as anoxidizing agent

20.17 Aqueous sulphurdioxide as a reducing agent

Oxidation and reductionUnit 20 20.1 – 20.13

Summary

1 The following table summarizes thedefinitionsof oxidation and reduction.

Defining in terms of Oxidation Reduction

Gainingor losingoxygen

process inwhich a speciesgainsoxygen

process inwhich a specieslosesoxygen

Gainingor losinghydrogen

process inwhich a speciesloseshydrogen

process inwhich a speciesgainshydrogen

Electron transfer process inwhich a speciesloses electron(s)

process inwhich a speciesgains electron(s)

Change inoxidationnumbers

process inwhich theoxidationnumberof an element in aspecies increases

process inwhich theoxidationnumberof an element in aspeciesdecreases

2 Oxidation and reduction always occur at the same time. The combined process iscalled a redox reaction.

3 A reducing agent is a species that causes reduction. An oxidizing agent is a speciesthat causesoxidation.

Species is a / an Species itself undergoes Species is Change in

oxidation number

reducing agent oxidation oxidized increase

oxidizing agent reduction reduced decrease

4 The oxidation number of an atom is an imaginary oxidation number assigned tothe atomaccording to a set of rules.

5 Balanced equations for redox reactions canbeobtainedby:

a) combining ionichalf-equations; and

b)oxidationnumbermethod.

6 In the electrochemical series,

a) theoxidizingpowerof oxidizing agents increasesdown the series;

b)the reducingpowerof reducing agentsdecreasesdown the series.

7 The electrochemical series canbeused topredictpossible redox reactions.


ExamtipsExamtipsExamtipsExamtips ♦ State oxidation numbers in the format of+1,+2, etc, NOT1+, 2+. ✔ ✔ ✘ ✘

+2 0 –1

♦ The oxidationnumber of cobalt inCo(NH3)4Cl2 is +2,NOT2. ✔ ✘

♦ Questionsoftenaskstudentstogivetheoxidationnumberofacertainelement in an unfamiliar species (e.g. the oxidation number of V inVO2

+).

♦ Questionsoftenaskabouttheoxidationnumberofsulphurindifferentsubstances.

Substance Oxidation number of S in the substance

H2S –2

Na2S2O3 +2

SO2 +4

Na2SO4 +6

H2S2O7 +6

♦ Questionsoftenaskabouttheoxidationnumberofnitrogenindifferentsubstances.

Substance Oxidation number of N in the substance

NH3 –3

NO +2

NO2 +4

HNO3 +5

♦ Questions often ask students to decide whether a certain equationrepresentsa redox reaction. Justdeterminewhether therearechangesin oxidationnumbers of the elements.

– The following are redox reactions:

CH4(g) +Cl2(g) CH3Cl(g) +HCl(g)

2KNO3(s) + S(s) + 3C(s) K2S(s) +N2(g) + 3CO2(g)

– The following areNOT redox reactions:

2CrO42–(aq)+2H+(aq) Cr2O7

2–(aq)+H2O(l)

CaCO3(s) CaO(s) +CO2(g)

Example

Hydrogenperoxide (H2O2) is anoxideofhydrogen.

a) Hydrogen peroxide decomposes at room temperature according to the followingequation:

2H2O2(aq) 2H2O(l) +O2(g)

Explain,intermsofoxidationnumbers,whythedecompositionisadisproportionationreaction. (2marks)

b)In the presence of a dilute acid, hydrogen peroxide oxidizes iodide ions to iodineand it is reduced towater.

i) Write an ionichalf-equation for the reductionofhydrogenperoxide. (1mark)

ii)State the expectedobservation andwrite an equation for the reaction involved. (2marks)

♦ Questionsoftenlistequationsofunfamiliarreactionsandaskstudentstoidentifytheoxidizingagents(i.e.speciesbeingreduced)orthereducingagents (i.e. species being oxidized).

e.g.

+4 0

SO2+2H2S 3S+2H2O

SO2 is theoxidizing agent; it is being reduced.

–3 0

2NH3+3CuO 3Cu +N2+3H2O

NH3 is the reducing agent; it is being oxidized.

+2 0

CuSO4+ Zn ZnSO4+Cu

CuSO4 is theoxidizing agent; it is being reduced.

2S2O32–(aq)+ I2(aq) S4O6

2–(aq)+2I–(aq)

I2 is theoxidizing agent; it is being reduced.

+4 +2

PbO2+NO2–+2H+ Pb2++NO3

–+H2O

PbO2 is theoxidizing agent; it is being reduced.

♦ Acidified potassium permanganate solution and acidified potassiumdichromate solution are important oxidizing agents. Do NOT confusetheir colour changes in redox reactions.

♦ Questionsoftenaskwhatwillhappenwhenadrunkendriverbreathesinto a breathalyzer containing dichromate ions.

When orange dichromate ions are brought into contact with ethanol,the following redox reaction occurs.

2Cr2O72–(aq)+3CH3CH2OH(aq)+16H+(aq)

4Cr3+(aq)+3CH3COOH(aq)+11H2O(l)

Green chromium(III) ions form.

The oxidation number of chromium changes from +6 to+3.


➤Questionsoftenaskstudentstowriteredoxequationsforunfamiliarreagents.Justcombineionichalf-equationsforchemicalchangesoftheoxidizingagentand reducing agent to obtain a balanced redox equation for a reaction.

RemarksRemarks*

c) In the presence of another powerful oxidizing agent, hydrogen peroxide can showreducingproperties.Inalkalineconditions,iron(III)ionsarereducedtoiron(II)ions.Oxygen andwater are alsoproduced.

i) Write an ionic half-equation for the oxidation of hydrogen peroxide in alkalineconditions. (1mark)

ii)State the expectedobservation andwrite an equation for the reaction involved. (2marks)

Answer

a) Theoxidationnumberof oxygen changes from–1 to –2 inH2O. (0.5)

Theoxidationnumberof oxygen changes from–1 to0 inO2. (0.5)

Hydrogenperoxide is simultaneously reduced andoxidized. (1)

b)i) H2O2(aq) + 2H+(aq) + 2e– 2H2O(l) (1)

ii)The colourof the reactionmixture changes fromcolourless tobrown. (1)

H2O2(aq) + 2I–(aq) + 2H+(aq) I2(aq) + 2H2O(l) (1)

c) i) H2O2(aq) + 2OH–(aq) O2(g) + 2H2O(l) + 2e– (1)

ii)The colourof the reactionmixture changes fromyellow-brown topale green. (1)

H2O2(aq) + 2OH–(aq) + 2Fe3+(aq) O2(g) + 2H2O(l) + 2Fe2+(aq) (1)

20.14 – 20.17

Summary

1 The following table summarizes the oxidizing properties of chlorine, dilute andconcentratednitric acid, and concentrated sulphuric acid.

Reaction Equation

Chlorine

aqueouschlorineoxidizesbromide ions tobromine Cl2(aq) + 2Br–(aq) 2Cl–(aq) +Br2(aq)

aqueouschlorineoxidizesiodide ions to iodine Cl2(aq) + 2I–(aq) 2Cl–(aq) + I2(aq)

reaction with cold anddilute sodium hydroxidesolution*

Cl2(g) + 2NaOH(aq)NaCl(aq) +NaOCl(aq) +H2O(l)

reaction with hot andconcentrated sodiumhydroxide solution*

3Cl2(g) + 6NaOH(aq) 5NaCl(aq) +NaClO3(aq) + 3H2O(l)

*Chlorine is oxidized and reduced simultaneously (disproportionation).

Reaction Equation

Concentrated/ dilutenitric

acid

concentrated nitric acidoxidizes metals; nitrateions are reduced tonitrogendioxide

Cu(s) + 2NO3–(aq) + 4H+(aq)

Cu2+(aq) + 2NO2(g) + 2H2O(l)

Mg(s) + 2NO3–(aq) + 4H+(aq)

Mg2+(aq) + 2NO2(g) + 2H2O(l)

hot concentrated nitricacid oxidizes non-metals;nitrateionsarereducedtonitrogendioxide

C(s) + 4HNO3(aq)CO2(g) + 4NO2(g) + 2H2O(l)

S(s) + 4HNO3(aq) SO2(g) + 4NO2(g) + 2H2O(l)

dilute nitric acid oxidizescopper; nitrate ions arereduced to ni t rogenmonoxide

3Cu(s) + 2NO3–(aq) + 8H+(aq)

3Cu2+(aq) + 2NO(g) + 4H2O(l)

Concentratedsulphuric

acid*

it oxidizes most metals;it is reduced to sulphurdioxide

Cu(s) + 2H2SO4(l)CuSO4(aq) + SO2(g) + 2H2O(l)

it oxidizes most non-metals

C(s) + 2H2SO4(l) CO2(g) + 2SO2(g) + 2H2O(l)

S(s) + 2H2SO4(l) 3SO2(g) + 2H2O(l)

it reacts with sodiumbromidetogivehydrogenbromide; it then oxidizeshydrogen bromide tobromine

NaBr(s) +H2SO4(l) NaHSO4(s) +HBr(g)

2HBr(g) +H2SO4(l) Br2(g) + SO2(g) + 2H2O(l)

it reacts with sodiumiodide to give hydrogeniodide; it then oxidizeshydrogen iod ide toiodine

NaI(s) +H2SO4(l) NaHSO4(s) +HI(g)

8HI(g) +H2SO4(l) 4I2(s) +H2S(g) + 4H2O(l)

*Concentrated sulphuric acid reactswith sodiumchloride to givehydrogen chloride.

NaCl(s) +H2SO4(l) NaHSO4(s) +HCl(g)

Notice that this isNOTa redox reaction.

2 The following table summarizes the reducingpropertyof aqueous sulphurdioxide.

Reaction Equation

SO2(aq) reduces permanganate ionstomanganese(II) ions

5SO32–(aq) + 2MnO4

–(aq) + 6H+(aq)

5SO42–(aq) + 2Mn2+(aq) + 3H2O(l)

SO2(aq) reduces dichromate ions tochromium(III) ions

3SO32–(aq) +Cr2O7

2–(aq) + 8H+(aq)

3SO42–(aq) + 2Cr3+(aq) + 4H2O(l)

SO2(aq) reducesaqueousbromine tobromide ions

SO32–(aq) +H2O(l) +Br2(aq)

SO42–(aq) + 2H+(aq) + 2Br–(aq)

SO2(aq) reduces iron(III) ions toiron(II) ions

SO32–(aq) +H2O(l) + 2Fe3+(aq)

SO42–(aq) + 2H+(aq) + 2Fe2+(aq)

Topic 520 Redox Reactions, Chemical Cells and Electrolysis 21Unit 20 Oxidation and reduction

ExamtipsExamtipsExamtipsExamtips ♦ Questions often ask about the reaction between chlorine and sodiumhydroxide solution.

♦ Chlorine bleach is made by reacting chlorine with sodium hydroxidesolution.Theactiveingredientissodiumhypochlorite,NaOCl.Householdbleach shouldNOTbeused togetherwithacidsbecauseOCl–(aq) ionsreact withH+(aq) ions to giveCl2(g),which is toxic.

Cl–(aq)+OCl–(aq)+2H+(aq) Cl2(g) +H2O(l)

♦ Questions often ask about aqueous bromine.

– Aqueous bromine is a common oxidizing agent.

– It changes from yellow-brown to colourless when reduced.

Br2(aq)+2e– 2Br–(aq)

– The oxidizing power of the halogens is in the order of Cl2 > Br2 >I2.

– There is NO reaction between aqueous bromine and potassiumchloride solution.

♦ Questions often ask about the similarity in chemical properties ofhalogens.

e.g.

– Both Cl2(aq) and Br2(aq) can oxidize SO32–(aq) ions to SO4

2–(aq)ions.

X2(aq)+ SO32–(aq)+H2O(l) 2X–(aq)+ SO4

2–(aq)+2H+(aq)

– BothCl2(aq)andBr2(aq)can reactwithFe2+(aq) ions togiveFe3+(aq)ions.

X2(aq)+2Fe2+(aq) 2X–(aq)+2Fe3+(aq)

– Both Cl2(aq) andBr2(aq) can reactwith I–(aq) ions to give I2(aq).

X2(aq)+2I–(aq) 2X–(aq)+ I2(aq)

– Both Cl2(aq) andBr2(aq) canundergo disproportionation in alkalis.

X2(aq)+2OH–(aq) X–(aq)+OX–(aq)+H2O(l)

♦ Iodideionhasthehighestreducingpoweramongchlorideion,bromideion and iodide ion. Thus,NaI(aq) can reactwith Br2(aq).

♦ Questions often ask students to compare the properties of differentacids.

♦ Dilute nitric acid can react with copper while dilute ethanoic acidCANNOT. This is because dilute nitric acid acts as anoxidizing agent.

♦ Students should be able to explain the different behaviours of NaCl(s)and NaI(s) towards concentrated H2SO4(l) in terms of the differentreducing power ofHCl(g) and HI(g).

HI(g) is a stronger reducing agent than HCl(g).

Acid

ObservationProperty

Dilute HNO3(aq) Dilute H2SO4(aq)

Reactionwith alkali salt and water are produced

Reactionwith carbonate(or hydrogencarbonate)

carbon dioxide gas is given off

Actionon litmus solution give a red colour

Reactionwith zinc abrown gas is given off a colourless gas is given off

Reactionwith bariumchloride solution

no white precipitate forms awhite precipitate forms

TitrationwithNaOH(aq)more NaOH(aq) is needed to reach the end point forH2SO4(aq) than HNO3(aq)

Acid

ObservationProperty

Concentrated H2SO4(l) Dilute H2SO4(aq)

Reactionwith alkali salt and water are produced

Reactionwith carbonate(or hydrogencarbonate)

carbon dioxide gas is given off

Actionon litmus solution give a red colour

Oxidizingproperty canoxidize copper no reaction with copper

Non-volatility

non-volatile, candisplaceother volatile acids (e.g.HCl(g) and HNO3(g)) fromtheir salts

not non-volatile

Topic 522 Redox Reactions, Chemical Cells and Electrolysis 23Unit 20 Oxidation and reduction

➤Questions often ask whether there will be a colour change when sulphurdioxide is passed into a certain solution.

e.g.

Acolourchangewilloccurwhensulphurdioxideispassedintothefollowingsolutions:

– Br2(aq) (from yellow-brown to colourless);

– I2(aq) (frombrown to colourless);

– Fe3+(aq) (from yellow-brown topale green);

– acidified KMnO4(aq) (frompurple to colourless);

– acidified K2Cr2O7(aq) (fromorange to green).

These colour changes are due to the reducing action of aqueous sulphurdioxide.

➤Sulphur dioxide does NOT reactwith FeSO4(aq).

➤Solutionsof sodium sulphite and sodium sulphate canbedistinguishedbyusing

– acidified potassium permanganate solution;

– acidified potassium dichromate solution; or

– aqueous bromine.

RemarksRemarks*

b)Anyoneof the following:

• Copperdissolves. (1)

• Ablue solution results. (1)

• A colourless gas evolves. (1)

• A choking smell is detected. (1)

Cu(s) + 2H2SO4(l) CuSO4(aq) + SO2(g) + 2H2O(l) (1)

c) Foodpreservation (1)

♦ QuestionsoftenaskaboutmethodsfordistinguishingbetweenFe2+(aq)ions and Fe3+(aq) ions.

– Colour

Fe2+(aq) ions are pale green in colour

while Fe3+(aq) ions are yellow-brown in colour.

– Acidifiedpotassium permanganate solution

Fe2+(aq)ionscandecolorizeacidifiedpotassiumpermanganatesolutionbut Fe3+(aq) ions cannot.

MnO4–(aq)+5Fe2+(aq)+8H+(aq)

Mn2+(aq)+5Fe3+(aq)+4H2O(l)

– Dilute sodiumhydroxide solution / dilute aqueous ammonia

Fe2+(aq) ions give a green precipitate with dilute sodium hydroxidesolution / dilute aqueous ammonia.

Fe3+(aq) ions give a reddish brown precipitate with dilute sodiumhydroxide solution / dilute aqueous ammonia.

Fe2+(aq)+2OH–(aq) Fe(OH)2(s)

Fe3+(aq)+3OH–(aq) Fe(OH)3(s)

– Concentrated nitric acid

Fe2+(aq) ions give a brown gas with concentrated nitric acid butFe3+(aq) ions donot.

Example

Sulphurdioxidecanbeobtainedfromthereactionbetweenhotconcentratedsulphuricacid and copper turnings.

a) Draw a labelled diagram to show the set-up in preparing and collecting sulphurdioxide. (2marks)

b)State theobservation andwrite a chemical equation for the reaction thatoccurs. (2marks)

c) Somepeoplewoulduse sulphurdioxide to treat food.

Suggest apurposeof treating foodwith sulphurdioxide. (1mark)

Answer

a)

(2)heat

copper turnings

concentratedsulphuric acid

delivery tube

gas jar

Topic 52� Redox Reactions, Chemical Cells and Electrolysis 2�Unit 21 Oxidation and reduction in chemical cells

21.1 Oxidation and reduction in a simple chemical cell

21.2 Redox reactions in a zinc-carbon cell

21.3 Redox reactions in simple chemical cells with inertelectrodes

21.4 The chemistryof a lead-acid accumulator

21.5 Fuel cells

21.6 The chemistryof lithium ion secondary cells

21.7 Redox reactions in everyday life

Oxidation and reduction in chemical cellsUnit 21 21.1 – 21.7

Summary

1 Thefollowingtablesummarizes reactionsthatoccurat theelectrodesofachemicalcell.

Negative electrode (anode) Positive electrode (cathode)

Electrodewhere oxidationoccurs reductionoccurs

Directionof electron flow fromthenegativeelectrodetothepositiveelectrodeintheexternalcircuit

2 In a zinc-carbon cell:

a) The carbon rod is thepositive electrode.

The zinc case is thenegative electrode.

The electrolyte is amoist pasteof ammonia chloride.

b)At thenegative electrode:


At thepositive electrode:

2NH4+(aq) + 2e– 2NH3(aq) +H2(g)

c) Manganese(IV)oxide,anoxidizingagent,isusedtoremovethehydrogenformedat thepositive electrode.

3 In a lead-acid accumulator:

a) At thenegative electrode:

Pb(s) + SO42–(aq)

discharge

charge PbSO4(s) + 2e–

b)At thepositive electrode:

PbO2(s) + 4H+(aq) + SO42–(aq) + 2e–

discharge

charge PbSO4(s) + 2H2O(l)

4 In a hydrogen-oxygen alkaline fuel cell, hydrogen and oxygen react to generateelectricity,heat andwater.With thenon-stop supplyof fuel (hydrogen), a fuel cellcan run forever andproduce a steady supplyof electrical energy.

a) At thenegative electrode:

H2(g) + 2OH–(aq) 2H2O(l) + 2e–

b)At thepositive electrode:

O2(g) + 2H2O(l) + 4e– 4OH–(aq)

5 Directmethanolfuelcellsusemethanol,insteadofpurehydrogen,asthefuel.Theseappear to be the most promising alternative power source for portable electronicdevices.


ExamtipsExamtipsExamtipsExamtips ♦ Questionsmayaskaboutreactionsthatoccurinanunfamiliarchemicalcell.

chromium rod


salt bridge

iron rod

chromium(III)chloride solution

iron(II) sulphate solution

Astimegoesby,thecolourof iron(II)sulphatesolutiongraduallyfadesout.

At the iron rod: Fe2+(aq)+2e– Fe(s)

At the chromium rod:Cr(s) Cr3+(aq)+3e–

Thus, electrons flow from the chromium rod to the iron rod in theexternal circuit.

♦ Questionsmayaskaboutchangeinthecolourintensityoftheelectrolyteofachemicalcell, i.e.change in theconcentrationofcoloured ions inthe electrolyte.

silver rod(+)

copper rod(–)

Cu2+

Cu2+

electron flow

copper(ll) sulphatesolution

Copper rod is the anodewhereoxidation occurs.

Cu(s) Cu2+(aq)+2e–

Silver rod is the cathodewhere reduction occurs.

Cu2+(aq) +2e– Cu(s)

The net effect is the transfer of copper from the copper rod to thesilverrod.Theconcentrationofcopper(ll)ionsintheelectrolyteremainsthe same. Theblue colour of the solution doesNOT change.

♦ Questionsmayaskstudentstopredictthechemicalchangesthatwouldoccur in a chemical cell based on the relative oxidizing power of thereagents involved.

e.g.

Consider the chemical cell shown below.

carbon rod X

digital multimeter

salt bridge

carbon rod Y

mixture of FeSO4(aq) and

Fe2(SO4)3(aq)

mixture of KBr(aq)and Br2(aq)

The question gives the information that Br2(aq) is a stronger oxidizingagent than Fe3+(aq) ion.

oxidizing powerincreasing

oxidizing powerincreasing

Br2(aq) + 2e– 2Br–(aq)

Fe3+(aq) + e– Fe2+(aq)

oxidize

higher in electrochemical

series

reduced form of half-equation

oxidized form of half-equation

lower in electrochemical

series

Hence it can be deduced that Br2(aq) acts as the oxidizing agent andFe2+(aq)ionactsasthereducingagent.Thefollowingchemicalchangeswould occur:

Fe2+(aq) Fe3+(aq)+ e–

Br2(aq)+2e– 2Br–(aq)

♦ Questions often ask about the functions of various components of azinc-carbon cell.


➤Questions often ask about fuel cells.RemarksRemarks*

Example

The followingdiagram represents an alkaline fuel cell.

electrode Y electrode X

O2(g)

O2(g)

H2(g)

H2(g) + H2O(g)

concentratedNaOH(aq)

load

Hydrogen andoxygen arepassed into the fuel cell.

a) Write ionic half-equations for the chemical changes occurring at electrodes X andY respectively. (2marks)

b)Identity thedirectionof electron flow in the external circuit. (1mark)

c) Suggest TWO advantages of using fuel cells over using conventional chemicalcells. (2marks)

d)SuggestwhyfuelcellsareNOTwidelyusedindailylifedespitetheirhavinganumberof advantages. (1mark)

Answer

a) At electrodeX: H2(g) + 2OH–(aq) 2H2O(l) + 2e– (1)

At electrodeY: O2(g) + 2H2O(l) + 4e– 4OH–(aq) (1)

b)FromelectrodeX to electrodeY (1)

c) Any twoof the following:

• Fuel cells donot emit air pollutants. (1)

• Fuel cellshavehigh fuel efficiency. (1)

• Fuel cells can operate continuously if the flow of hydrogen and oxygen can bemaintained (theydonot rundownor require charging). (1)

d)Anyoneof the following:

• Currently theH2(g)used in fuel cells isderived from fossil fuels.Theuseof fuelcells cannothelp save fossil fuels. (1)

• Hydrogen is explosive.Mishandlingof the gas canbedisastrous. (1)

♦ When using a zinc-carbon cell in electrolysis, the carbon electrode(positive electrode) is connected to the anode of the electrolytic cell.Electrons in the external circuit flow to the carbon electrode of thezinc-carbon cell.

electronflow

moist paste of ammonium chloride

manganese(IV) oxide and carbon powder

carbon rod

zinc case

cathode of electrolytic cell

anode of electrolytic cell

♦ The state of charge of a lead-acid accumulator can be estimated bymeasuring thedensity of thebattery acid.

Theoverall reaction for thedischargingprocessof a lead-acid cell canbe representedby the following equation:

Pb(s) + PbO2(s) + 4H+(aq)+2SO42–(aq) 2PbSO4(s) + 2H2O(l)

During the discharging process, sulphuric acid is consumed. Thus, thedensity of thebattery acid decreases.

During the charging process, sulphuric acid is produced. Thus, thedensity of thebattery acid increases.

♦ A type of breathalyzer consists of a chemical cell. If the breath of adrivercontainsethanol,theethanolwouldbeconvertedtoethanoicacidat the anode (where oxidation occurs) of the cell. An electric currentwouldbeproduced.

CH3CH2OH+H2O CH3COOH+4H++4e–

Ahigherconcentrationofethanolwouldproducealargercurrent.Thus,the amount of ethanol in thebreathof thedriver canbe estimated.

Topic 530 Redox Reactions, Chemical Cells and Electrolysis 31Unit 22 Electrolysis

22.1 Electrolysis: chemical reactions fromelectricity

22.2 Comparing a chemical cell and an electrolytic cell

22.3 Electrolysis of molten sodium chloride using carbonelectrodes

22.4 Some knowledge related to aqueous electrolytes

22.5 Electrolysis of aqueous solutionsof ionic compounds

22.6 Factors affecting the order of discharge of ions during theelectrolysis of aqueous solutions

22.7 Thepositionofionsintheelectrochemicalseriesandtheorderofdischargeof ions

22.8 The effect of concentration of ions in the solution and theorderofdischargeof ions

22.9 The natureof electrodes and theorderofdischargeof ions

22.10 Industrial usesof electrolysis

22.11 Environmental impactof the electroplating industry

ElectrolysisUnit 22 22.1 – 22.9

Summary

1 Electrolysis is the chemical reaction that occurs when electricity passes through anelectrolyte inmolten stateor in aqueous solution.

–

+

+

+

+

–

–

–

electron flow

anode(anion

dischargedhere)

cathode(cationdischargedhere)

key:cation

anion

–+

electrolyte in molten state or in aqueous solution

2 Anelectrolyticcellisacellinwhichelectrolysisoccurs.Thefollowingtablesummarizessome facts concerning the electrodesof an electrolytic cell.

Electrode Connected to which electrode of d.c. supply

Ion discharged here

Process that occurs here

Anode positive anion oxidation

Cathode negative cation reduction

3 Factors affecting the order of discharge of ions during the electrolysis of aqueoussolutions include:

a) thepositionof ions in the electrochemical series;

b)the concentrationof ions in the solution; and

c) thenatureof the electrodesused.

4 The following table summarizes theproductsof electrolysisof somemolten liquidsand aqueous solutions.

SubstanceMaterial of Product at Change in the

solutionanode cathode anode cathode

Moltenliquid sodiumchloride carbon carbon chlorine

gas sodium —

Topic 532 Redox Reactions, Chemical Cells and Electrolysis 33Unit 22 Electrolysis

SubstanceMaterial of Product at Change in the

solutionanode cathode anode cathode

Aqueoussolution

dilute sulphuricacid

dilutehydrochloricacid

platinum

carbon

platinum

carbon

oxygengas

chlorinegas

hydrogengas

hydrogengas

becomesmoreconcentrated

becomesmoredilute

verydilutesodiumchloride

dilute sodiumchloride

carbon

carbon

carbon

carbon

oxygengas

chlorinegas

hydrogengas

hydrogengas


becomes sodiumhydroxide solution

dilute sodiumsulphate

dilute sodiumnitrate

carbon

carbon

carbon

carbon

oxygengas

oxygengas

hydrogengas

hydrogengas



concentratedsodiumchloride

concentratedsodiumchloride

carbon

carbon

carbon

mercury

chlorinegas

chlorinegas

hydrogengas

sodium

becomes sodiumhydroxide solution

becomesmoredilute

dilute copper(II)sulphate




carbon

carbon

copper

copper

carbon

copper

copper

carbon

oxygengas

oxygengas

copper(II)ions

copper(II)ions

copper

copper

copper

copper

becomes sulphuricacid

becomes sulphuricacid

remains the same

remains the same

ExamtipsExamtipsExamtipsExamtips ♦ Redox reaction must be involved in aworking electrolytic cell.

♦ Questions may ask about the electrolysis of dilute copper(ll) sulphatesolutionusing carbon anode and copper cathode.

copperelectrode(cathode)copperdepositoxygen

gas

carbonelectrode

(anode)

dilute copper(II)sulphate solution

At the anode:

4OH–(aq) O2(g) + 2H2O(l) + 4e–

At the cathode:

Cu2+(aq)+2e– Cu(s)

Thebluecolourofthesolutionfadesgraduallybecausetheconcentrationof copper(II) ions in the electrolyte decreases.

Copper(II) ions and hydroxide ions are consumed in the electrolysis.Hydrogen ions and sulphate ions remain in the solution. Thus, thesolution eventually becomes sulphuric acid.

Example

Averydilute sodiumchloride solution iselectrolyzedusing inertelectrodes fora longperiodof time.

a) Statetheexpectedobservationatthecathode.Explainyouranswerwithanappropriateionichalf-equation. (3marks)

b)StateALLexpectedobservationsattheanode.Explainyouranswerwithappropriateionichalf-equations. (6marks)

c) Explainwhether the resulting solution is acidic, alkalineorneutral. (1mark)

Answer

a) Colourless gas bubbles evolve. (1)

Ahydrogenionisastrongeroxidizingagentthanasodiumion.Hydrogenionsarepreferentiallydischarged (reduced) to formhydrogengas. (1)

2H+(aq) + 2e– H2(g) (1)

♦ Questionsmayalsoaskabouttheelectrolysisofdilutecopper(II)sulphatesolution using copper anode and carbon cathode.

copper anode carbon cathode

dilute copper(II)sulphatesolution

At the anode:

The copper anode dissolves to form copper(II) ions.

At the cathode:

Copper(II) ions are discharged to form a deposit of copper on thecathode.

Changes in the solution:

Theneteffectisthetransferofcopperfromtheanodetothecathode.The concentration of copper(II) ions in the electrolyte remains thesame. Theblue colour of the solution does NOT change.

♦ Questionsoften list somesolutionsandaskwhichof themwouldgivehydrogen upon electrolysis using carbon electrodes.

e.g.

The electrolysis of concentrated potassium sulphate solution andconcentrated calcium chloride solution gives hydrogen at the carboncathode.

Topic 53� Redox Reactions, Chemical Cells and Electrolysis 3�Unit 22 Electrolysis

b)Colourless gas bubbles evolve at thebeginning. (1)

A hydroxide ion is a stronger reducing agent than a chloride ion. Hydroxide ionsarepreferentiallydischarged (oxidized) to formoxygengas. (1)

4OH–(aq) O2(g) + 2H2O(l) + 4e– (1)

The sodium chloride solution becomes more concentrated as water is decomposedin the electrolysis. When the concentration of chloride ions is much higher thanthatofhydroxideions,chlorideionsarepreferentiallydischarged(oxidized)toformchlorine gas. (1)

2Cl–(aq) Cl2(g) + 2e– (1)

A greenishyellowgas evolves. (1)

c) The resulting solution is alkaline. After electrolysis, the concentration of OH–(aq)ions ishigher than thatofH+(aq) ions. (1)

➤Questionsoftenaskaboutthemembranecellforthemanufactureofsodiumhydroxide from concentrated sodium chloride solution.

Cl–(aq)OH–(aq)

Na+(aq)

chlorine hydrogen

brine

steel electrode

sodium hydroxide solution

titaniumelectrode

used brine

water

membrane

At the anode: 2Cl–(aq) Cl2(g) + 2e–

At the cathode: 2H+(aq)+2e– H2(g)

The membrane allows sodium ions to pass through. Water dissociatescontinuouslytoreplacethehydrogenionsdischargedatthecathode.Thus,there is an excess of hydroxide ions in the cathodic compartment. Thus,sodium hydroxide solution is obtained.

(This cellwill be further discussed in Topic 14 Industrial Chemistry.)

➤Studentsshouldbeabletodrawadiagramoftheset-upfortheelectrolysisof sodium chloride solution.

chlorine

carbonelectrode

(anode)

carbonelectrode(cathode)

dilute or concentratedsodium chloride solution

hydrogen

➤Questionsoftenaskabouttheelectrolysisofverydilute/dilute/concentratedsodium chloride solution.

➤Noticethathydrogenions/chlorideionsarepreferentiallydischarged,NOThydrogen / chlorine.

➤In the electrolysis of concentrated sodium chloride solution using carbonelectrodes,gaseousproducts (hydrogenand chlorine) are liberatedatbothelectrodes.

Theoverall cell reaction is

2H+(aq)+2Cl–(aq) H2(g) +Cl2(g)

Thetheoreticalvolumeratioofthegascollectedatthecathodetothegascollected at the anode is 1:1.

RemarksRemarks*

22.10 – 22.11

Summary

1 The following table summarizes two industrial usesof electrolysis.

Use Anode Cathode Electrolyte Process

Refiningofcopper

impurecopper

pure copper copper(II)sulphatesolution+sulphuricacid

at the anode


Fe(s) Fe2+(aq) + 2e–

Cu(s) Cu2+(aq) + 2e–

at the cathode


overall cell reaction

Cu(s) Cu(s)(anode) (cathode)

Electroplating platingmetal object tobeplated

aqueoussolutionofa salt of theplatingmetal

example—plating anobjectwith copper

at the anode

Cu(s) Cu2+(aq) + 2e–

at the cathode


overall cell reaction

Cu(s) Cu(s)(anode) (cathode)

Topic 53� Redox Reactions, Chemical Cells and Electrolysis 3�Unit 22 Electrolysis

2 Methods to controlpollution from the electroplating industry include:

a) reducing thevolumeofwaste solutions;

b)recyclingwaste solutions; and

c) treating effluents beforedischarge.

ExamtipsExamtipsExamtipsExamtips ♦ Questions often ask about the set-up for electroplating nickel on anobject.

Anode Cathode Electrolyte

nickel object to beplated nickel(II) sulphate solution

nickelanode

object to be plated (cathode)

nickel(ll) sulphatesolution

♦ After chromium plating, it is a common practice to convert anychromium(VI) compounds left in the bath to chromium(III) compoundsbefore discharge. This is because chromium(VI) compounds are moretoxic than chromium(III) compounds.

c) Insoluble impurities deposit under the impure copper as ‘sludge’. According to theinformationgiven, suggestwhatsubstances thesludgewouldcontain.Explainyouranswer. (2marks)

d)Explainwhy theblue colourof the copper(II) sulphate solution fades gradually. (3marks)

Answer

a) The impure copper (1)

b)The impure copper givesup electrons to formcopper(II) ions. (1)

Copper(II) ions aredischarged to formcopperon thepure copper. (1)

c) Gold and silver (1)

Comparedwith copper, thesemetals form ions less readily. (1)

d)The concentrationof copper(II) ions in the solutiondrops gradually. (1)

At the anode, zinc dissolves as ions because it forms ions more readily thancopper. (1)

At the cathode, copper(II) ions are alwayspreferentiallydischarged. (1)

➤At the impure copper anode:Zn(s) Zn2+(aq)+2e–

Cu(s) Cu2+(aq)+2e–

At thepure copper cathode: Cu2+(aq)+2e– Cu(s)

RemarksRemarks*

Example

The copper extracted from a copper ore contains impurities including metals such asgold, silverandzinc.The impurecopper ispurifiedbyelectrolysisas illustrated in thediagrambelow.

impure copper

sludge

pure copper

copper(II) sulphate solution

d.c. supply

a) What is the anode, the impure copperor thepure copper? (1mark)

b)Explain briefly how impure copper can be purified by electrolysis as illustrated inthediagramabove. (2marks)

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