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By nadiasima Oct 14, 2014 2601 Words
Fathi Habashi

Montana

School

of

Kinetics of Corrosion of Metals

Mines
Butte

For better understanding of corrosion
processes, it is essential a t first to consider the simple
cases of corrosion when no insoluble products, scales or
films are formed. Once the mechanism of these processes is well established, it may be possible to understand corrosion processes in general. Corrosion processes in this category (no insoluble end

products are formed) can he divided into two broad
groups: (1) corrosion with no hydrogen evolution, and
(2) corrosion with hydrogen evolution.
Hydrogen is liberated in those processes when H+ ion
is acting as depolarizer (electron acceptor). If, however, the solution contains a depolarizer other than H+ ions, as for example Oz, H202,etc., then corrosion proceeds with hydrogen evolution. It is the purpose of this paper to report that all these corrosion processes follow

a single kinetic law.

concentration, the rate depends only on the concentration of the acid, i.e., on the concentration of hydrogen ions, and is independent of the concentration of the
depolarizer. At high acid concentration, the rate is
independent of the acid concentration, depending only
on the concentration of the depolarizer. Further, the
ratio of the concentrations a t which the rate changes its
dependence is a constant (0.12), as can he seen from the
points where the lines change slope in Figure l(a).
Table 1.

Corrosion Reactions with N o Hydrogen Evolution

Examples

Mplecular complexmg agent

Zinc in dilute acetic
acid
Iron in dilute HCI
Copper in aqueous PITH3
Cop er in aqueous
et&lenediamine
Copper, silver, and
rold in CN-solution

Acid

Corrosion With N o Hydrogen Evolution

Three groups of reactions of this type can he identified. These are shown in Table 1, and the kinetics of these processes are plotted in Figures 1-3.
In all these reactions, the aqueous phase is composed
of t,wo components: depolarizer (02, H202, etc.) and H+
ions or complexing agent (NHa, CN-, etc.). These
processes have two things in common: depending upon
experimental conditions, the rate of corrosion is controlled by only one of the components of the aqueous phase; arid there is always a constant ratio between the
concentrations of the two components of the aqueous
~ h a s a t which the rate changes its dependence from one
e
of the components to the other.
The example shown in Figure l(a) illustrates these
two points. This shows the corrosion of zinc in acetic
acid in presence of KN03 as depolarizer. At low acid

Depolarizer

Medium

Anion complexing
went

Ref.

KXO?

(1)

H?02
O2
O?

(3,

O2

(6-9)

(b)

Mechanism. The electrochemical theory of corrosion assumes that when a metal comes into contact with water to which oxygen or any depolarizer is added, the
depolarizer takes up electrons a t one part of the surface
(the cathodic zone) while the metal gives them up at
another (the anodic zone) (10). The cathodic reduction of oxygen for example, a t the surface of the metal may lead to the formation of either hydrogen peroxide
or hydroxyl ions as follows:
O2

--

+ 2Hp0 + 2ec H202+ 20H40HO2 + 2H90 + 48-

-

The anodic reaction may he represented as follows:
31

31"+

+ ne-

0 0 Moles11 HCI

a
.
c

0

I

I

0.02

004

I

006

I

I

008

ACETIC ACID C O N C E N T R A T I O N I M a l e s l l )

((11
Figure 1.

Corrosion of metals in acid,.

0

v

I
008
H,O,

0 24
0 32
CONCENTRATION (Males 1 1
01
6

Ibl

lo) Zinc in acetic acid in the presence of KNO, at 25'C(I1. (b) lronin HCI in the presence of HIOl at 2S°C 121.

318 / Journal of Chemical Education

4)

(6)

OZ PRESSURE

(ATM.1

Ibl

Corrosion of copper in molecular complexing
.hove). 10) Ammonio (3). Ibl NHCHCHcNHs 151.

Figure 2.

Or

PRESSURE

agents

(left a n d

(ATM.1

Id

The liberated metal ion hydrolyzes forming insoluble
compounds', as follows:
M"+ + n HnO M(OH), + n H+

-

The insoluble hydrolytic products form a thin 6lm
which adheres firmly on the anodic zone (Fig. 4), thus
preventing or slowing down further dissolution of the
metal. If,however, the aqueous phase contains an acid
or a complexing agent, these insoluble products will be
dissolved as soon as formed.
Kinetic Consideration. Like any other heterogeneous
process, the surface of the metal in contact with an
aqueous phase is thought to be covered by a thin layer
of a stagnant liquid (the Nernst boundary layer),
through which the reactants have to diffuse before
1 Evans and Davies (11) described experiments in which pure
metallio einc was suspended in distilled water through which g ion per
oxygen was continuously bubbled. After 4 days,
liter of Znz+ was detected, and this amount did not increase. This was in agreement with the coneentmtion predicted from
the solubility product of einc hydroxide.

reaching the metal surface. If the rate of the chemical
reaction at the interface is much higher than the rate of
diffusion, the process will be diffusion-controlled. On
the other hand, if the rate of chemical reaction is much
slower than the rate of diiusion, then the process will be
chemically-controlled.
The rate of the cathodic reaction can be given in its
simplest form, when first order kinetics are followed,
by the equation:
V = ktA,[DI
z
(1)
where k,

velocity constant'
surface area. of the cathodic zone
[Dl = concentration of the depolarizer

A1

=
=

The rate of the anodic reaction can be given by the
1 When the rate is diffusion-controlled, the velocity constant = D/6 where D = the diffusion coefficient of the reagent, and 6 = the thickness of the Nernst boundary layer. When the rate is chemically-controlled, k = A ecEIRT the familiar Arrhenius

equation. A diffusion-controlled process is always first order; a ehemically-controlled process can be any order. However, the discussion here is limited to processes which follow first order kinetics, established by either mechanism.

"

ol/

25
1

5

'

'

I0

'

15~10~'

KCN concentrotion

F:g.,=

3.

.gent

IKCN solulion a l 2SLC.).

Corrorion of metals :n ononic complexing
(01 Copper 181. (bl

Silver

191. CI Gdd 181.

Po,=021 ( A T M )

n
0

5
1
0
K G N CONCENTRATION

(Moleslll
(cl

Volume 42, Number 6, June 7 965

/

319

equation:
Vz

=

kzA4CI

(2)

where k2 = velocity constantP
Az = the surface area. of the anodic zone
[C] = concentration of H + ions, or the complexing agents

At the steady state, the rate of the cathodic reaction =
the rate of t,he anodic reaction, i.e.
k,A,[DJ = MzICl
(3)
But, since
A

=

A,

+ A*

(4)

(where A is the total surface mea of the metal in contact with the solution) therefore:

Under these conditions, the rate depends on both
concentrations of C and D. These are the points where
t,he rate curve changes its direction. This change takes
place at a certain ratio of [C]/[D] as can be deduced
from equation (8)

Equation (5), derived theoretically, is therefore in agreement with experimental data and is a general equation for a large number of corrosion processes.
Corrosion With Hydrogen Evolution

Evolution of hydrogen during the corrosion of metals
may take place either in acidic or in alkaline medium.
I n these cases hydrogen ion is acting as depolarizer, and
the reactions taking place are:

-

+ + nec
Hz

cathodic reaction: 2H+ 2eanodic reaction: M - M e f

At low concentration of D and high concentration of
C, the first term in the denominator may he neglected in
comparison with the second, and the velocity equation simplifies to: Rste = klAID]
(6)

The two cases will be considered separalely.
Aeidic medium. Beside acting as depolari~er,~
HC
ions also prevent the hydrolysis of metal ions:

The rate of corrosion under these conditions depends
only on the conceutration of the depolarizer. This is in
agreement with the experimental facts plotted in Figures 1-3. On the other hand, at high concentration of D and low concentration of C, the second term in the
denominator may he neglected in comparison with the
first, and the rate equation simplifiesto:

preventing the deposition of insoluble products that
may block the anodic zone. Such mechanism is shown
schematically in Figure 5. The kinetic equation for
this process can be deduced from equation (5), by suhstituting the hydrogen ion concentration for [Dl and [CI:

Rate = &A [C]

(7)

The rate of corrosion in this case is only a function of
the hydrogen ion concentration or the complexing
agent concentration. This is again in agreement with
the experimental facts plotted in Figures 1-3.

= k

+

A [H+]

(11)

where k = klkJkl
kr. Equation (11) is the usual
velocity equation for the corrosion of metals in acids
with H, liberation. Thus Straumanis and Chen (12)
reported that the corrosion rate of titanium in hydrofluoric acid, (Ti 3 HF + TiF,
HZ),in absence

+

+

a Hydrogen overvoltage of the metal may be rate-determining
in this process. This discussion applies only to cases of first order kinetics.

A,
Calhodlc Area

HydrolytC Products:

M"'+ n H & l z M I O H h + n H f

nydrolytic Products:
M "'tn

i
Figure 4. Schematic representation of corrosion of metals
gen evolution.

320

/

Journal of Chemical Education

HzO=

M iOHl.+n~'

Figure 5. Schematic representation of corrosion of metois in acids with hydrogen evolution.

0
Figure 8.

I

4
6
8
1
0
12
1
4
1
6
1
8
20
N o O H CONCENTRATION
IMoles / I )
R o k of dissolution of aluminum in sodium hydroxide (141.

where K

2

=

[H+] [OH-].

When

H F CONCENTRATION ( M o l e s N )
Figure 6.

Rote of dirrolvtion of titanium in hydrofluoric acid (I?).

of air, increases linearly vith increasing acid concentration as shown in Figure 6. Alkaline medium. An example of this case is the
dissolution of aluminum in NaOH. This process
follows the same equations for cathodic and anodic
reactions; hydrogen ions also act as the depolarizer.
The function of OH- is to complex the liberakd metal
ion into a soluble form:
M(0H).
mOH- M(OH),&

+

-

-

l n the case of aluminum, aluminate ion is formed:
AI(OHh + OH- Al(OH),This mechanism is shown schematically in Figure 7.

equation (12) reduces to:
This coiricides with the experimental results by Straumanis and BrakSs (15), who found that the rate of dissolution of high purity aluminum in NaOH increases linearly with NaOH concentration in the range 0.3 to

3.0 molar. When

equation (12) reduces to
Rate

=

kXA

1
[OH-I

(14)

Equation (14) shows that beyond a certain hydroxide
ion concentration, the rate of corrosion should decrease
with increasing hydroxide concentration. This agrees
with the experiment,al results of Balezin and IUimov
(14) in their systematic study of the dissolution of
aluminum in different alkalies a t variable concentrations. These authors reported that beyond 5 M NaOH the rat,e of dissolution decreases with increasing NaOH,
as shown in Figure 8.

COthOdlC Are0

Autocatalytic Corrosion
Anodic A r m

Figvre 7. Schematic representation of corrosion of metals in alkalies with hydrogen evolution.

The kinetic equation of this process is also equation
( 5 ) , uhich after substituting the concentrations of the

depolarizer (H+ ion) and the complexing agent (OHion), becomes: Rate = kAA[H+I[OH-]
k,[Htl
blOH-I

+

When the product of corrosion reacts further with
the metal undergoing corrosion, the rate of corrosion of
the metal will be accelerated and the process will be
autocatalytic. An example of this is the corrosion of
copper in dilute sulphuric acid in presence of oxygen.
The copper(I1) ion formed reacts with metallic copper
to form copper(1) ion which is rapidly oxidized by oxygen to regenerate the copper(I1) ion. The kinetics of two such superimposed processes can
be established by cousidering each process separately.
The overall velocity equation will be the sum of the
velocities of these two parallel reactions, namely, the
electrochemical dissolution reaction and the dissolution
due to the autocatalytic action, (the catalytic term).
The catalytic term in this case can be derived by
assuming that the rate depends on the diffusion of the
copper@) ion away from the metal interface. But since
Cu

+ CUP+= 2cu+

Volume 42, Number 6, June 1965

/ 321

Therefore, the rate should be proportional to the square
root of the copper(I1) concentration:
Rate

=

K"[Cu2+1'/%

(17)

This can be verified if the experimental data by Kinevsky (15),who studied the effectof copper(I1) ion on the corrosion of copper in sulfuric acid, are plotted in the
form shown in Figure 9. In this plot, the linear proportionality between the rate of corrosion and the square root of the copper(I1) ion concentration, as
implied by the mechanism suggested, is demonstrated.

kinetic law applied, no matter what metal was undergoing corrosion, what the composition of the aqueous phase, or whether HI was evolved or not, provided only
that no insoluble end products were formed. The data
can be interpreted on the basis of the electrochen~ical
theory of corrosion and the derived rate equation (5)
which governs all these processes is in agreement with
experiments.
In the special case when H+ is acting as depolarizer,
equatiou (5) reduces to equation (11) in acidic medium
and to equation (12) in alkaline medium.
When the products of corrosion react further with the
metal, corrosion will be accelerated, and a rate equation of the type given in equatiou (18) is valid. Acknowledgmenl

This study was carried out in part during a fellowship awarded to the author by the National Research Council of Canada, to which the author wishes to extend his thanks. Literature Cited

0

L-Lpl--~~-~'
02

. -.
,

04

[c"'.

06
I+

I

8

10

08
iMOk./l

1'

Figure 9. Effect of sapperlil) on the corrosion of copper in HSOr (2.26 N and under Ph = 1 dm1 11.5).

(Decrease in rate a t high copper (11) concentration is
due to increased O2 solubility.) The overall kinetic
eqnation will be:

In the initial stage of corrosion, when only small
amounts of reaction products are formed, the catalytic
term can be neglected, and the rate equation simplifies to
equatiou ( 5 ) . When the reaction proceeds further, the
catalytic role becomes predominant and equation (18)
simplifies to equation (17).

(1) KINO, V., AND SCHACK, J . Am. Chem. Soe., 57, 1212
C.
M.,
(1935).
(2) ABRAMSON, B., AND KING,C . V., J . Am. Chem. Soc., 61,
M.
2290 (1939).
(3) HALPERN, J . Electrochem. Soc., 100,421 (1953).
J.,
(4) H ~ A S F., Ber. Bumengesellschajt Physik. Chem., 67, 402 H ~

- - -- .

(14fi2)
\
,

(5) HALPERN, MILANTS, AND WIYILES, R., J . ElectroJ.,
H.,
D.
chem. Soc., 106, 647 (1959).
(6) DEITZ, A,, AND HALPERN, J . Metah 5, I109 (1953).
G.
J.,
(7) NABASHI, In press.
F.,
(8) KaKovs~Ir, A,, AND KHOLMANSKIKH, B., h.
I.
Yu.
Akad.
Nauk SSSR Otd. Tekh. Nauk, Met. i Toplivo, No. 5, 207
(1QROI.
\----,.
(9) K,\~ovsnrr,I. A., AND KAOLMANSKIKH, B., IZV.M .
Yu.
A
Nauk SSSR Otd. Tekh. Nauk, Met. i Topliuo, No. 5, 97
(1959).
(10) EVANS, R., "The Corrosion and Oxidation of Metala,"
U.
2nd ed., Edward Arnold & Co., London, 1960.
(11) EVANS, R., AND DIVIES, D. E., J . Chem. SOC.,2607
U.
(19511.
s, M. E., r m CHEN,P. C . , J . Electroehem. Soe.,
134 (1951).
NAN~S, E., AND BRAKSS, J . Electrochem. Soc.
hl.
N.,

Conclusion

In all cases examined for corrosion of metals a single

322 / Journol of Chemical Education

2924a (1962).
(15) KINEYSKY, I., Zh. P ~ i k lKhim., 28, 1113 (1955).
A.
.

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