03 Electrochemistry Thermodynamics and Electrode Potential

8/10/2019 03 Electrochemistry Thermodynamics and Electrode Potential

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ELECTROCHEMISTRYELECTROCHEMISTRY

Deals with chemical changes produced by an electriccurrent and with the production of electricity by chemicalreactions

All electrochemical reactions involve transfer of electronsand are redox (oxidation-reduction) reactions

Electrochemical reactions take place in electrochemicalcell (an apparatus that allows a reaction to occur through

an external conductor)


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ELECTROCHEMICAL CELLSELECTROCHEMICAL CELLSTwo types:

1. Electrolytic cells: - these are cells in which an externalelectrical source forces a nonspontaneous reaction tooccur

(one common process is called electrolysis)

(not to be covered here)

2. Voltaic cells: - also called galvanic cells. In thesecells spontaneous chemical reactions generateelectrical energy and supply it to an external circuit

(corrosion and fuel cells are examples)


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Electric current enters and exits the cell by electrodes -electrodes are surfaces upon which oxidation or reductionhalf-reactions occur

Two kinds of electrodes:

Cathode: - electrode at which reduction occurs(electrons are gained by a species)

Anode: - electrode at which oxidation occurs (aselectrons are lost by some species)

ELECTRODESELECTRODES(as also covered earlier)(as also covered earlier)


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GALVANIC CELLSGALVANIC CELLS

Cells in which spontaneous reactions produces electrical

energy

The two half-cells are separated so that electron transferoccurs through an external circuit

Each half-cell contains the oxidized and reduced forms ofspecies in contact with each other

Half-cells linked by a piece of wire and a salt bridge


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+


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A salt bridge has three functions

1. It allows electrical contact between the two half-cells

2. It prevents mixing of the electrode solutions

3. It maintains electrical neutrality in each half-cellas ions flow into and out of the salt bridge

Note that: No electron flow between the two electrodes in separate solutionswithout the salt bridge.


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In galvanic cells, voltage/potential difference between the

electrodes drops to 0 as the reaction proceeds

Zn Zn 2+(1.0 M) Cu 2+(1.0 M) Cu

Electrode

Salt bridge

Species (withconcentrations) incontact with electrodes


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The Silver-Copper Cell

Composed of two half-cells:1. A strip of copper immersed in 1 M CuSO42. A strip of silver immersed in 1 M AgNO3

Experimentally we see:- Initial voltage (potential difference between electrodes

when not connected yet) is 0.46 volts

- When electrodes connected- The mass of the copper electrode decreases- The mass of the silver electrode increases- [Cu2+] increases and [Ag+] decreases


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Cu Cu2+ + 2e- (oxidation, anode)

2(Ag+ + e- Ag) (reduction, cathode)

2Ag+ + Cu Cu2+ + Ag (overall cell reaction)

Cu |Cu2+(1.0 M) ||Ag+(1.0 M) | Ag

Notice that in this case the copper electrode is the anode

The Silver-Copper Cell (cont.)


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The standard equilibrium cell potential (Ecell) is for the cell operatingunder standard state conditions

For an electrochemical cell, standard conditions are:- solutes at 1 M concentrations

- gases at 1 atm partial pressure

- solids and liquids in pure form

- all at some specified temperature, usually 298 K

STANDARD ELECTRODE POTENTIALSTANDARD ELECTRODE POTENTIAL


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EoZn= -0.762 vs SHEZn and Pt electrodes are at equilibrium with their environments at standardconditions and they are are not connected to eachother by a conductorso no current flows between Zn and Pt. The potential measured (EoZn) is theequilbrium potential at standard conditions vs SHE.

The electrode potentials are commonly measured versus the StandardHydrogen Electrode (SHE): E = 0.00 V


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EQUILIBRIUM POTENTIALEQUILIBRIUM POTENTIAL

When we immerse a metal in solution, there will be a tendency for the

metal to react with the solution, either with metal atoms dissolving ascations or cations already in the solution depositing as metal atoms:

Zn Zn2+ + 2e-

Zn2+ + 2e- Zn

As a result of these reactions, the metal will tend to accumulate anegative or positive charge. The build-up of this charge on the metal willchange its potential in such a way as to inhibit the reaction generatingthe charge until the potential reaches a value at which the rates of thetwo reactions are equal and opposite. This is known as the equilibriumpotential, and is the potential the metal will adopt in the solution in theabsence of any other reactions.


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For covenience all half-cell reactions are written in the table inreduction form and the electrode potentials are invariant

e.g. Zn=Zn2++2e- and Zn2++2e-=Zn are identical and represent zincin equilibrium with its ions with a potential of -0.762 V vs SHE

The more positive the E value for a half-reaction the greater thetendency for the reaction to proceed as written (in reduction-cathodicform) at standard conditions

The more negative the E value, the more likely is the reverse of thereaction as written (in oxidation-anodic form) at standard conditions

However, usually concentrations of reactants differ from one anotherand also change during the course of a reaction, in that case Ecell and

the actual Ecell are related by the Nernst Equation

STANDARD ELECTRODE POTENTIAL (cont.)STANDARD ELECTRODE POTENTIAL (cont.)


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E = equilibrium potential under the nonstandard conditions

E = equilibrium potential at standard state (all reactants and products at unit activity)R = gas constant, 8.314 J/mol.K

T = absolute temperature

n = number of equivalents (moles of electrons) transferred

F = Faradays constant, 96,485 J/V equivalent mol (= C/equivalent mol)

K =

p.s. If the equation is to be used for determining the equilibrium potential for a halfcell rxn, then the half cell rxn is to be written in reduction form or the (-) sign inthe equation is to be replaced by (+) sign in case the half cell rxn is written in

oxidation form.

tscoefficientheirofpowerthetoraisedreactantsofactivities

tscoefficientheirofpowerthetoraisedproductsofactivities

THE NERNST EQUATIONTHE NERNST EQUATION

E = E - (RT/nF) lnK


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Notation: E for the overall cell rxn, e for the half-cell rxns

p.s. If the equation is to be used for determining the equilibrium potential fora half cell rxn, then the half cell rxn is to be written in reduction form orthe (-) sign in the equation is to be replaced by (+) sign in case

the half cell rxn is written in oxidation form.


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Activity of a dissolved species A (A) is equal to its concentration inmoles per 1000 grams of water (molality) multiplied by the activity

coefficient, f.

Activity coefficients are extensively tabulated in numerous chemicaland electrochemical handbooks.

Activity of a gas is approximated at ordinary pressures by its partialpressure in atmospheres (atm).

The activities of pure solids and water are set equal to unity inaquaeous solutions.

At 25C, 2.303RT/F = 0.0592 V equivalent

ACTIVITYACTIVITY


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ACTIVITY (Cont.)ACTIVITY (Cont.)

Activity Coefficients of Strong Electrolytes (M=molality)

[Corrosion and Corrosion Control, H. H. Uhlig and R. W. Revie, John Wiley & Sons, 1985.]


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ACTIVITY (Cont.)ACTIVITY (Cont.)

Activity Coefficients of Strong Electrolytes (M=molality)



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ACTIVITY (Cont.)ACTIVITY (Cont.)Activity Coefficients of Strong Electrolytes (M=molality)



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PREDICTION OF SPONTANEITY

1. First write the half-cell rxn with the more positive (less

negative) E for the reduction-cathodic along with its halfcell electrode potential

2. Write the other half-cell rxn as an oxidation-anodic andinclude its half cell electrode potential

3. Balance the electron transfer

4. Obtain the cell rxn by adding the reduction and oxidation

half-cell rxns.5. Determine the overall cell potential, Ecell, from

Ecell= Ecathode- Eanode


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Note that half-cell reaction potentials are the same

regardless of the species stoichiometric coefficient inthe balanced equation.

Ecell > 0 Forward reaction (left-to-right) isspontaneous

Ecell < 0 Backward reaction (right-to-left) isspontaneous


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EXAMPLE PROBLEM 1EXAMPLE PROBLEM 1Notation: e for the half-cell rxns, E for the overall cell rxn.


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EXAMPLE BROBLEM 2EXAMPLE BROBLEM 2

Notation: E for the overall cell rxn, e for the half-cell rxns


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EXAMPLE PROBLEM 3EXAMPLE PROBLEM 3

Notation: e for the half-cell rxns, E for the overall cell rxn.

Assuming standard states for all reactants and products, determine thespontaneous direction of the following reactions by calculating the cell potential:

CuCl2 + H2 = Cu + 2HCl

CuCl2 + H2 Cu + 2HCl


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Note that the sign notation used in the text is different and

confusing.

The text uses

E = ec + ea

which is actually identical to the notation presented above

since the text takes ea as the negative value of the anodehalf-cell electrode potential.

A NOTE FOR THE SIGN NOTATION


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If a reference electrode other than SHE is used to

measure the equilibrium potential of a reaction, thepotential of the reference electrode relative to SHEshould be added to the measured potential if one wants

to determine the equilibrium potential of the reactionrelative to SHE

The reference electrodes are also used to measure the

corrosion potential of a corrosion cell (Ecorr) (to becovered later)

SECONDARY REFERENCE ELECTRODES


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Potential Values for Common Secondary Reference Electrodes

(Standard Hydrogen Electrode included for reference)


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EXAMPLE PROBLEMEXAMPLE PROBLEM((PrbPrb. 2.15 in. 2.15 in Principles and Prevention of CorrosionPrinciples and Prevention of Corrosion , Denny Jones, 1996), Denny Jones, 1996)

A corrosion potential of -0.229 V versus SCE was measured for a corroding alloy. What

is the potential versus (a) SHE, (b) Ag/AgCl (saturated), (c) Cu/saturated CuSO4?

(a)

(b)

(c) Home Exercise

The Pourbaix (Equilibrium Potential pH)


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The Pourbaix (Equilibrium Potential-pH)

Diagram

Potential

H2O is stable

H2 is stable

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2H+ + 2e- = H2

2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8-1.2

0

2H2O = O2 + 4H+ + 4e-O2 is stable


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Potential

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2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8-1.2

0

Pourbaix Diagram for Zinc

Equilibrium forZn Zn2+ + 2e-

Equilibrium forZn2+ + 2OH- Zn(OH)2

Equilibrium forZn + 2OH- Zn(OH)2 + 2e

-

Equilibrium forZn(OH)

2

+ 2OH- ZnO2

2- + 2H2

O

Equilibrium forZn + 4OH- ZnO2

2- + 2H2O + 2e-

Zn2+ stablein solution

Zn metalstable

Zn(OH)2stablesolid

ZnO22-

stable insolution


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Potential

7 14

2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8

-1.2

0

Pourbaix Diagram for Zinc

Corrosion is thermodynamically impossible

Corrosion is possible,but likely to be stifled by

solid corrosion productCorrosion

Immunity

Passi-vity

Corro-sion

Corrosion possible withhydrogen evolution

Corrosion possible withoxygen reduction

Corrosion requires strong oxidising agent

Corrosion requires strongoxidising agent


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2H+ + 2e- H2 hydrogen evolution in acids

2H2O + 2e- H2 + 2OH

- hydrogen evolution in water/bases

(the above two reactions are equivalent reactions)

O2 + 2H2O + 4e- 4OH- oxygen reduction in water/bases

O2 + 4H+ + 4e- 2H2O oxygen reduction in acids

(the above two reactions are equivalent reactions)


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Pourbaix Diagram for Gold

Potential

7 14

2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8

-1.2

0

Gold metal stable

Immunity region

CC

Passivity Gold cant corrodewith oxygen reductionor hydrogen evolution


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Pourbaix Diagram for Copper

Potential

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2.01.6

0.8

1.2

-0.4

0.4

0.0

-1.6

-0.8

-1.2

0

Cu metal stable

Cu2+

stablein solution

Cu oxidesstable

CuO2

2-stableinsoln.

No - hydrogen evolution

only occurs below thepotential for copper

corrosion

Will copper corrodein deaerated acid?


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Pourbaix Diagram for Copper (Cont.)

Potential

7 14

2.01.6

0.8

1.2

-0.4

0.4

0.0

-1.6

-0.8

-1.2

0

Cu metal stable

Cu2+

stablein solution

Cu oxidesstable

CuO2

2-stableinsoln.

Usually it will just passivate,but corrosion can occur inslightly acid solutions

Will copper corrodein neutral aeratedwaters?

P b i Di f I


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Pourbaix Diagram for Iron

Potential

7 14

2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8

-1.2

0

Fe metal stable

Fe3+

Fe oxidesstable

Will iron corrode indeaerated acid?

Fe2+ stable

Yes - there is a reasonablywide range of potentialswhere hydrogen can beevolved and iron dissolved


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P b i Di f I


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Pourbaix Diagram for Iron

Potential

7 14

2.0

1.6

0.8

1.2

-0.4

0.40.0

-1.6

-0.8

-1.2

0

Fe metal stable

Fe3+

Fe oxidesstable

Will iron corrode inalkaline solutions?

Fe2+ stable

No - iron forms a solid oxideat all potentials, and willpassivate

P b i Di f Al i


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Pourbaix Diagram for Aluminum

Potential

7 14

1.2

0.8

0.0

0.4

-1.2

-0.4-0.8

-2.4

-1.6

-2.0

0

Al

Al3+Al2O3

AlO2-


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Limitations of Pourbaix Diagrams

Tell us what can happen, not necessarily

what will happen

No information on rate of reaction

Can only be plotted for pure metals and

simple solutions, not for alloys


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EXAMPLE PROBLEMEXAMPLE PROBLEM((PrbPrb. 2.10 in. 2.10 in Principles and Prevention of CorrosionPrinciples and Prevention of Corrosion , Denny Jones, 1996), Denny Jones, 1996)

Using the Pourbaix diagram for nickel, give the anodic and cathodic reactions on Niin water for the following conditions, assuming activity of 10-6 for all soluble species:

(a) deaerated pH 2, (b) deaerated pH 10, (c) aerated pH 2, aerated pH 10.


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Home Work Problems

Determine whether silver will corrode with hydrogen evolution (PH2=1

atm) in deaerated KCN solution, pH=9, when CN-

activity=1.0 andAg(CN)2 activity=0.001. What is the cell potential in volts?

Ag(CN)2

- + e- Ag + 2CN- eo=-0.31 V

Prbs. 3, 5, 6 of Chapter 2

in Principles and Prevention of Corrosion, Denny Jones, Prentice-Hall, 1996.


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Home Exercise Problems

Prbs. 7, 8, 11 and 12 of Chapter 2

in Principles and Prevention of Corrosion, Denny Jones, Prentice-Hall, 1996.

R f


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References

Principles and Prevention of Corrosion, Denny Jones, Prentice-Hall, 1996.

Corrosion Engineering, Mars Fontana, McGraw-Hill, 1986.

Corrosion and Corrosion Control, H. H. Uhlig and R. W. Revie, John Wiley &

Sons, 1985.

Web Site of Dr. R. A. (Bob) Cottis. Web Site of Dr. Floyd Beckford

Documents

03 Electrochemistry Thermodynamics and Electrode Potential