ElektroKIMIA

UNIVERSITY INDONESIAOF

Elektro KimiaElektro Kimia

Galvanic CellsUNIVERSITY INDONESIAOF

Galvanic Cells

• Galvanic cell ‐ a spontaneous chemicalGalvanic cell a spontaneous chemicalreaction generates an electric current.

• Electrolytic cell an electric current drives a• Electrolytic cell ‐ an electric current drives anonspontaneous reaction.


Galvanic Cells

• Redox reaction.edo eact o .– Oxidation ‐ a loss of electrons (an increase in oxidationnumber).

– Reduction ‐ a gain of electrons (a decrease inoxidation number).

– Represent oxidation and reduction aspects of the– Represent oxidation and reduction aspects of thereaction with half‐reactions.

– Oxidizing agent ‐ species that causes oxidation tooccur and is itself reduced.

– Reducing agent ‐ species that causes reduction tooccur and is itself oxidizedoccur and is itself oxidized.


Galvanic Cells

• Redox reaction.Redox reaction.– If a spontaneous reaction is carried out in abeaker:

• oxidizing agent and reducing agent are in direct contact• electrons are directly transferred

th l f ti i l t t th di f• enthalpy of reaction is lost to the surroundings for anexothermic reaction

– If spontaneous reaction is carried out in a galvanicp gcell:

• chemical energy released by the reaction is convertedto electrical energyto electrical energy


Galvanic Cells

• For the reaction:For the reaction:

• Zn (s) + Cu2+ (aq) Zn2+ (aq) + Cu (s).U D i ll ll t f l i ll t– Use a Daniell cell, a type of galvanic cell, to carryout the reaction.

• consists of two half cells• consists of two half‐cells– beaker with a strip of Zn in a solution of ZnSO4

– beaker with a strip of Cu in a solution of CuSO4

• electrodes ‐ strips of zinc and copper

• salt bridge ‐ a U‐shaped tube that contains a gelpermeated with a solution of an inert electrolytepermeated with a solution of an inert electrolyte


Galvanic Cells


Galvanic Cells

• The electrons can be transferred only throughThe electrons can be transferred only throughthe wire.– oxidation and reduction half reactions occur at– oxidation and reduction half‐reactions occur atseparate electrodes

– electric current flows through the wireelectric current flows through the wire


Galvanic Cells

• Anode ‐ the electrode at which oxidation takesAnode the electrode at which oxidation takesplace.– the negative ( ) electrode– the negative (‐) electrode

– produces electrons

C th d th l t d t hi h d ti• Cathode ‐ the electrode at which reductiontakes place.

h ( ) l d– the positive (+) electrode

– consumes electrons


Galvanic Cells

• Anode and cathode half‐reaction must add toode a d cat ode a eact o ust add togive the overall cell reaction.

• Salt bridge maintains electrical neutrality by ag y yflow of ions.– anions flow through the salt bridge from the cathodet th d t tto the anode compartment

– cations migrate through the salt bridge from theanode to the cathode compartmentp

• Electrons move through the external circuit fromthe anode to the cathode.


Galvanic Cells

Oxidation‐reduction reactions part I


Galvanic Cells


Galvanic Cells


Galvanic Cells

Galvanic cells I: the copper‐zinc cell


Galvanic Cells


Galvanic Cells


Galvanic Cells


Galvanic Cells


Galvanic Cells

Galvanic cells II: the zinc‐hydrogen cell


Galvanic Cells

• EXAMPLE:EXAMPLE:

• Describe how you would construct a galvanic cell based on the following reaction:cell based on the following reaction:

Pb2+ (aq) + Zn (s) Pb (s) + Zn2+ (aq)Pb2+ (aq) + Zn (s) Pb (s) + Zn2+ (aq)


Galvanic Cells

• Pb2+ (aq) + 2 e‐ Pb (s)( q) ( )• Zn (s) Zn2+ (aq) + 2 e‐

• Looking at the two half‐reactions, we find that the Pb2+is being reduced, and the Zn is being oxidized.

• Therefore, the anode compartment of our cell wouldconsist of a strip of zinc metal immersed in a solutionconsist of a strip of zinc metal immersed in a solutioncontaining Zn2+ ions (such as zinc nitrate).

• The cathode compartment would consist of a strip ofp plead immersed in a solution containing Pb2+ ions (suchas lead (II) nitrate).


Galvanic Cells

• Pb2+ (aq) + 2 e‐ Pb (s)Pb (aq) + 2 e Pb (s)• Zn (s) Zn2+ (aq) + 2 e‐

• The two half‐cells would be connected to each• The two half‐cells would be connected to eachother with a salt bridge and an external wire.

• Electrons flow through the wire from the zinc• Electrons flow through the wire from the zincanode to the lead cathode.

• Anions move from the cathode compartment• Anions move from the cathode compartmenttowards the anode while cations migrate fromthe anode compartment toward the cathode.the anode compartment toward the cathode.


Galvanic Cells

Redox chemistry of iron and copper

Shorthand Notation f l i ll


for Galvanic Cells

• Single vertical line represents a phaseSingle vertical line, , represents a phaseboundary.

• Double vertical line ǁ represents a salt bridge• Double vertical line, ǁ, represents a salt bridge.



for Galvanic Cells

• Shorthand for the anode half‐cell is alwaysS o t a d o t e a ode a ce s a ayswritten on the left of the salt‐bridge symbol,followed on the right of the symbol by theh h d f h h d h lf llshorthand for the cathode half‐cell.– Reactants in each half cell are written first, followedby productsby products.

– Electrons move through the external circuit from leftto right.

– For Zn (s) + Cu2+ (aq) Zn2+ (aq) + Cu (s):Zn (s)Zn2+ (aq) ǁ Cu2+ (aq)Cu (s).



for Galvanic Cells

• Cell involving a gasCell involving a gas.– Additional vertical line due to presence ofadditional phaseadditional phase.

– List the gas immediately adjacent to theappropriate electrode.appropriate electrode.

• Detailed notation includes ion concentrationsand gas pressuresand gas pressures.



for Galvanic Cells


• Give the shorthand notation for a galvanic cell that employs the overall reactionthat employs the overall reaction

Pb(NO ) (aq) + Ni (s) Pb(s) + Ni(NO ) (aq)Pb(NO3)2(aq) + Ni (s) Pb(s) + Ni(NO3)2(aq)

Give a brief description of the cellGive a brief description of the cell.



for Galvanic Cells

• SOLUTION: The two half‐reactions for thisSOLUTION: The two half reactions for this overall reaction are:

Pb2+ (aq) + 2 e‐ Pb (s)

Ni (s) Ni2+ (aq) + 2 e‐

From these half‐reactions, we know that lead is being reduced and nickel is being oxidizedis being reduced and nickel is being oxidized.



for Galvanic Cells

• Therefore, Ni is the anode and Pb is theTherefore, Ni is the anode and Pb is thecathode. The cell notation is:

• Ni(s)Ni2+(aq)ǁPb2+(aq)Pb(s)Ni(s)Ni (aq)ǁPb (aq)Pb(s)• This cell would consist of a strip of nickel asthe anode dipping into an aqueous solution ofthe anode dipping into an aqueous solution ofNi(NO3)2 and a strip of Pb as the cathodedipping into an aqueous solution of Pb(NO3)2.3 2

• The two half‐cells would be connected by asalt bridge and a wire.

Cell Potentials and Free‐Energy Changes


Free Energy Changes for Cell Reactions

• Electromotive force (emf) ‐ the driving forceElectromotive force (emf) the driving force(electrical potential) that pushes the negativelycharged electrons away from the anode and pullsg y pthem toward the cathode.– Also called the cell potential (E) or the cell voltage.p ( ) g

– Potential of a galvanic cell is a positive quantity.

• Coulomb (C) ‐ the amount of charge transferred( ) gwhen a current of 1 ampere (A) flows for 1 s.– 1 J = 1 C x 1 V




• Cell potential ‐measured with a voltmeterCell potential measured with a voltmeter.– Gives a positive reading when the + and ‐terminals of the voltmeter are connected toterminals of the voltmeter are connected tocathode (+) and anode (‐), respectively.

• can use voltmeter‐cell connections to determine whichelectrode is the anode and which is the cathode




• Two driving forces of a chemical reaction: cellTwo driving forces of a chemical reaction: cellpotential, E and free‐energy change, G.

Related by G = nFE– Related by G = ‐nFE.• n = number of moles of electrons transferred in thereaction

• F (faraday) ‐ the electrical charge on 1 mol of electrons– 1 F = 96,500 C/mol e‐

• G and E have opposite signs– spontaneous reaction has a positive cell potential but negativeG




• Standard cell potential Eo ‐ the cell potentialStandard cell potential, E the cell potentialwhen both reactants and products are in theirstandard statesstandard states.– Solutes at 1 M concentration.

Gases at a partial pressure of 1 atm– Gases at a partial pressure of 1 atm.

– Solids and liquids in pure form.

T 25oC– T = 25oC.

• Go = ‐nFEo.

Standard Reduction i l


Potentials

E E E

– Can't measure potential of a single electrode.

Measure a potential difference by placing a

Ecell Eanode Ecathode

– Measure a potential difference by placing avoltmeter between two electrodes.

– Develop a set of standard half‐cell potentials– Develop a set of standard half‐cell potentials.• choose an arbitrary standard half‐cell as a referencepoint and assign an arbitrary potential

• express the potential of all other half‐cells relative tothe reference half‐cell



Potentials

• Standard hydrogen electrode (S H E ) ‐Standard hydrogen electrode (S.H.E.)reference half‐cell.– Corresponding half reaction assigned an– Corresponding half‐reaction ‐ assigned anarbitrary potential of exactly 0 V.

• 2 H+ (aq, 1 M) + 2 e‐ H2 (g, 1 atm) Eo 0 V2 H (aq, 1 M) 2 e H2 (g, 1 atm) E 0 V

– Shorthand notation for S. H. E.• H+ (1 M)H2 (1 atm)Pt (s)( ) 2 ( ) ( )



Potentials• Determine standard potentials for half‐cells by constructing

a galvanic cell in which the half‐cell of interest is paired upwith the standard hydrogen electrode.– Standard oxidation potential ‐ the corresponding half‐cell

t ti l f id ti h lf tipotential for an oxidation half‐reaction.– Standard reduction potential ‐ the corresponding half‐cell

potential for a reduction half‐reaction.Whenever the direction of a half reaction is reversed the sign– Whenever the direction of a half‐reaction is reversed, the signof Eo must be reversed.

• The standard oxidation potential and the standard reduction potentialalways have the same magnitude, but they have opposite signs

• Construct a table of standard reduction potentials(Appendix D in text).



Potentials

• Conventions used in constructing a table ofConventions used in constructing a table ofhalf‐cell potentials:– The half reactions are written as reductions– The half‐reactions are written as reductions.

• oxidizing agents and electrons are on the reactant side

• reducing agents are on the product sidereducing agents are on the product side

– The half‐cell potentials are standard reductionpotentials.p

• also known as standard electrode potentials



Potentials

• Conventions used in constructing a table ofConventions used in constructing a table ofhalf‐cell potentials:– The half reactions are listed in order of decreasing– The half‐reactions are listed in order of decreasingstandard reduction potential.

• strongest oxidizing agents are located in the upper leftstrongest oxidizing agents are located in the upper leftof the table

• strongest reducing agents are in the lower right of thebltable



Potentials

• Conventions used in constructing a table ofConventions used in constructing a table ofhalf‐cell potentials:– Ordering of half reactions correspond to ordering– Ordering of half‐reactions correspond to orderingof the oxidation reactions in the activity series.

• the more active metals at the top of the activity seriesthe more active metals at the top of the activity serieshave the more positive oxidation potential (morenegative reduction potential)



Potentials



Potentials

Standard reduction potentials

Using Standard R d i P i l


Reduction Potentials

• Table of standard reduction potentials ‐ab e o sta da d educt o pote t a ssummarizes an enormous amount of chemicalinformation.– Can arrange any two or more oxidizing or reducingagents in order of increasing strength.Predict the spontaneity or nonspontaneity of– Predict the spontaneity or nonspontaneity ofthousands of redox reactions.

• combine half‐reactions of interest and use• may need to multiply half‐reactions by some factor toensure that electrons cancel

– do not multiply values of Eo for the half‐reactions by that factor




• Eo values are independent of the amount ofE values are independent of the amount ofreaction.

Go = nFEo– Go = ‐nFEo

• Go is an extensive property because it depends on theamount of substance

• change the amount of substance that reacts, Go

changes by the same amount as does n, the number ofelectrons transferredelectrons transferred

• Eo = ‐Go/nF remains constant




• Can predict the spontaneity of a reaction byCan predict the spontaneity of a reaction byknowing the location of the oxidizing andreducing agent in the tablereducing agent in the table.– An oxidizing agent can oxidize any reducing agentthat lies below it in the table.that lies below it in the table.

• Eo for overall reaction must be positive





• Write the balanced net ionic equation, andcalculate Eo for the following galvanic cell:calculate Eo for the following galvanic cell:

Al (s)Al3+ (aq) ǁ Cu2+ (aq)Cu (s)Al (s)Al3+ (aq) ǁ Cu2+ (aq)Cu (s)



Potentials




• Al (s)Al3+ (aq) ǁ Cu2+ (aq)Cu (s)Al (s)Al3+ (aq) ǁ Cu2+ (aq)Cu (s)• Al (s) is the anode and therefore undergoesoxidation while Cu (s) is the cathode and Cu2+oxidation while Cu (s) is the cathode and Cu2+

therefore, undergoes reduction.

Th h lf i d h i ll i l• The half‐reactions and their cell potentialsare:

• Al (s) Al3+ (aq) + 3 e‐ Eo=+1.66V

• Cu2+ (aq) + 2 e‐ Cu (s) Eo = +0.34 V( q) ( )




• Al (s)Al3+ (aq) ǁ Cu2+ (aq)Cu (s)• Al (s) Al3+ (aq) + 3 e‐ Eo=+1.66V• Cu2+ (aq) + 2 e‐ Cu (s) Eo = +0.34 V

V 00.2V 34.0V 66.1oCuCuAlAl

ocell +2+3

EEE o

• Notice that although we multiplied the coefficients in bothhalf‐reactions by the factor of 2 and 3 respectively, we didnot multiply the values of Eo by these factors.

• This is because Eo values are independent of the amount ofreaction.

Cell Potentials and Composition of the Reaction


pMixture: The Nernst Equation

• Cell potentials depend on temperature and onCell potentials depend on temperature and onthe composition of the reaction mixture.

G = Go + RT ln Q– G = Go + RT ln Q.• G = ‐nFE;

• Go = ‐nFEoG nFE

– ‐nFE = ‐nFEo + RT ln Q.




• Nernst equation:Nernst equation:

(i l 2 C)

E E

0.0592n

logQ

• (in volts at 25oC).– Enables us to calculate cell potentials under

d d dnonstandard‐state conditions.




• Calculate Ecell for the following cell reaction:Calculate Ecell for the following cell reaction:

• 2 Cr(s) + 3Pb2+(aq) 2 Cr3+ (aq) + 3Pb(s)[ b2+] 0• [Pb2+] = 0.15 M;

• [Cr3+] = 0.50 M



Potentials




• 2 Cr (s) 2 Cr3+ + 6 e‐ Eo = +0 74 V2 Cr (s) 2 Cr + 6 e E = +0.74 V

• 3 Pb2+ (aq) + 6 e‐ 3 Pb (s) Eo=‐0.13V




Ecell V 0 74 013 0 61. . . cell

E E 0.0592

nlogQ

n

Ecell Ecell

0.05926

logCr3+ 2Pb2 + 3

6 Pb2 + Ecell 0.61

0.05926

log0.5 2

0 15 3 0.59

6 0.15

Electrochemical D i i f H


Determination of pH

• Important application of Nernst equation• Important application of Nernst equation ‐electrochemical determination of pH using apH meterpH meter.

• Consider a cell with a hydrogen electrode asth d d d f l t dthe anode and a second reference electrodeas the cathode.

– Pt (s)H2 (1 atm)H+ (? M)ǁreference cathode.

Ecell 0.0592pH Eref



Determination of pH

Ecell 0.0592pH Eref

pH Ecell Eref

0 0592

• can measure the pH of a solution by

p0.0592

p ymeasuring Ecell

• Actual pH measurements use a glass electrodeActual pH measurements use a glass electrodewith a calomel electrode as the reference.



Determination of pH

• EXAMPLE:• EXAMPLE:

• The following cell has a potential of 0.49 V.C l l h H f h l i i h dCalculate the pH of the solution in the anodecompartment.

• Pt(s) H2(g) (1 atm)H+(pH = ?)Cl‐(aq) (1M)Hg2Cl2 (s) Hg (l)



Determination of pH

• The cell reaction is• The cell reaction is• Hg2Cl2 (s) + H2 (g) 2 Hg (l) + 2 Cl‐ (aq) + 2 H+ (aq)

••



Potentials



Determination of pH

• The cell reaction is• The cell reaction is• Hg2Cl2 (s) + H2 (g) 2 Hg (l) + 2 Cl‐ (aq) + 2 H+ (aq)

••V 28.0V 28.0V 00.0o

ClHg,ClHgo

HHo

-22

+2

EEE

pH Ecell Eref

0.0592

55.30592.0

V 28.0V 49.0pH

Standard Cell Potentials and


Potentials and Equilibrium Constants

• Standard free energy change for a reaction is• Standard free‐energy change for a reaction isrelated to both the standard cell potential andthe equilibrium constantthe equilibrium constant.– Go = ‐nFEo.

Go RT l K– Go = ‐RT ln K.




• Can combine the two equations• Can combine the two equations.

Eo RTnF

ln K 2.303RT

nFlogK

nF nF

E 0.0592

nlogK

• common use ‐ calculating equilibrium

n

constants from standard cell potentials.




• Equilibrium constants for redox reactions tend• Equilibrium constants for redox reactions tendto be either very large or very small incomparison with equilibrium constants forcomparison with equilibrium constants foracid‐base reactions.

P iti l f Eo d t K > 1– Positive value of Eo corresponds to K > 1.

– Negative value of Eo corresponds to K < 1.




• Three different ways to determine the value of• Three different ways to determine the value of an equilibrium constant K:

K C c D d

A a B b

ln K

G

RT

ln K

nFE

RT




• EXAMPLE:• EXAMPLE:

• Calculate the equilibrium constant for thef ll i i 25oCfollowing reaction at 25oC.

• 5 S2O82‐ (aq) + I2 (s) + 6 H2O (l) 10 SO4

2‐ (aq) + 2 IO3‐ (aq) + 12 H+ (aq)



Potentials




• S2O82‐ (aq) + 2 e‐ 2 SO4

2‐ (aq) Eo = +2 01 VS2O8 (aq) + 2 e 2 SO4 (aq) E = +2.01 V

• I2(s) + 6H2O (l) 2IO3‐ (aq) + 12 H+(aq) +10e‐ Eo = ‐1.20 V

h l f f hi i i 10

Ecell 2.01 1.20 0.81 V

• The value of n for this reaction is 10.

log K (10)(0.81)

137E 0.0592

logK log K (0.0592)

137

K = 10137

E

nlogK




BatteriesUNIVERSITY INDONESIAOF

Batteries

• Most important practical application ofMost important practical application ofgalvanic cells is their use as batteries.

• Features required in a battery depend on theFeatures required in a battery depend on theapplication.

• C. General features.C. General features.– Compact and lightweight.– Physically rugged and inexpensive.y y gg p– Provide a stable source of power for relativelylong periods of time.


Batteries

• Lead storage battery.Lead storage battery.– Used as a reliable source of power for startingautomobiles for more than three‐quarters of acentury.

– 12 V battery ‐ six 2 V cells connected in series.– Anode ‐ a series of lead grids packed with spongylead.Cathode a series of grids packed with lead– Cathode ‐ a series of grids packed with leaddioxide, dipped into an aqueous solution of H2SO4(38% w/w).


Batteries

• Lead storage battery.Lead storage battery.– Electrode half‐reactions and the overall cellreaction.

• Anode:• Pb (s) + HSO4

‐ (aq) PbSO4 (s) + H+ (aq) + 2 e‐ Eo = 0.296 V• Cathode:• Cathode:• PbO2 (s) + 3 H+ (aq) + HSO4

‐ (aq) + 2 e‐ 2 PbSO4 (s) + 2 H2O (l) Eo = 1.628 V• Overall:• Pb (s) + PbO (s) + 2 H+ (aq) + 2 HSO ‐ (aq) 2 PbSO (s) + 2 H O (l)• Pb (s) + PbO2 (s) + 2 H+ (aq) + 2 HSO4 (aq) 2 PbSO4 (s) + 2 H2O (l)• E0= 1.924 V

•


Batteries

• PbSO4 adheres to the surface of thePbSO4 adheres to the surface of theelectrodes.– recharge by using an external source of direct– recharge by using an external source of directcurrent to drive the cell reaction in the reverse,nonspontaneous directionp


Batteries

• Dry‐Cell Batteries (Leclanché cell) ‐ commonDry Cell Batteries (Leclanché cell) commonhousehold batteries.– Anode Zn metal can– Anode ‐ Zn metal can.

– Cathode ‐ inert graphite rod surrounded by apaste of solid MnO2 and carbon blackpaste of solid MnO2 and carbon black.

– Electrolyte ‐ a moist paste of NH4Cl and ZnCl2 instarch.starch.

• surrounds the MnO2 containing paste

• acidic ‐ causes corrosion of the Zn anode (Zn Zn2+)


Batteries

• Dry‐Cell Batteries (Leclanché cell) ‐ commonDry Cell Batteries (Leclanché cell) commonhousehold batteries.– Alkaline dry cell ‐modified version of Leclanché.Alkaline dry cell modified version of Leclanché.

• replace acidic NH4Cl (acidic) with NaOH or KOH

• electrode reactions ‐ oxidation of zinc and reduction ofd dmanganese dioxide

– produces ZnO due to basic conditions

– zinc corrodes more slowly

– battery has a longer life

• produces higher power and more stable current and voltage– more efficient ion transport in the alkaline electrolytemore efficient ion transport in the alkaline electrolyte


Batteries

• Mercury battery ‐ used in watches heartMercury battery used in watches, heartpacemakers, and other devices.– small size– small size

– anode ‐ Zn (same as dry cell)

cathode steel in contact with HgO in an alkaline– cathode ‐ steel in contact with HgO in an alkalinemedium of KOH and Zn(OH)2


Batteries

• Lithium batteries – light weight high voltageLithium batteries light weight, high voltage,rechargeable battery– Anode – lithium metal– Anode – lithium metal.

• highest standard oxidizing potential

– Cathode – metal oxide or sulfide that canCathode metal oxide or sulfide that canincorporate Li+.

– Electrolyte – lithium salt in an organic solvent.Electrolyte lithium salt in an organic solvent.

– Used in cell phones, laptop computers andcameras.


Batteries

• Fuel cells ‐ a galvanic cell in which one of theFuel cells a galvanic cell in which one of thereactants is a traditional fuel.– Reactants are not self contained within the cell– Reactants are not self‐contained within the cell.

• supplied from an external reservoir

– Best‐known ‐ hydrogen‐oxygen fuel cellBest known hydrogen oxygen fuel cell.• used in space vehicles as a source of electric power


Batteries


Batteries


Batteries


Batteries

CorrosionUNIVERSITY INDONESIAOF

Corrosion

• Corrosion ‐ the oxidative deterioration of aCorrosion the oxidative deterioration of ametal.

• Well known example of corrosion conversion• Well‐known example of corrosion ‐ conversionof iron to rust.

R i b th d t– Requires both oxygen and water.

– Involves pitting of the metal surface.i d i d l i h i ll d f• rust is deposited at a location physically separated from

the pits


Corrosion

• Proposed mechanism for formation of rust ‐ anoposed ec a s o o at o o ust aelectrochemical process in which iron is oxidizedin one region of the surface and oxygen isd d hreduced in another region.

– Anode region:• Fe (s) Fe2+(aq) +2 e Eo = 0 45 V• Fe (s) Fe2+(aq) +2 e‐ Eo = 0.45 V

– Cathode region:• O2 (g) + 4 H+ (aq) + 4 e‐ 2 H2O (l) Eo = 1.23 V2 2 o

– Electrons flow from the anode to the cathode throughthe metal.


Corrosion

• Proposed mechanism for formation of rust ‐Proposed mechanism for formation of rustan electrochemical process in which iron isoxidized in one region of the surface andoxidized in one region of the surface andoxygen is reduced in another region.– Ions migrate through the water droplets– Ions migrate through the water droplets.

• Fe2+ reacts with O2 and is oxidized to Fe3+

• Fe3+ reacts with H2O to form Fe2O3.xH2O (s) (rust)Fe reacts with H2O to form Fe2O3.xH2O (s) (rust)


Corrosion

• Proposed mechanism for formation of rust ‐ anProposed mechanism for formation of rust anelectrochemical process in which iron is oxidizedin one region of the surface and oxygen isg ygreduced in another region.– Explains why cars rust more rapidly when road salt isp y p yused to melt snow and ice.

• dissolved salt in water greatly increases the conductivity ofthe electrolytethe electrolyte

– O2 ‐ able to oxidize all metals except a few.• O2/H2O half‐reaction lies above the Mn+/M half‐reactionO2/H2O half reaction lies above the M /M half reaction


Corrosion

• Prevention of corrosion ‐ shield the metalPrevention of corrosion shield the metalsurface from oxygen and moisture.– Durable surface coating metals such as– Durable surface coating ‐ metals such aschromium, tin, or zinc.

– Galvanizing ‐ coating by dipping into a bath ofGalvanizing coating by dipping into a bath ofmolten zinc.

– Cathodic protection ‐ protecting a metal fromCathodic protection protecting a metal fromcorrosion by connecting it to a second metal thatis more easily oxidized.


Corrosion


Corrosion

Electrolysis and l l ll


Electrolytic Cells• Electrolytic cell ‐ an electric current is used toElectrolytic cell an electric current is used todrive a nonspontaneous reaction.– Processes occurring in galvanic and electrolytic– Processes occurring in galvanic and electrolyticcells are the reverse of each other.

• Electrolysis the process of using an electric• Electrolysis ‐ the process of using an electriccurrent to bring about chemical change.



Electrolytic Cells• Electrolytic cell.y

– Two electrodes that dip into an electrolyte and areconnected to a battery or some other source of directelectric current.electric current.

• battery ‐ an electron pump, pushing electrons into oneelectrode and pulling them out of the other electrode

– Anode ‐ electrode where oxidation takes placeAnode electrode where oxidation takes place.• positive sign• the battery pulls electrons out of it

3 Cathode electrode where reduction takes place– 3. Cathode ‐ electrode where reduction takes place.• negative sign• the battery pushes electrons into it



Electrolytic Cells• Electrolysis of molten NaClElectrolysis of molten NaCl.

– Cathode – attracts Na+.• Na+ + e‐ Na (l)• Na + e Na (l)

– Anode – attracts Cl‐.• 2 Cl‐ Cl2 (g) + e‐2 Cl Cl2 (g) + e



Electrolytic Cells• Electrolysis of aqueous NaCl.y q

– Electrode reactions in an aqueous solution can differfrom those for a molten salt.Cathode reaction can involve the reduction of Na+ or– Cathode reaction can involve the reduction of Na orthe reduction of water.

• reduction of water preferred.E8o less negative for H O than Eo– E8o less negative for H2O than Eo.

• 2 H2O (l) + 2 e‐ H2 (g) + 2 OH‐ (aq)– Anode reaction can involve the oxidation of Cl‐ or theoxidation of wateroxidation of water.

• actual reaction – the oxidation of Cl‐ due to overvoltage• 2 Cl‐ (aq) Cl2 (g) + 2 e‐



Electrolytic Cells• Electrolysis of aqueous NaCl.y q

– Overvoltage – amount of voltage needed above thecalculated standard reduction (or oxidation) potentialfor electrolysis to occur.for electrolysis to occur.

• needed when the half‐reaction has a substantial barrier forelectron transfer (slow rate).

– surmounts barrier– reaction proceeds at satisfactory rate

• small overvoltage needed for solution or deposition ofmetals

• large overvoltage needed for formation of O2 or H2• can’t predict; need experimental evidence if cell potentialsare similar



Electrolytic Cells• Electrolysis of aqueous NaClElectrolysis of aqueous NaCl.

– Overall cell reaction:– 2 Cl‐ (aq) + 2 H O (l) Cl (g) + 2 H (g) + 2 OH‐ (aq)– 2 Cl (aq) + 2 H2O (l) Cl2 (g) + 2 H2 (g) + 2 OH (aq)

• Na+ is a spectator ion and reacts with the OH‐ to formNaOH



Electrolytic Cells• F. Electrolysis of water.. ect o ys s o ate .

– Electrolysis of any aqueous solution requires thepresence of an electrolyte to carry the current inl tisolution.

– If the ions of the electrolyte are less easily oxidizedand reduced than water is, then water will react at,both electrodes.

• Anode: 2 H2O (l) O2 (g) + 4 H+ (aq) + 4 e‐.• Cathode 4 H O (l) + 4 e‐ 2 H (g) + 4 OH‐ (aq)• Cathode: 4 H2O (l) + 4 e‐ 2 H2 (g) + 4 OH‐ (aq).

– Overall cell reaction:• 2 H2O (l) O2 (g) + 2 H2 (g)2 2 g 2 g



Electrolytic Cells



Electrolytic Cells



Electrolytic Cells

Electrolysis of water



Electrolytic Cells



Electrolytic Cells



Electrolytic Cells



Electrolytic Cells

Electroplating



Electrolytic Cells

Commercial Applications of


Applications of Electrolysis

• Manufacture of sodium ‐ produced commercially in aDowns cell by electrolysis of a molten mixture of NaCland CaCl2.– Liquid Na produced at the cylindrical steel cathode is lessq p ydense than the molten salt and thus floats to the top partof the cell, where it is drawn off into a suitable container.

• Manufacture of chlorine and sodium hydroxide ‐a u ac u e o c o e a d sod u yd o deelectrolysis of aqueous NaCl.– Basis of chlor‐alkali industry.– Anode and cathode reactions for electrolysis of aqueous– Anode and cathode reactions for electrolysis of aqueousNaCl carried out in membrane cell.

• membrane keeps Cl2 and OH‐ apart but allows a current of Na toflow




• Manufacture of aluminum ‐ Hall‐Heroult process.p– Electrolysis of a molten mixture of Al2O3 and cryolite(Na3AlF6) at 1000o C in a cell with graphite electrodes.

• success of process ‐ the use of cryolite as a solventsuccess of process the use of cryolite as a solvent– Electrode reactions involve the formation of complexions.

• ions are reduced at the cathode to produce Al (l)• ions are reduced at the cathode to produce Al (l)• ions are oxidized at the anode to produce O2 (g)

– O2 (g) reacts with the graphite electrode to produce CO2 (g)– Requires frequent replacement of the anodes– Requires frequent replacement of the anodes

– Largest single consumer of electricity in the U.S.• 1 mol of electrons produces only 9 g of Al




• Electrorefining ‐ the purification of a metal byElectrorefining the purification of a metal bymeans of electrolysis.

• Electroplating the coating of one metal on• Electroplating ‐ the coating of one metal onthe surface of another using electrolysis.

C th d bj t t b l t d ( f ll l d)– Cathode ‐ object to be plated (carefully cleaned).

– Electrolytic cell contains a solution of ions of themetal to be depositedmetal to be deposited.

Quantitative Aspects of


Aspects of Electrolysis

• The amount of substance produced at anThe amount of substance produced at anelectrode by electrolysis depends on thequantity of charge passed through the cell.– Follows directly from the stoichiometry of thereaction and the atomic mass of the product.

• Moles of electrons passed through a cell aredetermined from the electric current and theti th t th t fltime that the current flows.

M l f - h (C)1 mol e -

Moles of e charge(C) 96,500 C




• Sequence of conversion used to calculate the mass orvolume of product produced by passing a knowncurrent for a fixed period of time.

current moles grams orand charge moles of e‐ of liters oftime product product

• Think of electrons as reactants in a balanced equationand proceed as with any other stoichiometry problem.p y y p








• How many grams of Cl2 would be produced inthe electrolysis of molten NaCl by a current ofthe electrolysis of molten NaCl by a current of4.25 A for 35.0 min?




• Remember that a coulomb is an A.s or that anampere is C/s.

• 2 Cl‐ Cl2 + 2 e‐

• moles of electrons = 2

4.25 Cs 35.0 min

60 s1 min

1 mol e -

96,500 C

1 mol Cl 22 mol e -

70.9 g Cl 21 mol Cl - 3.28 g Cl2

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