CH.CH.22. . TTHERMODYNAMICSHERMODYNAMICS OFOF CCORROSIONORROSION

1Image Source: Corrosion Doctors, www.corrosion-doctors.org

Why Do Metal Corrode?Why Do Metal Corrode?} Any spontaneous reaction in the universe is associated with a lowering in the free

energy of the system, i.e., a negative free energy change.} All metals except some noble metals have free energies greater than those of their

yy

} All metals except some noble metals have free energies greater than those of their compounds. Therefore, they tend to become their compounds through the process of corrosion.

} As an example, take the following examples:

kJ/mol119GCu(OH)O1

OHCu

kJ/mol596G,2Mg(OH)2O21

O2HMg

=++

J/molk66G,3Au(OH)2O43

O2H23

Au

kJ/mol119G,2Cu(OH)2O2O2HCu

+=++

=++

} It is obvious from G that Mg and Cu will corrode.} Au has a positive G and does not corrode.

T d t i M C A

2

} Tendency to corrosion: Mg > Cu >> Au

Electrochemical Nature of CorrosionElectrochemical Nature of Corrosion

} All metallic corrosion are electrochemical reactions, i.e., metal is converted to its ions and compounds with a transfer of electrons.converted to its ions and compounds with a transfer of electrons.

} The overall reaction may be split into oxidation (anodic) and reduction (cathodic) partial reactionsreduction (cathodic) partial reactions.

} Electrochemical reactions take place in electrochemical cell.

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Electrochemical CellElectrochemical Cell

Metallic PathMetallic Path

ee--

A C+ ions

A C+ ions

A C- ions

Electrolytic Path

A C ions

Electrolytic PathElectrolytic Path

Conventional Current Flow

Electrolytic Path

Conventional Current Flow

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Conventional Current FlowConventional Current Flow

Electrochemical CellsElectrochemical Cells

Two types:

1. Electrolytic Cells - these are cells in which an external electrical source forces a non-spontaneous reaction to occur.

(One common process is called electrolysis.)(not to be covered in this course)

2. Galvanic Cells also called voltaic cells. In these cells, ,spontaneous chemical reactions generate electrical energy and supply it to an external circuit.

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(e.g., corrosion, battery and fuel cell, etc.)

2e gainedper Cu2+ ionreduced2e

lostper Zn atom

e

Cu2+ eCuZn2+Zn

Voltmetere e

oxidized

Salt bridgeNa+Zn CuSO42

Anode()

Cathode(+)

Zn2+ Cu2+

Oxidation half-reaction

Reduction half-reaction

Zn(s)

Zn Cu

Zn2+(aq) + 2e

Overall (cell) reactionZn(s) + Cu2+(aq)

Cu2+(aq) + 2e Cu(s)

Zn2+(aq) + Cu(s)Image Source: M.S. Silberberg, Chemistry:

The Molecular Nature of Matter and Change, 2nd ed.

Cell NotationCell Notation

Zn (s) | Zn2+ (aq 1 0M) || Cu2+ (aq 1 0M) | Cu (s)Zn (s) | Zn (aq, 1.0M) || Cu (aq, 1.0M) | Cu (s)

C ll N i} Cell Notation Write components in sequence Separate phases with a single vertical line | A salt bridge or membrane is represented by a double vertical line || Included a specification of the species concentration

} In a galvanic cell, voltage/potential difference between the electrodes drops to 0 as the reaction proceed (to be covered in Chapter 3).

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Corrosion ProcessCorrosion ProcessConsider what happens when a metal electrode M is dipped into a solution containing the metal ions Mn+. The concentration of Mn+ in the metal is much larger than in the

l ti hi h t h i l d i i f th t t d t li th t tisolution, which creates a chemical driving force that tends to equalize the concentration. A charge separation occurs as a result of the metals tendency to oxidize.

M Mn+ (s) + ne- (M) : Dynamic equilibrium.M M (s) ne (M) : Dynamic equilibrium.

e-+

++++

++++

++

CuZn++

+

++++

++++++ + ++

++

++

WATERM t l

8

water+

WATER

Forces of ionization

Metal

Image source: Prof. H.S. Kwons Lecture Note (corrosion.kaist.re.kr)

Solvation Shells (Hydration Sheaths) of a CationSolvation Shells (Hydration Sheaths) of a Cation

Schematic of the primary and secondary solvent molecules for a cation in water.

Primary shell with completely oriented water moves as and where the ion moves. Secondary shell with partly-oriented water is a region of disturbance

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Secondary shell with partly oriented water is a region of disturbance..

Structure of Electrified InterfacesStructure of Electrified InterfacesThe metal electrode - solution interface is also hydrated because of the charge effects.

Schematic of a charged interface and the location of cations at the electrode surface.

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Structure of Electrified InterfacesStructure of Electrified InterfacesBecause of adsorbed H2O on surface and primary hydration shell, ions can only approach to within a fixed distance from the electrode: the OUTER HELMHOLTZ PLANE (OHP). The resulting double layer of charge is often treated theoretically as a capacitor..

--

-+++

--

---

Scattered ions

++ -

-- -

--

'Stuck' ions

M S

M

x

S

St d l f EDL

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Simplified double layer at a metal-solution interface.

Stern model for EDL

Structure of Electrified InterfacesStructure of Electrified InterfacesThe electric double layer (EDL; ) at the surface acts as a barrier to electron transfer.

A potential drop develops at the metal-solution interface(At equilibrium, this potential drop is representative of the reversible (electrode) potential )potential.)

This causes the surface to polarize, i.e., the potential has to change in order for the electrons to overcome the barrier.

This is called ACTIVATION POLARIZATION (CH.3).

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Corrosion ThermodynamicsCorrosion ThermodynamicsThe free energy for any species, G:

G = G0 + RT ln a

yy

Consider the generic chemical reaction:

aA + bB = cC + dD

The free energy change, G, which is a measure of driving force for the reaction to proceed, is given by the difference of the free energy of the products and reactants.

++==

BAdC

tstanreacproducts

)bGaG()dGcG(

GGG

+=

++++++=

bB

aA

dD

cC0

B0BA

0AD

0DC

0C

BAdC

aaaa

lnRTG

)alnbRTbGalnaRTaG()alndRTdGalncRTcG(

)()(

In generalized form:

+=

ktstanreac

jproducts0

)a()a(

lnRTGG

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IF G

Gibbs Function and WorkGibbs Function and Work} Start with the First Law of Thermodynamics and some standard thermodynamic

relations. We find

Gibbs Function and WorkGibbs Function and Work

dU = dq + dwdq = T dSdw = PdV + dwelectricalelectrical

dHP = dUP + PdVdU = T dS PdV + dwelectrical

P PdGT = dHT T dS

= dUT ,P + PdV T dS= T dS PdV + dwelectrical + PdV T dS

dGT ,P = dwelectrical

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Thus the Gibbs function is at the heart of electrochemistry, for it identifies the amount of work we can extract electrically from a system.

Free Energy and Electrode PotentialFree Energy and Electrode Potential} Here we can easily see how this Gibbs function relates to a

potential.

gygy

welectrical = V Qsince Q = n F

} By convention we identify work which is being done by the system

= n F E

} By convention, we identify work which is being done by the system on the surroundings has a negative sign. And negative free energy change is identified as defining a spontaneous process.

GT ,P = welectrical = n FE

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} Note how a measurement of a cell potential directly calculates the Gibbs free energy change for the process.

Free Energy and Electrode PotentialFree Energy and Electrode Potential

} Associated with each electrochemical (galvanic) cell is a potential difference between the electrodes called the equilibrium cell potential,

( l ll d ibl l h i l i l l i

gygy

Ecell (also called reversible or electrochemical potential or electromotive force emf)

E = E E = (e +e )Ecell = Ecathode Eanode = (ea+ec)

} The electrodes are considered not connected to each other so no reaction takes place on them yet, i.e., each electrode is at equilibriumreaction takes place on them yet, i.e., each electrode is at equilibriumwith its environment.

} E measures the spontaneity of the cells redox reaction

G = -nFEcell

} Higher cell potentials (more negative G) indicate a great driving force f th ti itt (i f d di ti i f l ft t i ht)

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for the reaction as written (in forward direction, i.e., from left to right).

HalfHalf--cell potentialcell potential

} Synonyms (Single) electrode potential

pp

(Single) electrode potential Redox (reduction/oxidation) potential

Cathode

Cathode Potential DifferenceElectrolyte

Ecell

A d P t ti l Diff

Cat ode ote t al e e ce

Anode

Anode Potential Difference

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Standard state298K, at unity activity

Selection of hydrogen evolution reaction as having a standard half-cell potential of 0 000Vcell potential of 0.000V.

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HalfHalf--cell Potentialcell Potential

We do not have a means of measuring the potential difference between the electrode and the solution.

?

Cathode

?

Electrolyte?

S l Ecell

Anode

Solution:

Adopt an arbitrary reference electrode.

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?

HalfHalf--cell (electrode) Potentialcell (electrode) Potential( )( )

Measurement of voltage with electrical circuitMeasurement of voltage with electrical circuitConnection for an electrochemical measurement

Source: P.J. Moran, E. Gileadi. J. Chem. Educ. 66 (1989) 912.

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Inclusion of second interface allows combined voltage of the two interfaces to be measured with a voltmeter

Standard Hydrogen ElectrodeStandard Hydrogen Electrodey gy g

} The convention is to select a particular electrode and assign its d d d i i l h l f 0 0000V Thi l d istandard reduction potential the value of 0.0000V. This electrode is

the Standard Hydrogen Electrode (SHE).

2H+(aq) + 2e H2(g)H2

V000.0e0H/H 2=+

H2

PtThe standard aspect to this cell is that the activity of H2(g) and that of H+(aq) are both 1. This means that the pressure of H2 is 1 atm

d h i f H i 1M i h

H+

and the concentration of H+ is 1M, given that these are our standard reference states.

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H+

HalfHalf--Cell Potential MeasurementCell Potential Measurement

SHE0

Zn/ZnV762.0e 2 =+

} Zn and Pt are at equilibrium with their environments at standard conditions.} The measured potential is the equilibrium potential at standard conditions

vs SHE

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vs. SHE.

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Standard Electrode PotentialStandard Electrode Potential

} For convenience, all half-cell reactions are written in the table in reduction form. Zn=Zn2++2e- and Zn2++2e-=Zn are identical and represent zinc in

equilibrium with its ions with a potential of -0.762V vs. SHE.

Th iti th 0 l f h lf ll ti th t} The more positive the e0 value for a half-cell reaction the greater the tendency for the reaction to proceed as written (in reduction-cathodic form) at standard conditions.

} The more negative the e0 value, the more likely is the reverse of the reaction as written (in oxidation-anodic form) at standard conditions.

} However, usually concentrations of reactants differ from one another and also change during a reaction, in that case E0cell (or e0half-cell) and

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the actual Ecell (ehalf-cell) are related by the Nernst Equation.

Nernst EquationNernst EquationConsider the following reaction involving electron transfer

b d

qq

aA + bB + cC + dD +

nFEaaaa

lnRTGG bB

aA

dD

cC0 =

+=

And for the following half-cell reaction in reduction form

KlognF

RT303.2E

aaaa

lognF

RT303.2EKln

nFRT

Eaaaa

lnnFRT

EE 0bB

aA

dD

cC00

bB

aA

dD

cC0 =

==

=

g

aA + bB + + ne- cC + dD +

dcdc aaRT3032aaRT

O1 + O2 + + ne- R1 + R2 +

=

= b

BaA

DC0bB

aA

DC0

aaaa

lognF

RT303.2e

aaaa

lnnFRT

ee

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[ ][ ]

+=

RO

lnnFRT

ee 0

ActivityActivity

} Activity of a dissolved species A (aA or (A) in textbook) is equal to its concentration in moles per 1000 grams of water (molality) multiplied

yy

concentration in moles per 1000 grams of water (molality) multiplied by the activity coefficient, f.

} Activity coefficients are extensively tabulated in numerous chemical and electrochemical handbooks.a d e ect oc e ca a dboo s

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

} The activities of pure solids and water are set equal to unity in aqueous solutions.

} At 25C 2 303RT/F 0 0592 V equivalent} At 25 C, 2.303RT/F 0.0592 Vequivalent

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Activity CoefficientActivity Coefficientyy

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Source: Corrosion and Corrosion Control, H.H. Uhlig and R.W. Revie, John Wiley and Sons, 1985.

Prediction of SpontaneityPrediction of Spontaneity

1. First, write the half-cell reaction with the more positive (less negative) e0 for the reduction-cathodic along with its half-cell

p yp y

negative) e for the reduction cathodic along with its half cell electrode potential.

2. Write the other half-cell reaction as an oxidation-anodic and include its half-cell electrode potentialts a ce e ect ode pote t a

3. Balance the electron transfer4. Obtain the cell reaction by adding the reduction and oxidation half-

cell reactionscell reactions5. Determine the overall cell potential, Ecell from

EEcell = ec - ea

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Example Example 11(Notation: e for the half(Notation: e for the half cell rxn E for the overall cell rxn )cell rxn E for the overall cell rxn )

For the reaction 2H+ + 2e- = H2 at 25oC and 1 atm.

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

220

220

)H()H(

log2

0592.0e

)H()H(

lognF

RT303.2ee ++ ==

2H

)H(

Plog

20592.0

0e 2+=

[ ]205920 [ ]2H )Hlog(Plog20592.0e 2 +=[ ])Hlog(2Plog0592.0e += [ ])Hlog(2Plog

2e

2H=

[ ]pH2Plog2

0592.0e

2H+= )Hlog(pH +=

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2

pH0592.0e = when PH2 = 1 atm

Example 2Example 2Lets consider the emf (E) of the following Cu-Zn cell at 250C.

pp

Zn Cu

Zn2+ Cu2+

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Example 2Example 2Reduction reaction for Cu: Cu2+ + 2e- Cu e0 = 0.337 V

and

pp

)Culog(0592.0

3370e 2++=and

Reduction reaction for Zn: Zn2+ + 2e- Zn e0 = -0.763 V

and

)Culog(2

337.0e +=

)Znlog(2

0592.0763.0e 2++=

Lets assume that copper corrodes (i.e., Cu is anode, Zn is cathode).

Overall reaction = [reduction rxn at cathode] [oxidation rxn at anode]

2

Cu + Zn2+ Cu2+ + Zn

emf (E) of the cell is then;

)Zn(05920 2+

Lets assume that the activities of Cu & Zn to be the same, then the emf, E = -1.100V (vs. SHE).

Negative E (i.e., positive G) means that the reaction is not spontaneous as written, but goes instead in the opposite direction

)Cu()Zn(

log2

0592.0100.1eeeeE 2CuZnanodecathode ++===

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in the opposite direction.In other words, the Zn electrode corrodes (anode), and Cu2+ plated out on the Cu electrode (cathode).

Example Example 3 3 ((PrbPrb. . 22--33a)a)Assuming standard states for all reactants and products, determine the spontaneous direction of the following reactions by calculating the cell potential:

pp (( ))

Cu + 2HCl = CuCl2 + H2

Anodic Cu2+ + 2e- = Cu ea = ea0 = -0.342 V vs. SHECathodic 2H+ + 2e- = H2 ec = ea0 = 0.000 V vs. SHE

E = ec ea = 0.342 V > 0

The reaction is spontaneous as written ().

NOTEA spontaneous reaction does not necessarily mean a fast reaction !!!

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A spontaneous reaction does not necessarily mean a fast reaction !!!

A Note for The Sign NotationA Note for The Sign Notation

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

gg

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 anode half-cell electrode potential.

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Secondary Reference ElectrodesSecondary Reference Electrodes

} If a reference electrode other than SHE is used to measure the equilibrium potential of a rxn, the potential of the reference

yy

equilibrium potential of a rxn, the potential of the reference electrode relative to SHE should be added to the measured potential if one wants to determine the equilibrium potential of the reaction relative to SHE.

} The reference electrodes are also used to measure the corrosion potential of a corrosion cell (Ecorr) (to be covered later).p ( corr) ( )

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Potential Values for Common Secondary Reference Electrodes Potential Values for Common Secondary Reference Electrodes (Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)

Image source: A J Bard L R Faulkner Electrochemical Methods 2nd ed John Wiley & Sons Inc 2001

Name Half-Cell Reaction Potential, V vs. SHE

Mercury-Mercurous Sulfate HgSO4 + 2e- = Hg + SO42- +0.615

Image source: A.J. Bard, L.R. Faulkner, Electrochemical Methods, 2nd ed., John Wiley & Sons, Inc., 2001.

Copper-Copper Sulfate CuSO4 + 2e- = Cu + SO42- +0.318

Saturated Calomel Hg2Cl2 + 2e- = 2Hg + 2Cl- +0.241

Silver Silver Chloride (saturated) AgCl + e- = Ag + Cl- +0 222

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Silver-Silver Chloride (saturated) AgCl + e = Ag + Cl +0.222

Standard Hydrogen 2H+ + 2e- = H2 +0.000

Reference Electrode Conversion ScaleReference Electrode Conversion Scale

Image source: R.G. Kelly et al., Electrochemical Techniques in Corrosion Science and Engineering, Marcel Dekker Inc., 2002.

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Example ProblemsExample Problems

A corrosion potential of -0.229 V versus SCE was measured for a corroding metal. What is the potential versus (a) SHE, (b) Ag/AgCl (saturated), (c)

pp

Cu/saturated CuSO4?

(a) Ecorr vs. ref 1 = Ecorr vs. ref 2 + eref 2 vs. ref 1Ecorr = Ecorr vs SCE + eSCE vs. SHEeSCE = 0.241 V vs. SHEEcorr = -0.229 + 0.241 = 0.012 V vs. SHE

(b) & (c)Home Exercise

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Homework ProblemsHomework Problems

} Problems 3, 5, 6, 12 of Chapter 2 in textbook.

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