3. Corrosion
(Ref.: Pletcher: Industrial Electrochemistry; M. Pourbaix: Lectures on Electrochemical
Corrosion; Bockris et al: Comprehensive Treatise on Electrochemistry, Vol. 4; Bockris
and Reddy, Modern Electrochemistry, Vol. 2.)
Corrosion is the spontaneous dissolution of a metal by the reactions:
M + (n/4) O2 + (n/2)H2O → Mn+ + nOHM + nH2O → Mn+ + (n/2) H2 + nOH-
and / or
These are the sum of two simultaneous reactions occurring on surface of the metal:
M → Mn+ + ne
and either
O2 + 2H2O + 4e → 4OHO2 + 4H+ + 4e → 2H2O
(in basic media)
(in acidic media)
2H2O + 2e → H2 + 2OH-
or
Thus, rate of corrosion will depend on presence of water, concentration of O2, pH and
composition of aqueous electrolyte. (Also T, nature of metal,..)
3.1 Thermodynamics of Corrosion
M → Mn+ + ne
Consider the reaction:
The reversible electrode potential of this couple is given by:
E = Eo +
RT
ln aM n +
nF
(1)
and is independent of pH.
The cathodic reactions:
2H+ + 2e → H2
or
O2 + 4H+ + 4e → 2H2O
have corresponding potentials:
2
E=E
o
H2
RT aH +
RT
RT
RT
RT
+
ln
= E0 −
pH −
ln aH2 = E o O2 −
pH +
ln aO2
2F a H 2
F
2F
F
4F
or, for aH 2 = aO2 =1,
E = Eo - 0.059 pH
(2)
(3)
i.e. E is a function of pH.
The pH dependence of the electrode potentials can thus be represented in a Pourbaix
diagram:
57
1.23
O 2 + 4 H + + 4e = 2 H 2 O
Eo
0
2 H + + 2e = H 2
− 0.441
Fe = Fe 2+ + 2e
pH
The complete Pourbaix diagram for the iron system (taking all chemical and
electrochemical equilibria into account) is given below. Details on drawing the diagram
are given in the last page of this chapter.
Fe3+
FeO4
1.23
2−
O2 + 4 H + + 4e = 2 H 2O
E
Fe2O3
Fe
0
2+
2 H + + 2e = H 2
Fe(OH ) 2
− 0.618
Fe
9
pH
Corrosion is said to take place if the concentration of the metal ion (Fe2+) is of order of
10-6 M. Thus, assuming the immediate product of iron dissolution is Fe2+, application of
Nernst equation yields a potential of: E = -0.441 + (0.059/2) log 10-6 = -0.618. Thus the
equilibrium line between Fe and Fe2+ shall be at -0.618 V, independent of pH. Because
the hydrogen evolution line lies above the Fe / Fe2+ line, iron is unstable in aqueous
solutions at all pH’s. Other equilibria that may exist are that between Fe2+ and
Fe(OH)2 (s):
Fe(OH)2 + 2H+ = Fe2+ + 2H2O
The equilibrium constant is K = [Fe2+]/[H+]2 = 1013.29, or log [Fe2+] = 13.29 – 2pH. If
[Fe2+] = 10-6 M, then pH = 9.6.; i.e Fe(OH)2 is stable for pH > 9.6. The Fe2+ / Fe(OH)2
58
equilibrium depends only on pH and is independent of potential; therefore, it is a
vertical line.
Exercise. Draw the Pourbaix diagram for the nickel system. Use the electrode potentials
given at the end of this chapter.
3.2 Kinetics of Corrosion
Consider a metal M in O2 – free aqueous solution.
i
2 H 2 O + 2e = H 2 + 2OH
Ecorr
−E
log i
−
M − ne = M n +
2 H 2 O + 2e = H 2 + 2OH −
log icorr
M − ne = M n +
log io
M
log io
H
E corr
E eq
E eq
M
−E
H
At E = Ecorr , i = icorr, no net current flows, iM / M n + = iH + / H = icorr
2
Equation for the two Tafel lines are:
η H = ( E − E eq H ) = − β H (ln io H − ln i )
(4)
η M = ( E − Eeq ) = β M (ln i − ln io )
M
where
βH =
RT
;
α nF
M
βM =
RT
(1 − α )nF
(5)
i = icorr at E = Ecorr.. Thus,
(a)
( E corr − E eq ) = β H [ln io − ln icorr ]
(6)
(b)
( E corr − E eq ) = β M [ln icorr − ln io ]
(7)
H
H
M
M
Eliminating Ecorr from (a) and (b), and solving for ln icorr
( E eq − E eq ) + ln[(io ) β H (io ) β M ]
H
ln icorr =
M
H
βH + βM
M
(8)
59
β1
or,
icorr = (io ) (io
where
β1 =
H
M
⎡ E eq H − E eq M
) exp ⎢
⎢⎣ β H + β M
β2
βH
β2 =
βH + βM
⎤
⎥
⎥⎦
(9)
βM
βH + βM
(10)
Multiplying (a) by βM , (b) by βH and adding ,
β M ( E corr − E eq H ) + β H ( E corr − E eq M ) = β H β M ln io H − β H β M ln io M
Solving for Ecorr,
Exercise:
E corr
⎛i H
β M E eq H + β H E eq M
β β
=
+ H M ln⎜⎜ o M
βH + βM
β H + β M ⎝ io
(11)
⎞
⎟
⎟
⎠
(12)
Confirm the last equation.
3.3 Passivation
Passivity explains the stability of metals such as Al, Cr, Ni and Pt in damp air, where a
thin layer of non-porous and insoluble oxide film protects the metal. In passive region,
metal is covered by an adjacent and non-porous film (1-15 nm thick).
O 2 evolution +
i
breakdown of film
active corrosion
M − ne = M n +
red ' n of O 2 or H 2 O
Flade potential
passive region
E
(+ve potential )
Theoretical conditions of corrosion, passivation and immunity for iron:
Passivation
E
Corrosion
Im munity
Corrosion
pH
60
Natural passivation occurs if, in the solution, there is a species, e.g. O2, or a redox
reagent, capable of taking the surface potential into a passive region. Even when
passivation is observed, there is further phenomena which must sometimes be
considered – viz pitting (see below).
3.4 Types of Corrosion
Uniform corrosion
A case of uniform corrosion is the corrosion of reinforcing steel in concrete. Uniform
corrosion is basically thermodynamically controlled, with the redox potentials and the
Nernst equation dictating the process until concentration polarization takes place when
the transport of oxygen is limited. For iron, the anodic reaction is:
2+
Fe → Fe + 2e-
while the cathodic reaction is oxygen reduction to form hydroxyl ions:
-
O2 + 2 H2O + 4e- → 4 OH
As the whole surface of steel corrodes the anodic sites and cathodic sites become the
same surface.
Galvanic Corrosion.
Galvanic corrosion results from two different metals being in contact in the
environment. Examples would be brass plumbing fittings on a cast iron pipe. In this case
several reactions are possible, but in general the corrosion rate of the most anodic or
active metal is increased and the corrosion rate of the more cathodic metal is decreased.
A good example is zinc in hydrochloric acid. The anode reaction favored is:
2+
Zn → Zn + 2eCathode:
+
2H + 2e- → H2
Factors Affecting Galvanic Corrosion.
ƒ
Area Effect. When current flows between the anode and cathode, the current
will be the same in the anode and cathode independent of the surface area of
each electrode. It is the current rather than the current density which is equal for
the anodic and cathodic reactions. Therefore, if the current flowing between the
2
anode and the cathode is one amp and the surface areas are one cm , then the
2
current density in each electrode is one A/cm . However, if the area of the anode
2
is only 0.1 cm , then the current density in the anode with the same one amp
flowing is 10 A/cm2. From Faradays Law, the corrosion rate depends on the
current density in the anode. In this case decreasing the surface area of the anode
increases the corrosion rate by a factor of 10.
61
As a general rule to minimize galvanic corrosion, the anode area should be large
and the cathode area should be small.
ƒ
ƒ
ƒ
For protection from galvanic corrosion, the cathode of the system should be
painted if a coating is applied. This arises from the area effect, in that if the paint
is damaged. by a scratch for example, then a small cathode to large anode area
ratio is formed which results in minimizing corrosion rates. If the anode is
painted, then damage to the paint results in a large cathode to small anode ratio
which results in large corrosion rates in the anode and rapid penetration into the
metal.
Physical Distance Effect. Galvanic corrosion rates are the largest at the interface
between the anode and cathode and decrease with distance away from the
contact region. As the anodes and cathodes are both good conductors, electron
transport is very good. What is more difficult is the ionic transport in the
electrolyte. As the distance between the anodic reaction site and the cathode
reaction site increases the transport of the ions becomes more difficult and the
corrosion rate decreases. Essentially the resistance between the anode and
cathode increases with distance.
Distance Apart in the Galvanic Series. For selection purposes metals close
together on the list are desirable as there is little driving force for corrosion to be
accelerated.
Temperature Effects. With increasing temperature above 180oF, zinc will form a
protective layer and become cathodic to iron. The zinc becomes nonprotective
and aggressive to the iron.
Crevice corrosion.
Crevice corrosion is a geometrically controlled form of corrosion. It occurs below rivet
heads, between lap joints, in threads and anywhere a small crevice is formed in which at
least one side is a metal.
Mechanism of Crevice Corrosion.
The general conditions for crevice corrosion include a stagnant solution and a gap
between two surfaces, one of which is metal, of the order of 0.002". Initially, the usual
anodic and cathodic reactions occur over the surface of the metal. The general anodic
reaction is:z+
M → M + zeThe general cathodic reaction is :
O2 + 2H2O + 4e- → 4 OHThese initially occur over the whole surface. However a restriction occurs in the crevice
region such that the dissolved oxygen in the crevice cannot easily be replaced. The
region inside the crevice cannot then support a cathodic reaction. It can still support an
anodic reaction of the type shown above. Outside the crevice region the cathodic
reaction proceeds but anodic reaction ceases as it is concentrated in the crevice. An
electrical charge imbalance exists between the high positive charge within the crevice
from metal ions and the negative charge outside the crevice. As a result, negative ions
62
are attracted into the crevice. The limit is the small size of the crevice. Chloride ions are
the favored ions to be attracted into the crevice. Associated with the negative chloride
ion is the very small positive hydrogen ion. Both the chloride ion concentration and the
hydrogen ion concentration increase within the crevice. That is the pH in the crevice
decreases from values of 6 to 2 - 3. The effect of this acidification is that the corrosion
rate inside the crevice increases. Reactions inside the crevice include:+
+
M + Cl → M Cl
z+
-z
-
M Cl + zH2O → M(OH)z + z H+Cl
H+Cl → H+ + Cl
This results in acidification within the crevice. Note that only the region inside the
crevice will be corroded. This is also important as the anodic area is localized and small
in comparison to the cathodic area. The area effect then also comes into play with a
small anode carrying the same current as the cathode, leading to an increased current
density and corrosion rate.
Pitting corrosion
Pitting corrosion is a form of localized corrosion as it does not spread laterally across an
exposed surface rapidly but penetrates into the metal very quickly, usually at an angle of
90o to the surface.
Stagnant solution conditions favor pitting corrosion. The presence of halide ions,
chloride, fluoride bromide and iodide, can all pit metals.
The metals recognized for pitting are materials which were designed as passive metals,
such as the stainless steels, aluminum alloys and nickel alloys. This tends to be due to
the difficulty in obtaining uniform corrosion in these alloys.
From a mechanistic point of view, the growth of a pit can be regarded as similar to the
corrosion process in a crevice. The exposed surface outside the growing pit is
cathodically protected by supporting the reduction of oxygen to hydroxyl ion reaction:O2 + 2H2O + 4e- → 4 OHAs this cathodically protects the region outside the pit, the metal dissolution region
cannot spread laterally across the surface. In addition the large cathodic surface can
maintain this reaction and form a large cathode to small anode ratio which will
accelerate the anodic reaction.
Within the pit, which is regarded as a small hemisphere at this stage, the metal
dissolution reaction is taking place. This is the general anodic reaction of:
z+
M → M + ze-
63
However, it is the only reaction within the pit and results in an electrical imbalance
again which attracts negatively charge ions, usually chloride ions. The autocatalytic
reaction to form hydrochloric acid in the pit is initiated and continues:
-
Mz+Clz- + zH2O → M(OH)z + zH+Cl
Pitting, like crevice corrosion, is an autocatalytic reaction once it is started and the pH
decreases while chloride ion concentration increases inside the pit.
Other forms of corrosion reduce the stress bearing capability of the material, such as
stress corrosion cracking (partly due to mechanical forces applied to metals), corrosion
fatigue, fretting fatigue and hydrogen embrittlement (when corrosion reaction occurs
with evolution of hydrogen which enters the metal lattice and thus reduce the strength of
inter-atomic bonds). In these cases the material will fail at stress levels below those
expected.
3.5 Corrosion Prevention by Electrochemical Methods
3.5.1 Assuming that α = ½ , the corrosion current can be expressed as:
⎡ F ( Eeq H − Eeq M ) ⎤
M H
icorr = (io io )1 / 2 exp ⎢
(13)
⎥
4 RT
⎢⎣
⎥⎦
M H
Hence, to reduce corrosion current, the term (io io )1/ 2 must also be reduced.
This can be accomplished as follows: if H2 evolution is the cathodic reaction,
addition of P, As, Sb, compounds educe the exchange current density for H2
evolution. If O2 reduction is the cathodic reaction, adding substances which react
with dissolved O2 reduces the concentration of O2 and therefore also the
corresponding io. Examples are: hydrazine and sulfite; N2H4 + 5O2 → 4NO2- +
4H+ + 2H2O; SO32- + O2 → 2SO42-.
Alternatively, ioM can be reduced by adding substances which adsorb on metal
electrode; examples are aliphatic and aromatic compounds; thiourea and
derivatives; amines; sulfur compounds and carbonyl compounds; etc)
3.5.2 The complete log I vs E characteristic for a metal suggests two potential regions
where potential could be usefully held: at –ve potentials or in the passive region.
Former is cathodic protection, and latter anodic protection.
Cathodic Protection
This form of corrosion prevention involves making the surface to be protected
the cathode of the system. There are two ways by which cathodic protection may
be applied; these are by the use of sacrificial anodes or via an impressed current
system.
64
O 2 + 4 H + + 4e = 2 H 2 )
with inhibitor
logi
M − ne = M n +
icorr
2 H + + 2e = H 2
log(icorr ) O2
icorr
log(icorr ) noO2
E
log i
reduction of O 2 or H 2 O
cathodic protection
anodic protection
icorr
E
−
+
e
soil
pipeline
ƒ
soil
sacrificial anode
pipeline
auxilliary electrode
Sacrificial anodes
The application of sacrificial anodes in cathodic protection is based on the
differences in electrochemical reactivity of metals. A quick look at the galvanic
series of metals reveals that if zinc and iron are connected together and
immersed in a corrosive electrolyte a galvanic couple is set up, whereby the zinc
is preferentially oxidized. Zinc galvanisation is a good example of cathodic
protection, where the Zn layer as well as forming a protective coating forms a
sacrificial anode which cathodically protects the underlying metal.The sacrificial
anode should possess fast or facile anodic kinetics (relatively large currents at
small polarisations) and must not passivate.
65
ƒ
Cathodic Protection by Impressed Current.
The objective here is to ensure the component requiring protection is maintained
in its cathodic region by the application of a voltage or cathodic current. The
system is shown schematically below:
DC rectifier
- ve
+ve
electrons
Structure to protect
Anode in impressed
current system
Anodic protection.
An impressed current technique can be applied if the material passivates in the
particular environment. In this case the structure is made more anodic by
drawing electrons out of it until it enters the passive region.
3.6 Other Methods of Corrosion Prevention
Inhibitors. Inhibitors are used to reduce and block corrosion.
Adsorption inhibitors.
Adsorption inhibitors protect by adsorption on to the metal or metal oxide film
exposed to electrolyte. Organic inhibitors are aliphatic and aromatic amines (N
compounds), thiourea( S compounds) and aldehydes (O compounds). All these
have a charged state, for example aliphatic amines have ammonium cations
present, R3NH+. The S and O compounds have a negative charge on them.
Thiourea bonds strongly to a metal by sharing its electrons with the metal
surface. This blocks solvating water molecules and also stops hydrogen gas
molecule formation.
Poisons.
These type of inhibitors block either of the hydrogen ion reduction or formation
of hydroxyl ions cathodic reduction reactions. The hydrogen ion reduction
reaction is inhibited by the group V metals or metalloids such as P, As or Sb.
As2O3 is added at about 0.25M. The combination of hydrogen atoms to
hydrogen molecules is blocked in a reaction of the form:AsO+ + 2Hads + e- → As + H2O
Alternatively:
As2O3 + 6Hads → 2As + 3H2O
Scavengers
66
Scavengers act to remove the oxygen preferentially before it can be used in the
cathodic reactions. Two popular examples are hydrazine and the sulfite ion.
+
N2H4 + 5/2 O2 → 2 NO2- + 2 H + H2O
2-
SO3 + ½ O2 → SO4
2-
Filming Inhibitors.
The addition of specific ions with high redox reaction potentials will produce
local reactions to form protective films. Two ions of this type are the chromate
and nitrite ions. The redox reactions are:
+
+
NO2- + 8 H + 6e → NH4 + 2H 2O Eo = + 0.9V
2-
+
2 CrO4 + 10 H + 6e → Cr2O3 + 5 H2O Eo = +1.31V
Both these reactions induce iron to dissolve in the ferric state with 3+ rather than
in the ferrous state as 2+. The ferric oxides are stable on the surface and block
further corrosion.
+
3+
Fe + 3 H2O → Fe2O3 + 6 H
Vapor phase.
These tend to be nitrites, carbonate and benzoate filming inhibitors attached to
parachutes of an organic cation. An example is dicyclohexyl ammonium nitrite.
The inhibitor evaporates onto the metal surface.
Simple calculations of corrosion rates
The measured current and the corrosion rate are simply and directly related as
follows.
In the case of ferrous metals for which Fe = Fe2+ + 2e, the corrosion of 1 mol Fe
= 56 g = 0.056 kg Fe releases 2 mol electrons = 2 x 96485 C/mol ; therefore a
current of 1 A corresponds to a corrosion rate of 0.056/(2 x 96485) kg/s iron.
More useful units are g/year; or if density and surface area are know mm/y depth
of penetration.
_
Example: If, by extrapolation of the η vs. log i plot, the corrosion current is found to be:
log icorr = -3.13, icorr = 7.41 x 10-4 A, then 7.41 x 10-4 A = 7.41 x 10-4 / 2 x 96485 =
3.84 x 10-9 moles of Fe corroding per second = 0.0185 g /day .
Exercise
For a metal M that corrodes with an evolution of H2, what is the corrosion current (in A
cathodic transferH coefficients are each equal to
m-2) if it is assumed that the anodic and
M
E
0.5, the equilibrium potentials are eq = -0.60 V and E eq = -0.165 V, and the exchange
current densities are ioM = 0.01 A m-2 and ioH = 0.05 A m-2 . (n = 2).
67
Annex I
Electrode potentials for the Iron system
Fe + 2e = Fe
Fe3+ + e = Fe2+
Eo (V)
-0.409
0.771
Fe2O3 + 6H+ + 6e = 2Fe + 3H2O
Fe3O4 + 8H+ + 8e = 3Fe + 4H2O
Fe(OH)2 + 2e = Fe + 2OHFe(OH)3 + 3H+ + 3e = Fe + 3H2O
Fe2O3 + 6H+ + 2e = 2Fe2+ + 3H2O
Fe3O4 + 8H+ + 2e = 3Fe2+ + 4H2O
Fe(OH)3 + 3H+ + e = Fe2+ + 3H2O
3Fe2O3 + 2H+ + 2e = 2Fe3O4 + H2O
Fe3O4 + 2H+ + 2e = 3FeO + H2O
Fe2O3 + 2H+ + 2e = 2FeO + H2O
-0.051
-0.085
-0.877
0.059
0.728
1.230
0.939
0.221
-0.197
-0.057
2+
Electrode potentials for the Nickel system
Ni2+ + 2e = Ni
Ni(OH)2 + 2H+ + 2e = Ni + 2H2O
Ni2O3 + 6H+ + 2e = 2Ni2+ + 3H2O
Ni3O4 + 8H+ + 2e = 3Ni2+ + 4H2O
Ni)2 + 4H+ + 2e = Ni2+ + 2H2O
-0.257
0.110
1.753
1.977
1.593
Exercise
Iron corrodes in de-aerated water at pH 2.8 to give a solution that is 0.01 mol dm-3 in
Fe2+. The exchange current density is 0.01 A m-2, the transfer coefficient is 0.5 and the
formal electrode potential is -0.66 V. If the cathodic reaction is the evolution of
hydrogen (exchange current density of 0.05 A m-2, α = 2), what is the corrosion
potential?
68
Annex II
Details on drawing the Pourbaix diagram for the iron system
The iron system:
(a)
Fe3O4 + 8H+ + 8e = 3Fe + 4H2O
Eo = -0.085
0 . 059
1
E = − 0 . 085 −
log
= − 0 . 085 + 0 . 059 log[ H
8
[ H + ]8
Fe3O4 + 8H+ + 2e = 3Fe2+ + 4H2O
(b)
E = 1 . 23 −
0 . 059
[ Fe
log
2
[H
2+
+
]
]8
3
= 1 . 23 + 0 . 531 − 0 . 236 pH
(c)
(d)
(e)
= 1 . 23 − 3 ( 0 . 0295 ) log[ Fe
2+
] + 8 ( 0 . 0295 ) log[ H
+
]
= 1 . 761 − 0 . 236 pH
Eo = 0.221 V
0 . 059
1
log
= 0 . 221 + 0 . 059 log[ H + ] = 0 . 221 − 0 . 059 pH
2
[ H + ]2
Fe2O3 + 6H+ + 2e = 2Fe2+ + 3H2O
E = 0 . 728 −
] = − 0 . 085 − 0 . 059 pH
Eo = 1.23 V
3Fe2O3 + 2H+ + 2e = 2Fe3O4 + H2O
E = 0 . 221 −
+
2+
Eo = 0.728
2
0 . 059
[ Fe ]
log
= 0 . 728 + 0 . 354 − 0 . 177 pH
2
[ H + ]6
Fe2O3 + 6H+ + 2e = 2Fe2+ + 3H2O
2Fe2+ = 2Fe3+ + 2e
Fe2O3 + 6H+ = 2Fe3+ + 3H2O
Eredo
Eredo = 0.728
= 0.771
For Fe2O3 + 6H+ + 2e = 2Fe2+ + 3H2O, ΔGo = -2(96,500)(0.728) = -140,504 J mol-1
For Fe2+ = 2Fe3+ + 2e, ΔGo = +2(96,500)(0.771) = 148,803 J mol-1
Therefore, for Fe2O3 + 6H+ = 2Fe3+ + 3H2O,
ΔGo = (148,803 – 140,504) J mol-1 = 8,299 J
-1
mol
= - RT ln K = -2.303 RT log K = 8, 299 J mol-1
3+ 2
from which, log K = -7.714, where K = [Fe } / [H+]6; i.e. log K = 2 log [Fe3+] – 6 log [H+] =
-7.714. With [Fe3+] = 10-6, pH = 0.714, (i.e. a vertical line at pH = 0.714).
(f)
For the equilibrium between Fe(OH)2 and Fe2+, consider the following:
(i)
(ii)
(iii)
Fe(OH)2 + 2e = Fe + 2OHFe = Fe2+ + 2e
Fe(OH)2 = Fe2+ + 2OH-
Eo = -0.877 V
E = -0.409 V
o
For (i), ΔGo = -2(96,500)(-0.877) = 169,261 J mol-1
For (ii), ΔGo = 2(96,500)(0.409) = -78,937 J mol-1. Therefore, for (iii), ΔGo = 90,324 J mol-1 =
- RT ln K, from which log K = -15.83, where K = [ Fe 2 + ][OH − ]2 . Thus, log K = log [ Fe 2+ ]
+ 2 log [OH-]
With [Fe2+] = 10-6, -15.83 = -6 + 2 (pH – 14) ; pH = 9.1 (i.e. a vertical line at pH 9.1)
69