Download - 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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OxidationReductionReactions
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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OxidationReductionReactions
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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BalancingRedox Equationsbythe
HalfReactionMethod
4.
Balance
each
half
reaction.a. BalanceelementsotherthanHandO.
b. BalanceObyaddingH2O.
c. BalanceHbyaddingH+.d. Balancechargebyaddingelectrons.
5. Multiplythehalfreactionsbyintegersso
thattheelectronsgainedandlostarethesame.
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BalancingRedox Equationsbythe
HalfReactionMethod
6.
Add
the
half
reactions,
subtracting
things
thatappearonbothsides.
7. Makesuretheequationisbalanced
accordingtomass.8. Makesuretheequationisbalanced
accordingtocharge.
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BalancingRedox Equationsbythe
HalfReactionMethod
ConsiderthereactionbetweenMnO4 andC2O4
2 :
MnO4
(aq) +C2O42
(aq) Mn2+(aq) +CO2(aq) (acidicsolution)
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BalancingRedox Equationsbythe
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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BalancingRedox Equationsbythe
HalfReactionMethod
MnO4 + C2O42
Mn2+ +CO2
+7 +3 +4+2
TheMnO4 istheoxidizingagent
TheC2O42 isthereducingagent
Determinetheoxidizingandreducingagents
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BalancingRedox Equationsbythe
HalfReactionMethod
Theoxidationhalfreactionis
C2O42 CO2
Balance
the
half
reactionC2O4
2 2CO2
Thisbalancesboththecarbonatomsand
theoxygenatoms
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BalancingRedox Equationsbythe
HalfReactionMethod
Balancethechargebyadding2electronsto
therightside.
C2O42
2
CO2+
2
e
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BalancingRedox Equationsbythe
HalfReactionMethod
Thereductionhalfreactionis
MnO4
Mn2+
The
manganese
is
balancedInordertobalancethe4oxygens,weadd4watermoleculestotherightside.
MnO4
Mn2+ +4H2O
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BalancingRedox Equationsbythe
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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BalancingRedox Equationsbythe
HalfReactionMethod
8
H+
+
MnO4
Mn2+
+
4
H2O
Balancethecharges,add5e totheleftside.
5e +8H+ +MnO4
Mn2+ +4H2O
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BalancingRedox Equationsbythe
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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BalancingRedox Equationsbythe
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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BalancingRedox Equations
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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BalancingRedox Equations
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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BalancingRedox Equations
inBasicSolution
Thereductionhalfreactionis
MnO4
MnO2
Balancethehalfreaction
MnO4
MnO2
Balance
the
oxygen
atoms,
add
H+
and
H2O4 H+ + MnO4
MnO2+2H2O
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BalancingRedox Equations
inBasicSolution
Balancethecharges,add3e totheleftside
3e + 4 H+ + MnO4
MnO2+2H2O
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BalancingRedox Equations
inBasicSolution
Theoxidationhalfreactionis
CN CNO
SincetheCatomsarebalanced,balancethe
oxygen.
AddH2OtotheleftsideandH+ totherightside.
H2O
+
CN
CNO
+
2
H+
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BalancingRedox Equations
inBasicSolution
H2O + CN
CNO
+2H+
Balancethecharge,add2e totherightside
H2O + CN
CNO
+2H+
+ 2e
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BalancingRedox Equations
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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BalancingRedox Equations
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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BalancingRedox Equations
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
-
Eo for cell = -2.19 V
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
Eo for cell = -1.02 V
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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