INTRODUCTION
THERMODYNAMICS OF CORROSION
KINETICS OF CORROSION
GALVANIC CORROSION
CORROSION PROTECTION

Corrosion – from Latin “to gnaw”
Corrodere – “to gnaw to pieces”
Corrosion is the degradation of a metal by electrochemical reaction with its environment.
It was calculated that in the UK, 1 ton of steel is converted completely to rust every 90 s
We will mainly consider corrosion of metals in aqueous environment .

Generally, when oxidation of a metal occurs the product formed could be:
 a soluble metal ion or complex, or
 an insoluble oxide, hydroxide or other salt.
Common oxidising agents:
Factors affecting corrosion:
 presence of O
2
 presence of complexing agent
 pH

Reaction of most metals with O
2
 thermodynamically favourable

Some form a protecting oxide layer (passive layer) e.g.
Al
 very reactive toward O
2
 oxide layer very thin and very protecting
Ti
 non-corrodable due to oxide layer formed
(also resistant to sea water and Cl
2
)
Titanium hip prosthesis
Stainless steel: steel is made corrosion resistant by alloying with Cr
 forms Cr
2
O
3 layer

Layer is too thin to be visible
 metal remains lustrous.

Layer cannot be penetrated by water and air,
 metal beneath is protected.

Layer quickly reforms when the surface is scratched.
Chrysler Building - type 302 stainless steel (chromium-nickel alloy)

Reaction of most metals with O
2
 thermodynamically favourable

Some have slow reaction kinetics
Metals such as Zn, Mg, Cd corrode slowly even though

G < 0
Galvanised metal sheeting
Graphite releases large amounts of energy upon oxidation, but the process is so slow kinetics that it is effectively immune to electrochemical corrosion under normal conditions.

Why don’t precious metals corrode????
 e.g. Au, Pt
Au + 3 /
2
H
2
O + 3 /
4
O
2

Au(OH)
3
Au nuggets
That is why they can be found in metallic form on Earth, and it is a large part of their intrinsic value.
Pt nuggets
Au ore body

Iron objects were found remarkably preserved after centuries of immersion at the bottom of a peat bog. Why???

Is the corrosion of copper in an acidic solution spontaneous? Always?
Consider:
Copper metal is in contact with a 1 M acid solution containing 10 -6 M Cu 2+ .
Calculate the equilibrium potential for this solution:
E

(Cu 2+ /Cu) = +0.34 V (vs SHE)

Cu 2+ (aq) + 2e  Cu(s) E

= +0.34 V
E

E o 
RT ln Q nF

E

E o 
0 .
05916 log Q n

In an aerated 1 M acid solution:
O
2
+ 4H + + 4e  2H
2
O
Cu 2+ + 2e  Cu
Overall:
2Cu + O
2
+ 4H +  2Cu 2+ + 2H
2
O
E

= 1.23 V
E = 0.16 V
In a deaerated 1 M acid solution:
2H + + 2e  H
2
Cu 2+ + 2e  Cu
Overall:
Cu + 2H +  Cu 2+ + H
2
E

= 0 V
E = 0.16 V
Is the corrosion of copper in an acidic solution spontaneous? Always?

Pourbaix diagram: for copper in a non-complexing aqueous soln at 25

C
Pourbaix diagrams give info about thermodynamics only
Kinetic factors may predominate in many situations

What info can be found on a Pourbaix diagram?
 Potentials for redox couples as a function of pH e.g. M/M n+ and M n+ /M (n+1)+
 Most stable metal compounds as a function of pH
 predict corrosion products
 Zones where metal would corrode or not corrode or become passive
Passivation
 dissolution occurs only to a point such that a maximum of 10 -6 M is in solution
In these diagrams we get 4 types of lines:
1) horizontal
2) vertical
3) sloping
4) dashed

Vertical lines:
Equilibria involving hydrolysis, but
NOT e transfer e.g. Cu 2+ + H
2
O  CuO(s) + 2H +
At pH 7: Cu 2+ concentration is reduced below 10 -6 M
 passivation.
Above pH 7, Cu 2+ will not be the major corrosion product.
Horizontal lines:
Equilibria involving e transfer, but
NOT H + /OH e.g. Cu 2+ + 2e 
Cu(s)
Between pH -2 to 6 Cu dissolves for potentials ~0.16 V.
Sloping lines:
Equilibria involving both hydrolysis and e transfer e.g. 2Cu(s) + 2H
2
O

Cu
2
O(s) + 2H + + 2e pH 6-14: corrosion product may be Cu
2
O, but this may oxidise further.
pH > 7: Cu 2+ will not be the major corrosion product if other oxidising agents are present.

slope

0 .
05916

H
 n
V per pH unit at 25 o
C
B
O
2
+ 2H
2
O + 4e  2OH E

= 0.4 V
O
2
+ 4H + + 4e  2H
2
O E

= 1.23 V
A H
2
O is stable in the region between the lines
2H + + 2e  H
2
E

= 0 V
2H
2
O + 2e 
H
2
+ 2OH E

= -0.83 V
Dashed lines:
Equilibria involving the redox couples A = H + /H
2 function of pH and B = H
2
O/O
2 as a
Slope = 0.059 V per pH unit.
If the dashed line is above the solid line, the corrosion reaction obtained by adding the two equilibria will be spontaneous .
If the dashed line is below the solid line, the corrosion process is thermodynamically unfavourable and the metal is immune to corrosion .

Corrosion potential:
- the potential of the metal surface in contact with electrolyte where corrosion occurs.
- no net current flows at the corrosion potential.
Oxidation = corrosion of metal i a i c
Reduction of substance in contact with the metal
Corrosion current:
- the exchange current at the corrosion potential.

How is the rate of corrosion determined?

Measure steady state current for metal oxidation and H
2 function of potential.

Plot graph of log
 i
 vs E
 a Tafel plot

Extrapolate lines till they overlap i.e. log
 i a

= log

-i c

= log
 i corrosion
 evolution as a
Change in i o or the Tafel slope
 change in corrosion rate

Galvanic corrosion: The electrochemical process in which one metal corrodes preferentially when it is in contact with a different type of metal and both metals are in an electrolyte .
Cu 2+ (aq) + 2e  Cu(s)
Fe 2+ (aq) + 2e  Fe(s)
Zn 2+ (aq) + 2e  Zn(s)
E

= +0.34 V
E

= -0.44 V
E

= -0.76 V
When different types of metal come into contact in the presence of an electrolyte a galvanic couple is set up as different metals have different electrode potentials .
The electrolyte provides a means for ion migration from the anode to the cathode.

 The anodic metal corrodes faster than it would otherwise.
 Corrosion of the cathodic metal is retarded even to the point of stopping.
 The presence of electrolyte and a conducting path between the metals may cause corrosion where otherwise neither metal alone would have corroded.
Factors that influence galvanic corrosion:
Relative size of anode and cathode
Degree of electrical contact
Aeration of electrolyte
Electrical resistance of electrolyte
Type or concentration of electrolyte
Temperature
Humidity
Potential difference between the two metals
Oxide formation
Covering by bio-organisms

1) CATHODIC PROTECTION
The potential of the metal is shifted more negative
 lower oxidation rate .
i) Electrolysis
Surround metal to be protected by inert anodes and pass a current (i cath
) between the metal and anodes.
ii) Sacrificial anode
Another metal with a more negative
 m potential is place in good electrical contact with the metal to be protected.
log log log log
Rate of metal dissolution reduced from i corrosion to i protected
.
Sacrificial metal will enforce its corrosion potential on the metal surface

i) Electrolysis ii) Sacrificial anode log log log log
Problem:
H
2 evolution also increases.
Some metals absorb this hydrogen at grain boundaries or into the metal lattice
 can change metal structure and hence chemical and physical properties of metal

Leads to hydrogen embrittlement

Example of sacrificial anodes used in cathodic protection:
Al anodes mounted on a steel jacket structure
Common sacrificial anodes: Zn, Mg, Al
Zn
Mg
Al
E o /V
-0.76
-2.36
-1.66
Sacrificial anodes will corrode at a higher rate than protected metal
 anodes need to be replaced periodically

2) ANODIC PROTECTION
The potential of the metal is shifted more positive to a region where it is passivated .
The thin layer of corrosion product on metal surface can act as a barrier to further oxidation of the metal.
Achieve passivation by: i) Electrochemical means
Surround metal by cathodes and apply a potential (e.g. anodisation of Al) ii) Chemical means
Add an oxidising agent to the solution (e.g. dichromate) OR add an alloying element to the metal which act as small local cathodes which can lead to film formation (e.g. Cr to stainless steel)

3) MEDIUM MODIFICATION

Useful for closed systems i) Remove aggressive species from medium to reduce corrosion e.g. O
2
, acid, number of ions in the electrolyte ii) Add inhibitors
 to catalyse passive film formation
 to act as redox reagents
 shifts metal potential to regions where metal is anodically or cathodically protected
 to adsorb on to metal surface to decreases rate of anodic and/or cathodic reaction
 adsorption must occur close to the corrosion potential
4) SURFACE COATINGS
Reduce rate of corrosion by “removing” metal from the environment .
Examples of surface coatings:

Plating with others metals which corrode more slowly

Forming oxide films

Coating with organic polymers (e.g. paint)
Localised damage to coating could lead to rapid corrosion in that region

Self-study!