CORROSIONCORROSION

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

INTRODUCTIONINTRODUCTION

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 O2

presence of complexing agent

pH

Reaction of most metals with O2 thermodynamically favourable

Some form a protecting oxide layer (passive layer)

Al very reactive toward O2

oxide layer very thin and very protecting

Ti non-corrodable due to oxide layer formed

(also resistant to sea water and Cl2)

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.

Stainless steel: steel is made corrosion resistant by alloying with Cr

forms Cr2O3 layer

Chrysler Building - type 302 stainless steel (chromium-nickel alloy)

e.g.

Titanium hip prosthesis

Reaction of most metals with O2 thermodynamically favourable

Some have slow reaction kinetics

Metals such as Zn, Mg, Cd corrode slowly even though G < 0

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.

Galvanised metal sheeting

That is why they can be found in metallic form on Earth, and it is a large part of their intrinsic value.

Why don’t precious metals corrode???? e.g. Au, Pt

Au + 3/2H2O + 3/4O2 Au(OH)3

Au nuggets Pt nuggets

Au ore body

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

Consider:

Copper metal is in contact with a 1 M acid solution containing 10-6 M Cu2+.

E(Cu2+/Cu) = +0.34 V (vs SHE)

Calculate the equilibrium potential for this solution:

Cu2+(aq) + 2e- Cu(s) E = +0.34 V

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RTEE o Qlog

n

05916.0EE o

THERMODYNAMICS OF CORROSIONTHERMODYNAMICS OF CORROSION

Is the corrosion of copper in an acidic solution spontaneous? Always?Is the corrosion of copper in an acidic solution spontaneous? Always?

In an aerated 1 M acid solution:

O2 + 4H+ + 4e- 2H2O E = 1.23 V

Cu2+ + 2e- Cu E = 0.16 V

Overall:

2Cu + O2 + 4H+ 2Cu2+ + 2H2O

In a deaerated 1 M acid solution:

2H+ + 2e- H2 E = 0 V

Cu2+ + 2e- Cu E = 0.16 V

Overall:

Cu + 2H+ Cu2+ + H2

Is the corrosion of copper in an acidic solution spontaneous? Always?Is the corrosion of copper in an acidic solution spontaneous? Always?

Pourbaix diagram: for copper in a non-complexing aqueous soln at 25C

Pourbaix diagrams give info about thermodynamics only

Kinetic factors may predominate in many situations

What info can be found on a Pourbaix diagram?What info can be found on a Pourbaix diagram?

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

Potentials for redox couples as a function of pH e.g. M/Mn+ and Mn+/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

Horizontal lines: Equilibria involving e- transfer, but

NOT H+/OH-

e.g. Cu2+ + 2e- Cu(s)

Between pH -2 to 6 Cu dissolves for potentials ~0.16 V.

Vertical lines: Equilibria involving hydrolysis, but

NOT e- transfer

e.g. Cu2+ + H2O CuO(s) + 2H+

At pH 7: Cu2+ concentration is reduced below 10-6 M passivation.

Above pH 7, Cu2+ will not be the major corrosion product.

Sloping lines: Equilibria involving both hydrolysis

and e- transfer

e.g. 2Cu(s) + 2H2O Cu2O(s) + 2H+ + 2e-

pH 6-14: corrosion product may be Cu2O, but this may oxidise further.

pH > 7: Cu2+ will not be the major corrosion product if other oxidising agents are present.

O2 + 2H2O + 4e- 2OH- E = 0.4 V

2H2O + 2e- H2 + 2OH- E = -0.83 V

H2O is stable in the region between the lines

2H+ + 2e- H2 E = 0 V

O2 + 4H+ + 4e- 2H2O E = 1.23 V

Dashed lines: Equilibria involving the redox couples A = H+/H2 and B = H2O/O2 as a function of pH

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.

Slope = 0.059 V per pH unit.

B

A

C25atunitpHperVn

H05916.0slope o

Corrosion potential:

- the potential of the metal surface in contact with electrolyte where corrosion occurs.

- no net current flows at the corrosion potential.

ia

ic

KINETICS OF CORROSIONKINETICS OF CORROSION

Corrosion current:

- the exchange current at the corrosion potential.

Oxidation = corrosion of metal

Reduction of substance in contact with the metal

How is the rate of corrosion determined?

Measure steady state current for metal oxidation and H2 evolution as a function of potential.

Plot graph of logi vs E a Tafel plot

Extrapolate lines till they overlap

i.e. logia = log-ic = logicorrosion

Change in io or the Tafel slope change in corrosion rate

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.

GALVANICGALVANIC CORROSIONCORROSIONGalvanic 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.

Cu2+(aq) + 2e- Cu(s) E = +0.34 V

Fe2+(aq) + 2e- Fe(s) E = -0.44 V

Zn2+(aq) + 2e- Zn(s) E = -0.76 V

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 cathodeDegree 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.

CORROSION PROTECTIONCORROSION PROTECTION

i) ElectrolysisSurround metal to be protected by inert anodes and pass a current (icath) between the metal and anodes.

log

log

ii) Sacrificial anodeAnother metal with a more negative m potential is place in good electrical contact with the metal to be protected.

log

log

Sacrificial metal will enforce its corrosion potential on the metal surface

Rate of metal dissolution reduced from icorrosion to iprotected.

i) Electrolysis

log

log

ii) Sacrificial anode

log

log

Problem:

H2 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:

Common sacrificial anodes: Zn, Mg, Al

Al anodes mounted on a steel jacket structure

Eo /V

Zn -0.76

Mg -2.36

Al -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 meansSurround metal by cathodes and apply a potential (e.g. anodisation of Al)

ii) Chemical meansAdd 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. O2, 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!