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Kwame Nkrumah University of

Science & Technology, Kumasi, Ghana

Corrosion Thermodynamics

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Course OutlineWeek Topic

1

Introduction:

Reactivity types, corrosion definition, atmospheric

corrosion, classification, effects, costs, risk management

2-3

Corrosion thermodynamics:

Corrosion reactions, cell requirements, free

energy change, electrochemical potential, Nernst

equation, Eh-pH (Pourbaix) diagrams, reference

electrodes

4-6

Corrosion kinetics:

Electrical double layer, exchange current density,

activation and mass transport control, mixed potential

theory, polarization diagrams, passivity

7 Mid Semester Exams

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Why study thermodynamics?

• Provides understanding of energy changes involved in

electrochemical corrosion reactions

• Show how conditions may be adjusted to control

corrosion

– Immunity

– Passivity

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Guiding Principles

• Corrosion is only a question of time – “rust never sleeps”

• Corrosion of metals is electrochemical in nature

o Reactions involve both charge (electron) and solvated proton (H+) transfer

Fe + 2H2O = Fe(OH)2 + 2H+ + 2e

• Any reaction which can be divided into two (or more) partial

reactions of oxidation and reduction is termed electrochemical.

– Eg. Reaction of Zn in HCl

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Guiding Principles• Chemical reaction – metal (iron) either changes into rust or

dissolves

o Acid-base reactions involving only solvated proton (H+) transfer (no

electron transfer)

Fe(OH)2 = HFeO2- + H+

2Fe(OH)3 = Fe2O3.3H2O

• Corrosion rate depends on both the metal (and its condition) and

the environment.

• Cr2O3 (SS) is dynamic layer and dissolves in reducing

environment. Al2O3 is static.

• Presence of moisture is important

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• Two reactions are necessary:

-- oxidation reaction:

-- reduction reaction:

Zn Zn2 2e

2H 2e H2(gas)

• Other reduction reactions in solutions with dissolved oxygen:

-- acidic solution -- neutral or basic solution

O2 4H 4e 2H2O

O2 2H2O 4e 4(OH)

Electrochemical Nature of Corrosion

6

Zinc

Oxidation reactionZn Zn2+

2e-Acid solution

reduction reaction

H+H+

H2(gas)

H+

H+

H+

H+

H+

flow of e-

in the metal

Corrosion of zinc in an acid solution

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Corrosion Cell

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Four Prerequisites for Corrosion

1. An anode on a metal surface (oxidation rxn).

e.g. Fe = Fe2+ + 2e-

2. A cathode on a metal surface (reduction rxn).

e.g. O2 + 2H2O + 4e- = 4OH-

3. Electrolyte in contact with anode and cathode (path for ionic

conduction).

4. An electrical connection between anode and cathode (allows

electrons to flow between anode and cathode)

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Importance of the Four

Prerequisites of Corrosion

• Corrosion Control:

• Stopping the anodic reaction (e.g. passivation, cathodic

protection)

• Stopping the cathodic reaction (e.g. removing dissolved

oxygen)

• Stopping ion flow between anodes and cathodes (e.g. organic

coatings)

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Why do corrosion cells form?

• Corrosion cells form due to the difference in energy

between the metal and the environment.

• Variations could be

– Metallurgical – composition, microstructure, inclusions, heat

treatment, welding, fabrication etc.

– Environmental – temperature induced corrosion, microbial

induced corrosion, environmental induced SCC etc.

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Etching as Corrosion

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Impurity atom/compound

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Multiphase structure

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Anodic Processes• Consider similar cell reactions:

– Zn + 2HCl → ZnCl2 + H2

– Zn + H2SO4 → ZnSO4 + H2

– Fe + 2HCl → FeCl2 + H2

– 2Al + 6HCl → 2AlCl3 + 3H2

• Separation into anodic reactions:

– Zn → Zn2+ + 2e-

– Fe → Fe2+ + 2e-

– Al → Al3+ + 3e-

• Generalized anodic reaction:

– M → M z+ + ze-

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Some Key Anodic Processes

• Dissolution-Precipitation

– M → Mz+ + ze- (electrochemical reaction)

-- Mz+ + zH2O = M(OH)z↓ + zH+ (chemical reaction)

– Example: Pb + H2SO4 → PbSO4↓ + H2 (cell reaction)

• Direct Film Formation

– M + H2O = MOads + 2H+ + 2e-

– M + H2O = MO(oxide) + 2H+ + 2e-

– Example: 2Cr + 3H2O = Cr2O3(s) + 6H+ + 6e-

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Cathodic Processes• Hydrogen Evolution (Deaerated Acids):

– 2H+ + 2e- → H2

• Oxygen Reduction:

– O2 + 4H+ + 4e- → H2O (Acidic)

– O2 + 2H2O + 4e- → 4OH- (Neutral/Alkaline)

• Metal Ion Reduction:

– Fe3+ + e- → Fe2+

• Metal Deposition:

– Cu2+ + 2e- → Cu

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Types of Electrochemical Cells

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Voltaic Cell (Galvanic Cell)

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Electrolytic Cell

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Galvanic Cell vs. Electrolytic Cell

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Standard Cell Reaction Direction

1. Zn/HCl cell:

– Anode: Zn = Zn2+ + 2e; E0a = -E0c = -(-0.762V)

– Cathode: 2H+ + 2e = H2; E0c = 0V

– E0(cell) = E0(anode) + E0(cathode) = +0.762V

• Spontaneous as written

2. Cu/HCl cell:

– Anode: Cu = Cu2+ + 2e; E0a = -E0c = -(+0.342V)

– Cathode: 2H+ + 2e = H2; E0c = 0V

– E0(cell) = -0.342V

• Not spontaneous as written

Consider two corrosion cell reactions:

Zn(s) + HCl(aq) = ZnCl2(aq) + H2(g)

Cu(s) + HCl(aq) = CuCl2(aq) + H2(g)

Are the reactions spontaneous in the direction written?

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Cell Potential

• Corrosion cell is usually represented as:

Zn|Zn2+||Cu2+|Cu

• The two half-cell reactions are:

1. Zn = Zn2+ + 2e

2. Cu2+ + 2e = Cu

• The overall reaction is:

Zn + Cu2+ = Zn2+ + Cu

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Derivation of the Nernst Equation

• The free energy change, ΔG, is related to the cell potential by Faraday’s law:

nFEG

00 nFEG

were n = number of electrons transferred in corrosion

reaction

F = Faraday’s constant, 96,500 C/mole

E = cell potential in a given state (Volts)

ΔG = Joules

Under standard conditions:

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The Nernst Equation

• Consider the following reaction:

A + B = C + D

}]][[

]][[ln{0

BA

DCRTGG

}]][[

]][[ln{0

BA

DCRTnFEnFE

}]][[

]][[ln{0

BA

DC

nF

RTEE

One of the most

fundamental

equations in corrosion

engineering!

KRTGG ln0

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The Nernst Equation

• Under standard conditions: T = 298K, R = 8.3143 J(mol.K)-1

• Nernst equation becomes

}]][[

]][[log{

059.00

BA

DC

nEE

E = non equilibrium potential

Standard EMF Series

metalo

• Metal with smaller

V corrodes.

• EMF series

Au

Cu

Pb

Sn

Ni

Co

Cd

Fe

Cr

Zn

Al

Mg

Na

K

+1.420 V

+0.340

- 0.126

- 0.136

- 0.250

- 0.277

- 0.403

- 0.440

- 0.744

- 0.763

- 1.662

- 2.363

- 2.714

- 2.924

metal Vmetalo

mo

re a

no

dic

mo

re c

ath

od

ic

V =

0.153V

o-

1.0 M

Ni2+ solution

1.0 M

Cd2+ solution

+

25°C NiCd

Cdo

Nio

• Ex: Cd-Ni cell

V < V Cd corrodes

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Solution concentration and temperature

• Ex: Cd-Ni cell with

standard 1 M solutions

VNi

o VCd

o 0.153 V

-

Ni

1.0 M

Ni2+ solution

1.0 M

Cd2+ solution

+

Cd 25°C

• Ex: Cd-Ni cell with

non-standard solutions

Y

Xln

nF

RTVVVV o

Cd

o

NiCdNi

n = #e-

per unit

oxid/red

reaction

(= 2 here)F =

Faraday's

constant

= 96,500

C/mol.

- +

Ni

Y M

Ni2+ solution

X M

Cd2+ solution

Cd T

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∆G = -nF∆E www.knust.edu.gh

The Nernst Equation

Question

What is the thermodynamic tendency for tin (Sn) to

corrode in deaerated sulphuric acid (H2SO4) at pH = 2,

activity of Sn = 10-6, pH2 = 1.0 atm, at 25 °C?

E0 = -0.138 V

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Importance of Potential

• Interfacial potential (E)

– Potential of corroding metal minus potential in

electrolyte (next to metal surface)

• Importance

– Potential can be measured readily

– Potential affects rate of corrosion

• Require reference electrode for measurement

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Measurement of Cell Potential

M

Ref

V

ElectrolytePorous tip

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Common Reference Electrodes

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Conversion

between

reference

electrodes

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Hydrogen Evolution Reaction

• Standard State, ec° = 0 V

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Neutral or Alkaline Environment

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Half Cell Potential = F(pH)

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Oxygen Evolution Reaction

• Standard State, ec° = +1.229 V

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Neutral or Alkaline Environment

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Half Cell Potential = F(pH)

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Stability of H2O