Transcript ppt - Wits Structural Chemistry

CORROSION
INTRODUCTION
THERMODYNAMICS OF CORROSION
KINETICS OF CORROSION
GALVANIC CORROSION
CORROSION PROTECTION
INTRODUCTION
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 O2
 presence of complexing agent
 pH
Reaction of most metals with O2  thermodynamically favourable
 Some form a protecting oxide layer (passive layer)
e.g.
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)
Titanium hip prosthesis
Stainless steel: steel is made corrosion resistant by alloying with Cr
 forms Cr2O3 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 O2  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 + 3/2H2O + 3/4O2  Au(OH)3
Au nuggets
Pt nuggets
That is why they can be found in metallic
form on Earth, and it is a large part of their
intrinsic value.
Au ore body
Iron objects were found remarkably preserved after centuries of immersion
at the bottom of a peat bog. Why???
THERMODYNAMICS OF CORROSION
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 Cu2+.
Calculate the equilibrium potential for this solution:
E(Cu2+/Cu) = +0.34 V
(vs SHE)
 Cu2+(aq) + 2e-  Cu(s)
E  Eo 
RT
ln Q
nF
 E  Eo 
E = +0.34 V
0.05916
log Q
n
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?
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/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
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. 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.
Horizontal lines:
Equilibria involving e- transfer, but
NOT H+/OH-
Sloping lines:
Equilibria involving both hydrolysis
and e- transfer
e.g. Cu2+ + 2e-  Cu(s)
Between pH -2 to 6 Cu dissolves for
potentials ~0.16 V.
e.g. 2Cu(s) + 2H2O  Cu2O(s) + 2H+ + 2epH 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.
H
slope  0.05916 
V per pH unit at 25 o C
n
B
O2 + 2H2O + 4e-  2OH-
O2 + 4H+ + 4e-  2H2O
A
E = 0.4 V
E = 1.23 V
H2O is stable in the region
between the lines
2H+ + 2e-  H2
E = 0 V
2H2O + 2e-  H2 + 2OH-
E = -0.83 V
Dashed lines:
Equilibria involving the redox couples A = H+/H2 and B = H2O/O2 as a
function of pH
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.
KINETICS OF 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
ia
ic
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 H2 evolution as a
function of potential.
 Plot graph of logi vs E  a Tafel plot
 Extrapolate lines till they overlap
i.e. logia = log-ic = logicorrosion
Change in io or
the Tafel slope
 change in
corrosion rate
GALVANIC CORROSION
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.
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
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
CORROSION PROTECTION
1) CATHODIC PROTECTION
The potential of the metal is shifted more negative  lower oxidation rate.
i) Electrolysis
ii) Sacrificial anode
Surround metal to be protected by inert
anodes and pass a current (icath)
between the metal and anodes.
Another metal with a more negative m
potential is place in good electrical
contact with the metal to be protected.
log
log
Rate of metal dissolution reduced from
icorrosion to iprotected.
log
log
Sacrificial metal will enforce its
corrosion potential on the metal surface
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:
Al anodes mounted on
a steel jacket structure
Common sacrificial anodes: Zn, Mg, Al
Zn
Mg
Al
Eo /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. 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!