redox and electrochemistry.pdf

8/10/2019 Redox and Electrochemistry.pdf

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OxidationReduction

andElectrochemistry

DavidA.KatzDepartmentofChemistry

PimaCommunityCollege


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OxidationReductionReactions

Inanoxidationreduction(Redox)reaction,

electronsaretransferredfromonespeciestoanother.

Forexample,inasinglereplacementreaction

Cu(s) +2AgNO3(aq) 2Ag(s)+Cu(NO3)2(aq)

TheCuatomsloseelectronstoformCu2+intheCu(NO3)2(aq)andtheAg

+ gainselectronstoformmetallicAg


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Thiscanbemoreeasilyobservedbywritingthenet

ionicequationforthereaction:Cu(s) +2Ag

+(aq) 2Ag(s)+Cu

2+(aq)

ThemetallicCuatomsareuncombined,sotheyare

consideredtohaveanoxidationnumberofzero.

TheinitialcombinedAg+ ionsareina+1oxidationstate.

EachCuatomwilllose2electronsto2Ag+ ions TheresultingAgatomsareconsideredtohavean

oxidationnumberofzero


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Cu(s) +2Ag+

(aq) 2Ag(s)+Cu2+

(aq)


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Oxidation

Numbers

Inordertokeeptrack

ofwhatloses

electronsandwhat

gainsthem,welist

the

oxidation

numbersofeach

element.


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OxidationandReduction

Aspeciesisoxidizedwhenitloseselectrons.

Here,zinclosestwoelectronstogofromneutralzinc

metaltotheZn2+ ion.


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Oxidation

and

Reduction

Aspeciesisreducedwhenitgainselectrons.

Here,eachoftheH+ gainsanelectronandthey

combinetoformH2.


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Oxidation

and

Reduction

Thespeciesthatcontainstheelementthatisreducedistheoxidizingagent.

H+

oxidizes

Zn

by

taking

electrons

from

it. Thespeciesthatcontainstheelementthatis

oxidizedisthereducingagent. ZnreducesH+ bygivingitelectrons.


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Assigning

Oxidation

Numbers

1. Elementsintheirelementalformhavean

oxidationnumberof0.

2. Theoxidationnumberofamonatomicion

isthesameasitscharge.


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Assigning

Oxidation

Numbers

3. Theoxidationnumberofmetalsdepends

ontheirpositionintheperiodictable GroupIAelementsare+1

GroupIIAelementsare+2

GroupIIIAelementsare+3

GroupIVAmetalsareusually+2or+4

Group

VA

metals

are

usually

+3

or

+5


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Assigning

Oxidation

Numbers4. Nonmetalstendtohavenegative

oxidationnumbers,althoughsomearepositiveincertaincompoundsorions.

Oxygenalwayshasanoxidationnumberof

2,

except

in

the

peroxide

ion

in

which

it

has

anoxidationnumberof 1.

Hydrogenisalways 1whenbondedtoametal

Hydrogenis +1whenbondedtoanonmetal.


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Assigning

Oxidation

Numbers4. Nonmetals(continued).

Fluorinealwayshasanoxidationnumberof1.

Thehalogens(Cl,Br,andI)haveanoxidation

number

of 1

when

they

are

negative Thehalogens(Cl,Br,andI)willhave

positiveoxidationnumbersinoxyanions(ClO,ClO

2

,ClO3

,etc.)


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Assigning

Oxidation

Numbers

5. Thesumoftheoxidationnumbersina

neutralcompoundis0.6. Thesumoftheoxidationnumbersina

polyatomicionisthechargeontheion.


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BalancingOxidationReductionEquations

Oxidationreductionequationsmaybedifficultto

balance.

Generally,theeasiestwaytobalancetheequation

ofanoxidationreductionreactionisviathehalf

reactionmethod.

Thisinvolvestreatingtheoxidationandreduction

astwoseparateprocesses,balancingthesehalf

reactions,andthencombiningthemtoattainthe

balancedequationfortheoverallreaction.


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BalancingRedox Equationsbythe

HalfReactionMethod

1. Assignoxidationnumberstotheelements

intheequation

2. Determinewhatisoxidizedandwhatis

reduced.

3. Writetheoxidationandreductionhalf

reactions.


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HalfReactionMethod

4.

Balance

each

half

reaction.a. BalanceelementsotherthanHandO.

b. BalanceObyaddingH2O.

c. BalanceHbyaddingH+.d. Balancechargebyaddingelectrons.

5. Multiplythehalfreactionsbyintegersso

thattheelectronsgainedandlostarethesame.


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HalfReactionMethod

6.

Add

the

half

reactions,

subtracting

things

thatappearonbothsides.

7. Makesuretheequationisbalanced

accordingtomass.8. Makesuretheequationisbalanced

accordingtocharge.


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HalfReactionMethod

ConsiderthereactionbetweenMnO4 andC2O4

2 :

MnO4

(aq) +C2O42

(aq) Mn2+(aq) +CO2(aq) (acidicsolution)


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HalfReactionMethod

First,

assign

oxidation

numbers

(remember

oxygen

is

2)

MnO4 + C2O4

2Mn2+ +CO2

+7 +3 +4+2

Theoxidationnumberofmanganesegoesfrom+7to+2,itis

reduced.

Theoxidationnumberofcarbongoesfrom+3to+4,itis

oxidized.


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HalfReactionMethod

MnO4 + C2O42

Mn2+ +CO2

+7 +3 +4+2

TheMnO4 istheoxidizingagent

TheC2O42 isthereducingagent

Determinetheoxidizingandreducingagents


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HalfReactionMethod

Theoxidationhalfreactionis

C2O42 CO2

Balance

the

half

reactionC2O4

2 2CO2

Thisbalancesboththecarbonatomsand

theoxygenatoms


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HalfReactionMethod

Balancethechargebyadding2electronsto

therightside.

C2O42

2

CO2+

2

e


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HalfReactionMethod

Thereductionhalfreactionis

MnO4

Mn2+

The

manganese

is

balancedInordertobalancethe4oxygens,weadd4watermoleculestotherightside.

MnO4

Mn2+ +4H2O


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HalfReactionMethod

MnO4

Mn2+ +4H2O

Theadditionofwaterontherightsideofthe

equationincludedhydrogenatoms.

To

balance

the

8

hydrogens,

add

8

H+

to

the

leftside.

8H+ +MnO4

Mn2+ +4H2O


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HalfReactionMethod

8

H+

+

MnO4

Mn2+

+

4

H2O

Balancethecharges,add5e totheleftside.

5e +8H+ +MnO4

Mn2+ +4H2O


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HalfReactionMethodBeforethetwohalfreactionscanbeadded

together,

the

number

of

electrons

lost

must

be

equaltotheelectronsgained:

C2O42 2CO2+2e

(2e lost)

5

e

+

8

H+

+

MnO4

Mn2+

+

4

H2O

(5egained)

Toattainthesamenumberofelectronsoneach

side,multiplythefirstreactionby5andthe

second

by

2.5C2O4

2 10CO2+10e

10e +16H+ +2MnO4

2Mn2+ +8H2O


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HalfReactionMethod

5C2O42 10CO2+10e

10e +16H+ +2MnO4

2Mn2+ +8H2O

Addthesehalfreactions

10e

+16H+ +2MnO4

+5C2O42

2Mn2+ +8H2O+10CO2+10e

Cancelouttheelectronsfrombothsidestoget

16H+ +2MnO4 +5C2O4

2 2Mn2+ +8H2O+10CO2


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BalancingRedox Equations

inBasicSolution

If

a

reaction

occurs

in

basic

solution,

use

OH

andH2OinsteadoftheH+ andH2Ousedinacid

solution

Theeasiestmethodistobalancetheequation

asifitoccurredinacid. Oncetheequationisbalanced,addOH toeach

sidetoneutralizetheH+ intheequationand

create

water

in

its

place. Ifthisresultsinwateronbothsidesoftheequation,subtractwaterfromeachsideasafinalstep.


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inBasicSolution

Completeandbalancethefollowingredoxequation

that

takes

place

in

basic

solutionCN (aq) + MnO4

(aq) CNO

(aq) + MnO2(s) (basicsolution)

First,assignoxidationnumbers(rememberoxygenis 2)

+2 3 +7 2 +4 3 2 +4 2

CN(aq)

+ MnO

4

(aq)

CNO(aq)

+ MnO2

(s)

Note: InananionsuchasCN,C,whichcomesfirst,wouldbepositiveandN

wouldbenegative. Ifnegative,Nwouldbe 3.


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inBasicSolution

+2 3 +7 2 +4 3 2 +4 2

CN (aq) + MnO4

(aq) CNO

(aq) + MnO2(s)

The

oxidation

number

of

carbon

goes

from

+2

to

+4,

it

is

oxidized. (Clost2e)

Theoxidationnumberofmanganesegoesfrom+7to+4,itis

reduced. (Mngained3e)

MnO4 istheoxidizingagent

CN isthereducingagent


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inBasicSolution

Thereductionhalfreactionis

MnO4

MnO2

Balancethehalfreaction

MnO4

MnO2

Balance

the

oxygen

atoms,

add

H+

and

H2O4 H+ + MnO4

MnO2+2H2O


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inBasicSolution

Balancethecharges,add3e totheleftside

3e + 4 H+ + MnO4

MnO2+2H2O


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inBasicSolution

Theoxidationhalfreactionis

CN CNO

SincetheCatomsarebalanced,balancethe

oxygen.

AddH2OtotheleftsideandH+ totherightside.

H2O

+

CN

CNO

+

2

H+


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inBasicSolution

H2O + CN

CNO

+2H+

Balancethecharge,add2e totherightside

H2O + CN

CNO

+2H+

+ 2e


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inBasicSolutionBeforethetwohalfreactionscanbeadded

together,

the

number

of

electrons

lost

must

be

equaltotheelectronsgained:

H2O + CN

CNO +2H+ + 2e (2e lost)

3

e

+

4

H+

+

MnO4

MnO2+

2

H2O

(3

e

gained)Toattainthesamenumberofelectronsoneach

side,multiplythefirstequationby3andthe

second

by

2.3H2O + 3CN

3CNO +6H+ + 6e

6e +8H+ +2MnO4

2MnO2+4H2O


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inBasicSolution

Addthehalfreactions

3H2O + 3CN +8H+ +2MnO4

3CNO +6H+ +2MnO2+4H2O

CancelouttheH+ andH2Otoget

2 H+ +3CN +2MnO4

3 CNO + 2 MnO2+ H2O


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inBasicSolution

Thisreaction,however,takesplaceinbasicsolution

2 H+ +3CN +2MnO4

3 CNO + 2 MnO2+ H2O

Add2OH tobothsidestocancelouttheH+

2OH

+2 H+ +3CN

+2MnO4

3 CNO

+ 2 MnO2+ H2O +2OH

(Rememberthat H+ + OH H2O)

2H2O+3CN +2MnO4

3 CNO +2 MnO2+H2O+2OH

Thebalancedequationis

3CN +2MnO4 + H2O 3 CNO

+2 MnO2+2OH


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AHistoryofElectricity/Electrochemistry

Thales of Miletus (640546 B.C.) is

credited with the discovery that

amber when rubbed with cloth or

fur acquired the property of

attracting light objects.

The word electricity comes from

"elektron" the Greek word for

amber.

Otto von Guericke (16021686)invented the first electrostatic

generator in 1675. It was made of

a sulphur ball which rotated in a

wooden cradle. The ball itself wasrubbed by hand and the charged

sulphur ball had to be

transported to the place where

the electric experiment was

carried out.

ThalesofMiletus OttovonGuericke


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Eventually,aglassglobereplacedthe

sulfursphereusedbyGuericke

Later,largediskswereused


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EwaldJrgenvonKleist(17001748),

inventedtheLeydenJarin1745to

storeelectricenergy. TheLeydenJar

containedwaterormercuryandwas

placedontoametalsurfacewith

groundconnection.

In1746,theLeydenjarwas

independentlyinventedby physicist

PietervanMusschenbroek (1692

1761)and/orhislawyerfriend

AndreasCunnaeusinLeyden/the

Netherlands

Leydenjarscouldbejoinedtogether

tostorelargeelectricalcharges


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In1752,BenjaminFranklin(17061790)demonstratedthatlightningwaselectricityinhisfamouskiteexperiment

In1780,ItalianphysicianandphysicistLuigiAloisio Galvani(17371798)discoveredthatmuscleandnervecells

produceelectricity.Whilstdissectingafrogonatablewherehehadbeenconductingexperimentswithstaticelectricity,Galvanitouchedtheexposed

sciatic

nerve

with

his

scalpel,

which

had

pickedupanelectriccharge.Henoticedthatthefrogslegjumped.


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CountAlessandroGiuseppeAntonioAnastasio

Volta

(1745

1827)

developed

the

first

electric

cell,calledaVoltaicPile,in1800.

Avoltaicpileconsistofalternatinglayersoftwo

dissimilarmetals,separatedbypiecesof

cardboard

soaked

in

a

sodium

chloride

solution

orsulfuricacid.

Voltadeterminedthat

thebestcombinationof

metalswaszincand

silver

Voltaselectricpile(right)

AVoltaicpileatthe

SmithsonianInstitution,(far

right)


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In1800,EnglishchemistWilliamNicholson

(17531815)andsurgeonAnthonyCarlisle

(17681840)separatedwaterintohydrogen

andoxygenbyelectrolysis.

JohannWilhelmRitter(17761810)repeated

Nicholsonsseparationofwaterintohydrogen

andoxygenbyelectrolysis. Soonthereafter,

Ritterdiscoveredtheprocessof

electroplating. Healsoobservedthatthe

amountofmetaldepositedandtheamount

ofoxygenproducedduringanelectrolytic

processdependedonthedistancebetween

theelectrodes

HumphreyDavy(17781829)utilizedthe

voltaic

pile,

in

1807,

to

isolate

elemental

potassiumbyelectrolysiswhichwassoon

followedbysodium,barium,calcium,

strontium,magnesium.

WilliamNicholson

JohannWilhelm

Ritter

HumphreyDavy


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MichaelFaraday(17911867)beganhiscareerin1813asDavy's

LaboratoryAssistant.

In1834,Faradaydevelopedthetwolawsofelectrochemistry:

TheFirstLawofElectrochemistry

Theamountofasubstancedepositedoneachelectrodeofan

electrolyticcellisdirectlyproportionaltotheamountofelectricity

passingthroughthecell.

TheSecondLawofElectrochemistry

Thequantitiesofdifferentelementsdepositedbyagivenamountof

electricityareintheratiooftheirchemicalequivalentweights.


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Faradayalsodefinedanumberofterms:

Theanodeisthereforethatsurfaceatwhichtheelectriccurrent,accordingto

ourpresentexpression,enters:itisthenegativeextremityofthedecomposingbody;iswhereoxygen,chlorine,acids,etc.,areevolved;andisagainstoroppositethepositiveelectrode.

Thecathodeisthatsurfaceatwhichthecurrentleavesthedecomposingbody,andisitspositiveextremity;thecombustiblebodies,metals,alkalies,andbasesareevolvedthere,anditisincontactwiththenegativeelectrode.

Manybodiesaredecomposeddirectlybytheelectriccurrent,theirelementsbeingsetfree;theseIproposetocallelectrolytes....

Finally,Irequireatermtoexpressthosebodieswhichcanpasstotheelectrodes,or,astheyareusuallycalled,thepoles.Substancesarefrequentlyspokenofasbeingelectronegativeorelectropositive,accordingastheygounderthesupposedinfluenceofadirectattractiontothepositiveornegativepole...Iproposetodistinguishsuchbodiesbycallingthoseanionswhichgototheanodeofthedecomposingbody;andthosepassingtothecathode,

cations;andwhenIhaveoccasiontospeakofthesetogether,Ishallcallthemions.

thechlorideofleadisanelectrolyte,andwhenelectrolyzedevolvesthetwoions,chlorineandlead,theformerbeingananion,andthelatteracation.


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JohnFredericDaniell(17901845),professorofchemistryatKing's

College,London.

Daniell'sresearchintodevelopmentofconstantcurrentcellstook

placeatthesametime(late1830s)thatcommercialtelegraph

systemsbegantoappear.Daniell'scopperbattery(1836)became

thestandardforBritishandAmericantelegraphsystems.

In1839,Daniellexperimentedonthefusionofmetalswitha70

cellbattery.Heproducedanelectricarcsorichinultravioletrays

thatitresultedinaninstant,artificialsunburn.Theseexperiments

causedseriousinjurytoDaniell'seyesaswellastheeyesof

spectators.

Ultimately,Daniellshowedthattheionofthemetal,ratherthan

itsoxide,carriesanelectricchargewhenametalsaltsolutionis

electrolyzed.

Left:AnearlyDaniellCell

Right:Daniellcellsused

bySirWilliamRobert

Grove,1839.


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VoltaicCells

Inspontaneous

oxidation

reduction(redox)

reactions,

electronsare

transferredand

energyis

released.


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VoltaicCells

Ifthereactionisseparatedintotwoparts,wecanusethatenergytodoworkbymaking

the

electrons

flow

throughanexternaldevice.

Thistypeofsetup

iscalledavoltaiccell.


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VoltaicCells

Thisisatypicalvoltaiccell

Astripofzincmetalis

immersed

in

a

solution

ofZn(NO3)2 Astripofcoppermetal

isimmersedinasolutionofCu(NO3)2

ThetwosolutionsareconnectedbyasaltbridgecontainingNaNO3

Theoxidationoccursat

theanode(Zn) Thereductionoccursat

thecathode(Cu)


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VoltaicCells

Thesaltbridgeisusedto

preventelectronsflowing

directlyfromthezinctothe

copper ThesaltbridgeconsistsofaU

shapedtubethatcontainsa

saltsolution,sealedwith

porousplugs,oranagargel

solutionofthesalt Thesaltbridgekeepsthe

chargesbalancedandforces

theelectrontomovethrough

the

wire

Cationsmovetowardthe

cathode.

Anionsmovetowardthe

anode.


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VoltaicCells

Inthecell,electronsleavetheanodeand

flow

through

the

wiretothecathode.

Astheelectronsleavetheanode,the

cations

formed

dissolveintothesolutionintheanodecompartment.

Eventually,

if

the

cell

isusedforalongtime,theanode(zinc)willdissolve

l i ll


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VoltaicCells

Astheelectronsreachthecathode,cationsin

the

cathode

are

attractedtothenownegativecathode.

Theelectronsaretaken

by

the

cation,

and

the

neutralmetalisdepositedonthecathode.

Eventually,

if

the

cell

is

usedforalongtime,allthecopperionswillplateontothecopper

cathode

El t ti F ( f)


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ElectromotiveForce(emf)

Wateronly

spontaneously

flowsonewayinawaterfall.

Likewise,electrons

onlyspontaneouslyflowonewayina

redoxreaction

from

higher

to

lowerpotential

energy.


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Electromotive

Force

(emf) Thepotentialdifferencebetweentheanode

and

cathode

in

a

cell

is

called

the

electromotiveforce(emf).

Itisalsocalledthecellpotential,andis

designatedEcell.


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Cell

Potential

Cell

potential

is

measured

in

volts

(V).

1V=1J

C

Where J=Joules

C=Coulombs

Recallthat1electronhasachargeof1.6x1019 C

Standard Reduction Potentials


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StandardReductionPotentials

Thecellpotentialisthedifferencebetweentwo

electrodepotentials.

Byconvention,electrodepotentialsare

written

as

reductions

Reductionpotentialsformostcommonelectrodesaretabulatedasstandardreductionpotentials.

Standard Hydrogen Electrode


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StandardHydrogenElectrode

Electrodepotentialsarereferencedtoastandard

hydrogenelectrode(SHE).

Bydefinition,thereductionpotentialfor

hydrogenis0V:

2H+ (aq,1M)+2e

H2 (g,1atm)


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Standard

Cell

PotentialsThecellpotentialatstandardconditionsis

calculated

Ecell =Ered(cathode) Ered(anode)

Becausecellpotentialisbasedonthe

potential

energy

per

unit

of

charge,

it

is

anintensiveproperty.

Substancereduced Substanceoxidized

Cell Potentials


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CellPotentials

Oxidation: Ered = -0.76 V Reduction: Ered = +0.34 V

ll i l


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CellPotentials

Ecell = Ered (cathode) Ered (anode)

=+0.34V (0.76V)

=+1.10V

Generally, most of the common cells used, on the average,

generate approximately 1.5 V

Oxidizing and Reducing Agents


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OxidizingandReducingAgents

Thestrongestoxidizershavethe

mostpositivereductionpotentials.

Thestrongestreducershavethemostnegativereductionpotentials.

O idi i d R d i A t


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OxidizingandReducingAgents

Thegreaterthe

differencebetweenthe

two,thegreaterthe

voltageofthecell.

Free Energy


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FreeEnergy

Gforaredoxreactioncanbefoundby

usingtheequation

G= nFE

where:

nis

the

number

of

moles

of

electrons

transferredFisaconstant,theFaraday.

1F=96,485C/mol=96,485J/Vmol

E=ThestandardcellpotentialinV

Understandardconditions,

G = nFE

Nernst Equation


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NernstEquation

Rememberthat

G= G +RTlnQ

Thismeans

nFE= nFE +RTlnQ

Dividing

both

sides

by nF,

we

get

the

Nernst

equation:

or,usingbase10logarithms,

E=E RT

nF lnQ

E=E 2.303RT

nF logQ

Nernst Equation


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NernstEquation

Atroomtemperature(298K),and

R=8.314J/molK

F=96,485J/Vmol

ThefinalformoftheNernstEquationbecomes

E=E 0.0592n

logQ

2.303RT

F =0.0592V

Walther Hermann Nernst (1864 1941)


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WaltherHermannNernst(1864 1941)

Nernst'searlystudiesinelectrochemistrywereinspiredbyArrhenius'dissociationtheoryofionsinsolution.

In1889heelucidatedthetheoryofgalvaniccellsby

assumingan"electrolyticpressureofdissolution"which

forcesionsfromelectrodesintosolutionandwhichwas

opposedtotheosmoticpressureofthedissolvedions.

Also,in1889,heshowedhowthecharacteristicsofthe

currentproducedcouldbeusedtocalculatethefree

energychangeinthechemicalreactionproducingthe

current. Thisequation,knownastheNernstEquation,

relatesthevoltageofacelltoitsproperties.

IndependentlyofThomson,heexplainedwhy

compoundsionizeeasilyinwater.Theexplanation,

calledtheNernstThomsonrule,holdsthatitisdifficult

forchargedionstoattracteachotherthroughinsulating

watermolecules,sotheydissociate.

Concentration Cells


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ConcentrationCells

NoticethattheNernstequationimpliesthatacellcouldbe

createdthathasthesamesubstanceatbothelectrodes.

Forsuchacell,Ecell wouldbe0,butQwouldnot.

Therefore,aslongastheconcentrationsaredifferent,

Ewillnotbe0.


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ApplicationsofOxidationReductionReactions

Why Study Electrochemistry?Why Study Electrochemistry?


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y y yy y y

Batteries

Corrosion

Industrial productionof chemicals suchas Cl2, NaOH, F2

and Al Biological redox

reactions

The heme group

BATTERIES


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BATTERIESPrimary, Secondary, and Fuel Cells

Batteries


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BatteriesMost commercial batteries produce 1.5 V. To get a higher

voltage, batteries are joined together.

Dry Cell Battery


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Dry Cell Battery

Anode (-)

Zn Zn2+ + 2e-

Cathode (+)

2 NH4+ + 2e-2 NH3 + H2

Primarybattery usesredoxreactions

thatcannotberestoredbyrecharge.

Alkaline Battery


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Nearly same reactions as in common dry

cell, but under basic conditions.

Alkaline Battery

Anode(): Zn+2OH ZnO +H2O+ 2e

Cathode(+):2MnO2+H2O + 2e Mn2O3+2OH

Alkaline Batteries


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AlkalineBatteries

Lead Storage Battery


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Secondary

battery

Usesredoxreactionsthat

canbereversed.

Canberestoredbyrecharging

Lead Storage Battery


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Anode (-) Eo = +0.36 V

Pb + HSO4- PbSO4 + H

+ + 2e-

Cathode (+) Eo

= +1.68 VPbO2 + HSO4

- + 3 H+ + 2e- PbSO4 + 2 H2O

Ni-Cad Battery


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yAnode (-)

Cd + 2 OH- Cd(OH)2 + 2e-

Cathode (+)

NiO(OH) + H2O + e- Ni(OH)2 + OH-


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Fuel Cells: H2 as a FuelFuelcell reactantsare

suppliedcontinuouslyfrom

anexternalsource.

Carscanuseelectricity

generatedbyH2/O2fuel

cells.

H2carriedintanksor

generatedfrom

hydrocarbons.

HydrogenAir Fuel Cell


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HydrogenFuelCells


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y g

H as a Fuel


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2

Comparisonofthevolumesofsubstancesrequired

tostore4kgofhydrogenrelativetocarsize.


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Storing H2 as a Fuel

OnewaytostoreH2istoadsorbthegasontoa

metalormetalalloy.

Electrolysis


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Usingelectricalenergytoproducechemicalchange.

Sn2+(aq)+2Cl

(aq) Sn(s) + Cl2(g)

Electrolysis of Aqueous NaOH


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Anode (+)

4 OH- O2(g) + 2 H2O + 4e-

Cathode (-)

4 H2

O + 4e- 2 H2

+ 4 OH-

Eo for cell = -1.23 V

ElectricEnergyfChemicalChange

Anode Cathode

Electrolysis


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ElectricEnergy

fChemicalChange

BATTERY

+

Na+

Cl-

Anode Cathode

electrons

BATTERY

+

Na+

Cl-

Anode Cathode

electrons Electrolysisofmolten

NaCl. Hereabattery

pumpselectrons

fromCl toNa+.

NOTE:Polarityof

electrodes

is

reversed

frombatteries.

Electrolysis of Molten NaCl


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SeeFigure20.18

Electrolysis of Molten NaCl


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Anode (+)

2 Cl- Cl2(g) + 2e-

Cathode (-)

Na+ + e- Na

BATTERY

+

Na+Cl-

Anode Cathode

electrons

BATTERY

+

Na+Cl-

Anode Cathode

electrons

Eo forcell(inwater)=Ec Ea

= 2.71V (+1.36V)

= 4.07V(inwater)

ExternalenergyneededbecauseEo is().

Electrolysis of Aqueous NaCll t

l t


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Anode (+)2 Cl- Cl2(g) + 2e-

Cathode (-)2 H2O + 2e- H2 + 2 OH

-


Note that H2O is more

easily reduced than

Na+.

BATTERY

+

Na+Cl-

Anode Cathode

H2O

electrons

BATTERY

+

Na+Cl-

Anode Cathode

H2O

electrons

Also,Cl

isoxidizedinpreferencetoH2O

becauseofkinetics.Also,Cl

isoxidizedinpreferencetoH2O

becauseofkinetics.

Electrolysis of Aqueous NaCl


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Cells like these are the source of NaOH and Cl2.

In 1995: 25.1 x 109 lb Cl2 and 26.1 x 109 lb NaOH

AlsothesourceofNaOClforuseinbleach.

Electrolysis of Aqueous NaI


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Anode (+): 2 I- I2(g) + 2e-

Cathode (-): 2 H2O + 2e- H2 + 2 OH-

Eo

for cell = -1.36 V

Electrolysis of Aqueous CuCl2


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Anode (+)

2 Cl- Cl2(g) + 2e-

Cathode (-)

Cu2+ + 2e- Cu


Note that Cu is more

easily reduced thaneither H2O or Na

+.

BATTERY

+

Cu2+Cl-

Anode Cathode

electrons

H2O

BATTERY

+

Cu2+Cl-

Anode Cathode

electrons

H2O

Electrolytic Refining of Copper


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ImpurecopperisoxidizedtoCu2+ attheanode.Theaqueous

Cu2+ ionsarereducedtoCumetalatthecathode.

Thecopperatthecathodeisover99%pure

Producing Aluminumf


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2Al2O3 + 3Cf4Al + 3CO2

CharlesHall(18631914)developedelectrolysisprocess.

FoundedAlcoa.

Corrosionand


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CorrosionPrevention


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redox and electrochemistry.pdf

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