POLAROGRAPHY TO STUDY COMPLEX FORMATION

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CORROSION
INTRODUCTION THERMODYNAMICS OF CORROSION KINETICS
OF CORROSION GALVANIC CORROSION CORROSION
PROTECTION
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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.
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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
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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)
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Reaction of most metals with O2 ?
thermodynamically favourable
? Some have slow reaction kinetics
Metals such as Zn, Mg, Cd corrode slowly even
though ?G lt 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.
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Why dont precious metals corrode???? ?
e.g. Au, Pt
Au 3/2H2O 3/4O2 ? Au(OH)3
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
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Iron objects were found remarkably preserved
after centuries of immersion at the bottom of a
peat bog. Why???
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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? 0.34 V
?
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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?
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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
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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(n1) ? 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
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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 2e- pH 6-14
corrosion product may be Cu2O, but this may
oxidise further. pH gt 7 Cu2 will not be the
major corrosion product if other oxidising agents
are present.
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B
O2 2H2O 4e- ? 2OH- E? 0.4 V
O2 4H 4e- ? 2H2O E? 1.23 V
A
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.
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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.
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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
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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.
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? 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
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CORROSION PROTECTION
1) CATHODIC PROTECTION
The potential of the metal is shifted more
negative ? lower oxidation rate.
i) Electrolysis
ii) Sacrificial anode
Another metal with a more negative ?m potential
is place in good electrical contact with the
metal to be protected.
Surround metal to be protected by inert anodes
and pass a current (icath) between the metal and
anodes.
log
log
log
log
Sacrificial metal will enforce its corrosion
potential on the metal surface
Rate of metal dissolution reduced from icorrosion
to iprotected.
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i) Electrolysis
ii) Sacrificial anode
log
log
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
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Example of sacrificial anodes used in cathodic
protection
Al anodes mounted on a steel jacket structure
Eo /V Zn -0.76 Mg -2.36 Al -1.66
Common sacrificial anodes Zn, Mg, Al
Sacrificial anodes will corrode at a higher rate
than protected metal ? anodes need to be replaced
periodically
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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)
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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!