Day 1 Thermodynamics -Pourbaix_2

8/20/2019 Day 1 Thermodynamics -Pourbaix_2

1/19

To be a partner of choice in corrosion research

Thermodynamics-Pourbaix

11/2/2012 1

Corrosion Electrochemistry :

Thermodynamics


2/19

*Dr.M Pourbaix (1966)

E l e c

t r o d e p o t e n t

i a l , V

pH


3/19

E-pH Diagram: Represent states of lowest free energy

E.g: Fe-H2O system:

predicting

environment

changes that will

prevent or reducecorrosive attacks

large region

labelled Fe in the

potential-pH diagramfor Fe-H2O system

indicates that iron is

inert under these

conditions.


4/19

Use of the Pourbaix

Diagrams:

Predicting thespontaneous direction of

reactions

Estimating the

composition of corrosion

products

Predicting

environment changesthat will prevent or

reduce corrosive attacks

Limitation:

1. Tell us what can

happen, notnecessarily what will

happen

2. Cannot predict

corrosion rates. 3. Can only be plotted

for pure metals and

simple solutions,

NOT for alloys.


5/19

The diagramme shows conditions of solution

oxidizing power (potential) and acidity of alkalinity

(pH) for the various possible phases that are

stable in aqueous electrochemical system.

The most common reduction reactions are the

reduction reactions of dissolved oxygen and the

reduction of water ( H2 evolution)


6/19

2H+

+ 2e-

= H2 (acidic)2H2O + 2e- = H2 + 2OH

- ( neutral/alkaline)

The Nernst equation gives

For 1 atm. hydrogen gas

E = -0.059 pH vs SHE

pH059.0Hlog

2

059.0

HHlog

2303.2

2

2

2+

E

F

RT E = E

0

0


7/19

Aerated Acid O2 + 4 H

+

+ 4 e

-

= 2 H2OAcid

O2 + 2 H2O + 4 e- = 4 OH- Neutral

/Alkaline

V23.1

SHEvs pH059.0Olog4

059.0

0

2

0

E

E E

E = 1.23 - 0.059pH


8/19

Neutral/Basic O2 + 2 H2O + 4 e- = 4 OH-

4

2

220

4

2220

OHOHOlog

4303.2

OH

OHOln

4

F

RT E

F

RT E E

But [H2O]=1, and log([OH-])=-14+pH

56 pH4Olog4

303.2 therefore

2

0 F

RT E E


9/19

P

o t e n t i a l ( V v s .

S H E )

pH

Upper: water can beoxidized

and form O2:

O2 can be reduced below line b

Lower: water can be

reduced to form H2

Intermediate: wateris

thermodynamically

st le

0

_a

b

02

+4H++4e=2H2

O

2H++2e=H2

0H-+H+=H2O

O2

H2

1.23 _

0 1

6

*(a) and (b) are commonly being superimposed on Pourbaixdiagrams.


10/19

When metal surface is at a potential where one orboth of reduction reactions can occur, the

possibility of corrosion exists provided the metal

disssolution reaction is thermodynamically

favourable.


11/19

Reaction of pure charge transfer.•Only electrons involved & NO hydrogen ion

•e.g: Ni Ni2+ + 2e

]log[03.025.0

ln

2

2/

2

Ni

Ni

Ni

nF

RT E ENi

o

Ni

So potential depends on the activity

of [Ni2+]not

pH.

Select 4 activities: 1, 10-2, 10-4, 10-6 M:

then

E=(-0.25V), (-0.23+0.03log(10-2)=-0.31V,

E=(-0.37V), (-0.43V)

-0.25-

0.31-

0.37-

0.43

V

pH


12/19

Reaction involving both electrons & hydrogen ion•e.g: NiO + 2H++2e Ni + H2O; E

o=0.11V

pH0.0590.11Ethen

]log[H pHsince

]log[0295.0

]][[

]][[log2

059.0

2

22

H E E

H NiO

O H Ni

E E

o

o

0.11

NiO

0.05

V

pH

Ni

0


13/19

A rod of Ni is immersed in an aqueous deaerated acid solution with a pH of 1 that contains10-4 g-ion/L of Ni2+ions. The system is under 1 atm pressure. Will the nickel corrode? What

will happen at pH more than 8? Refer to the Pourbaix Diagram for nickel.

At the metallic Ni/water interface:Ni2++2e NiENi = -0.25 + 0.026log[Ni

2+]= - 0.37VDeaerated acid solution:

2H+

+2e=H2 EH = -0.059 pHEH = -0.06V at pH 1Since ENi more active than EH , sothe electron flow from Ni (-ve) to H2 (+ve). Nickel not stable at low pH inwater, so the corrosion occurs.

8

0 _

-

0.4 _

1

Ni2+

Ni

NiO2

Ni(OH)2

pH

E, V

At pH 6 and 8:Hydrogen more active than Niregion of immunity of NiNo corrosion


14/19

P o t e n t i a l

7 14

2.0

1.6

0.8

1.2

-0.4

0.4

0.0

-1.6

-0.8

-1.2

0

Cu metal stable

Cu2+ stablein solution

Cu oxides

stable

C u O 2

2 - s t a b l e i n s o l n .Will copper

corrode in

acid?

No - hydrogen

evolution only

occurs below the

potential for coppercorrosion

Will copper

corrode in

neutral waters? Usually it will just

passivate, but

corrosion can occurin slightly acid

solutions


15/19

Reversible potential for copper metal oxidation

reaction is above line “a” for all pH. Thus in theabsence of O2, metallic copper is

thermodymanically stable in pure water.


16/19

P o t e n t i a l

7 14

2.0

1.6

0.8

1.2

-0.4

0.4

0.0

-1.6

-0.8-1.2

0

Fe metal stable

Fe3+

Fe oxides

stable

Will iron

corrode in

acid?

Fe2+ stable

Yes - there is a

reasonably wide

range of potentialswhere hydrogen

can be evolved and

iron dissolved

Will iron

corrode in

neutral waters? Yes - although iron can

form an oxide in neutral

solution, it tends not toform directly on the

metal, as the potential

is too low.

Will iron corrode

in alkaline

solution?

No - iron forms a solid

oxide at all potentials,

and will passivate


17/19

E = -0.2 V-SHE ,pH =4Increase pH : Water

Treatment

Change potential:

Cathodic Protection

Anodic Protection

Extending passive

region: Alloying : SS

Passivating inhibitor

such as chromate ions


18/19

P o t e n t i a l

7 14

2.0

1.6

0.8

1.2

-0.4

0.4

0.0

-1.6

-0.8-1.2

0

Gold metal stable

Immunity

C

CPassivity

Gold can’t corrode

with oxygen reduction

or hydrogen evolution


19/19

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Day 1 Thermodynamics -Pourbaix_2