Schweitzer P.A. Fundamentals of corrosion. Mechanisms, causes, and preventative methods

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Corrosion Mechanisms 9

2.2.1 Dissimilar Electrode Cells

Dissimilar electrode cells may be formed when two dissimilar metals

are in contact or due to the heterogeneity of the same metal surface. The

Daniel cell is an example of the former. In practice, a copper pipe con-

nected to a steel pipe or a bronze propeller in contact with the steel hull

of a ship provides an example of this type of corrosion cell. This is often

referred to as galvanic coupling, in which the less noble metal becomes

the anode. Galvanic corrosion is discussed further in Chapter 3. A cold

worked metal in contact with the same metal annealed leads to a similar

situation (i.e., the cold worked metal remaining anodic). On the same metal

surface, such type of cell formation may result from dissimilar phases and

impurities, grain boundaries, differentially strained areas, and scratches

or abrasions. In a single crystal, the different crystal faces differ in their

electrochemical characteristics because of the difference in their atomic

orientation and, as a result, one crystal face tends to become anodic com-

pared to the others.

2.2.2 Concentration Cells

Concentration cells are formed when the electrodes are identical but are in

contact with solutions of differing composition. A salt concentration cell forms

when one electrode is in contact with a concentrated solution and the other

electrode with a dilute solution. On short-circuiting, the electrode in contact

with the dilute solution will be anodic (refer to Figure 2.1). The local variation

Cathode

CuCu

Anode

–+

Dilute CuSO

Conc. CuSO

FigurE 2.1

Salt concentration cell.

10 Fundamentals of Corrosion

of composition of the process stream inside the pipeline in a chemical plant

may lead to such a situation in practice.

A differential aeration cell forms when the identical electrodes are exposed

to solutions of identical chemical composition that differ in oxygen content,

which is illustrated in Figure 2.2. The electrode in contact with the less aer-

ated or oxygenated solution will act as the anode.

Differential aeration cell formation is quite common in operation and is very

important from the viewpoint of practical corrosion damage. A metallic bucket

half lled with water tends to corrode just below the water line because of the

lower oxygen concentration compared to the area just above it near the water

line. Corrosion damage invariably becomes pronounced underneath a corrosion

product or at crevices where oxygen availability is low. Formation of concentra-

tion cells of both kinds accounts for the initiation of pits (discussed in Chapter 3)

in stainless steels and in some other metals and alloys exposed to seawater.

2.2.3 Differential Temperature Cells

Differential temperature cells are formed when electrodes of the same metal,

each of which is at a different temperature, are immersed in an electrolyte

of the same initial composition. Such a situation may arise in practice in

components of heat exchangers, boilers, and similar heat transfer equip-

ment. Polarity developed in an electrode varies from system to system. For

a copper electrode in a copper sulfate solution, the electrode at the higher

temperature is the cathode; but for lead, the situation is just the reverse. For

iron immersed in dilute aerated sodium chloride solutions, the hot electrode

is initially anodic to the colder metal, but the polarity may reverse with the

progress of corrosion.

Anode

Cathode

Air

or O

Dilute NaClDilute NaCl

+–

FigurE 2.2

Differential aeration cell.

Corrosion Mechanisms 11

2.2.4 Oxygen Concentration Cells

The oxygen-reduction reaction that occurs in neutral or basic solutions, O

2Η

O + 4e

−

→ 40H

−

, plays a signicant role in many corrosion processes. It not

only contributes to sustaining a cathodic reaction, but also can induce one.

This occurs when substantial differences in dissolved oxygen content exist

at one area on the metal surface relative to another. The natural tendency

is to equal concentrations, and the means of achieving this by corrosion is

to lower the oxygen concentration at the region where it is the highest. The

oxygen-reduction reaction accomplishes this but the area where this occurs

becomes cathodic to the lower oxygen concentration region. Because of the

current ow created by this action, corrosion will occur at the anodic or low

oxygen concentration site.

2.2.5 Metal ion Concentration Cells

Metal ion concentration cells can also develop and fuel the corrosion process.

This situation arises when a signicant difference in metal ion concentration

exists over a metal surface. The tendency is to reach equilibrium ion con-

centration, and in a corrosive environment this is managed by putting more

metal ions into solution at the low-concentration area. This area becomes the

anode, and the current ow generated by this process can result in plating

out metal ions at the cathodic or high metal ion concentration region.

2.3 Potential–pH Diagrams

Potential–pH diagrams, also known as Pourbaix diagrams, are graphical rep-

resentations of the stability of a metal and its corrosion products as a func-

tion of the potential and pH (acidity or alkalinity) of the aqueous solution.

The potential is shown on the vertical axis and the pH on the horizontal axis.

Such diagrams are constructed from calculations based on the Nernst equa-

tion and the solubility data for various metal compounds. The potential–pH

diagram for an Fe–Η

Ο system is shown in Figure 2.3. In the diagram, the

horizontal lines represent pure electron transfer reactions dependent solely

on potential, but independent of pH:

Fe Fe e=+

+−2

(2.3)

Fe Fe e

+−3

(2.4)

12 Fundamentals of Corrosion

These lines extend across the diagram until the pH is sufciently high

to facilitate the formation of hydroxides, represented by vertical lines,

thereby reducing the concentration Fe

and Fe

ions. The boundary is

often set arbitrarily at the concentration of these ions at 10

−6

g-ions/liter,

which is indicative of a negligible dissolution or corrosion of the metal in

the medium.

The vertical lines in Figure 2.3 correspond to the reactions:

Fe 2H OFeOHH

()

(2.5)

Fe 3H OFeOHH

()

(2.6)

There is no electron transfer involved and the reactions are solely dependent

on pH.

The sloping lines in Figure 2.3 represent equilibria involving both electron

transfer and pH; for example:

Fe 3H OFeOHHe

()

+−

(2.7)

Fe 2H OHFeOHe

+= ++

−+−

(2.8)

1412108642

Potential E

(V)

–1.2

–0.8

–0.4

0.4

Hydrogen line

Oxygen line

0.8

1.2

Fe(OH)

HFeO

–

FigurE 2.3

Potential–pH (Pourbaix) diagram for Fe–H

O system.

Corrosion Mechanisms 13

The hydrogen and oxygen are also shown in the diagram by the dotted lines.

The hydrogen line represents the equilibria:

2H +2e=Hinacidsolutions

+–

(2.9)

22 2

HO eHOH in neutraloralkalinesolu

−

ttions

(2.10)

These two reactions are equivalent and their pH dependence of single elec-

trode potential is represented by:

EE pH

H/H

=−0059.

(2.11)

at pH = 0; that is, for [H

] = 1,

H+/H

= 0

and the slope is −0.059V. Similarly,

for oxygen equilibrium with water the corresponding reactions at lower and

higher pH are:

OHeHO

442++=

−

(2.12)

and

OHOe OH

244++=

−−

(2.13)

The pH dependence of single electrode potential is represented by:

EE pH

222

/H O

O/HO

=−0059.

(2.14)

at pH = 0,

O/HO

=1 226.

and at pH = 1 (i.e., for [OH

−

] = 1),

O/HO

= 0401.

Here again, the slope of the line is −0.059V. Water is stable in the area designated

by these two lines. Below the hydrogen line it is reduced to hydrogen gas, and

above the oxygen line it is oxidized to oxygen.

The potential–pH diagram shows three clear-cut zones:

1. Immunity zone. Under these conditions of potential and pH, iron

remains in metallic form.

2. Corrosion zone. Under these conditions of potential and pH, iron cor-

rodes, forming Fe

or Fe

or HFeO

−

14 Fundamentals of Corrosion

3. Passive zone. Under these conditions of potential and pH, protective

layers of Fe(OH)

form on iron and further corrosion of iron does not

take place.

Such diagrams can be used for:

1. Predicting the spontaneous direction of reactions

2. Estimating the stability and composition of corrosion products

3. Predicting environmental changes that will prevent or reduce

corrosion

With reference to Figure 2.3, corrosion prevention can be achieved by lower-

ing the electrode potential down to the zone of immunity, raising the elec-

trode potential up to the region of passivity, or raising the pH or alkalinity of

the solution so that a passive lm is formed.

There are, however, limitations in using such diagrams. The most impor-

tant of these is that they represent equilibrium conditions and hence cannot

be used for predicting the rate of a reaction. The tacit assumption that cor-

rosion products (oxides, hydroxides, etc.) lead to passivity may not always

be true because they may not always precipitate on the metal surface. The

possibility of precipitation of other ions such as chlorides, sulfates, and phos-

phates has been ignored. Finally, the pH at the metal surface may vary dras-

tically because of side reactions, and a prediction of corrosion based on the

bulk pH of the solution may be misleading.

2.4 Polarization

At an intermediate resistance in the circuit, some current begins to ow and

the potentials of both half-cell reactions move slightly toward each other.

This change in potential is called polarization. The resistance in the circuit

depends on a number of factors, including the resistivity of the media,

surface lms, and the metal itself. The relationships between polarization

reactions at each half-cell are represented in Figure 2.4. The intersection of

the two polarization lines (curves) closely approximates the corrosion cur-

rent and the combined cell potentials for the freely corroding situation.

Once the corrosion current is determined, the corrosion density can be

calculated by determining the surface area. Using Faraday’s laws, a corro-

sion rate in terms of metal loss per unit time can be determined. However,

polarization data can be more useful than just estimating corrosion rates.

The extent of polarization can help predict the type and severity of cor-

rosion. As polarization increases, corrosion decreases. Polarization may

Corrosion Mechanisms 15

be preferable to either cathodic or anodic reactions. Understanding the

inuence of environmental changes on polarization can offer insight into

controlling corrosion. For example, in the iron–hydrochloric acid exam-

ple, hydrogen gas formation at the cathode can actually slow the reaction

(increased current resistance) by blocking access of hydrogen ions to the

cathode site. This results in cathodic polarization and lowers the current

ow and corrosion rate. If oxygen is bubbled through the solution, the

hydrogen is removed more rapidly by combining to form water and the

corrosion rate increases signicantly. Although this is an oversimplied

view of the effects of oxygen, it does indicate that the degree of polar-

ization can be affected by changes in the environment, either natural or

induced.

There are three basic causes of polarization, termed activation, concentra-

tion, and potential drop. Potential drop is the change in voltage associated

with effects of the environment and the circuit between the anode and cath-

ode sites. It includes the effects of the resistivity of the media, surface lms,

corrosion products, etc.

2.4.1 activation Polarization

Activation polarization arises out of a slow step in the electrode reaction

for which an activation energy in the form of an increment in potential is

required for the reaction to proceed. This will best be illustrated by the

hydrogen evolution reaction.

The hydrogen evolution reaction consists of several steps as shown

in Figure 2.5. Either the electron transfer step (step 2) or the formation of

Current, i

corr

+ 2e

Fe Fe

+ 2e

–0.5

–0.3

–0.1

Potential, E

E (H

) Cathodic polarization curve

0.1

corr

E (Fe/Fe

)

Anodic polarization curve

FigurE 2.4

Polarization of iron in acid.

16 Fundamentals of Corrosion

hydrogen molecules (step 3) is deemed the slowest step in the reaction

sequence, and the rate of overall reaction will depend on how fast or slow it

proceeds. Therefore, to have a higher rate of reaction, expressed in terms of

increased current density, an increase in potential should be effected. The

relationship between reaction rate and change in potential (overvoltage) is

expressed by the Tafel equation:

ηβ

log

=±

(2.15)

where η

is overvoltage polarization (in volts), and β is a constant, called the

Tafel constant (also expressed in volts), and is usually on the order of 0.1V.

A graphical representation of Equation 2.15, as applied to the hydrogen

evolution reaction, with a β slope of 0.1V is shown in Figure 2.6. It can be

noted from the graph that 0.1V change in overvoltage can effect a tenfold

increase or decrease in the reaction rate.

Dissolution reactions (anodic) in corrosion are usually controlled by

activation polarization where the solution of ions is the probable rate-con-

trolling step. Hydrogen evolution reactions (cathodic reactions) are con-

trolled by activation polarization when the concentration of hydrogen ions

is high.

–

–2

–H

2–

Zinc

FigurE 2.5

Steps involved in hydrogen reduction reaction.

Corrosion Mechanisms 17

2.4.2 Concentration Polarization

A buildup or depletion of ions at the electrode surface as a result of reac-

tion will change the value of the electrode potential according to the Nernst

equation. For example, for a corroding zinc electrode, the concentration of

zinc will increase with dissolution in the vicinity of the electrode. The value

of a

oxid

in the equation increases, causing the electrode potential to shift in a

positive direction.

For the hydrogen evolution reaction, the higher rate of discharge of hydro-

gen ions at the electrode surface brings down the value of a

oxid

and the

electrode potential, according to the Nernst equation, shifts in a negative

direction. However, the rate of discharge of hydrogen ions at the electrode

surface depends on the diffusion of hydrogen ions from the bulk of the solu-

tion to the surface:

DnFC

(2.16)

1001010.10.01

–0.3

–0.2

–0.1

0.1

+ 2e

–

Noble

Active

Overvoltage η

Volts

0.2

Current Density

FigurE 2.6

Activation polarization curve of a hydrogen electrode.

18 Fundamentals of Corrosion

where i

is called the limiting diffusion current density (amp/cm

), D is the

diffusion coefcient for Η

ion, n is the number of electrons transferred, F is

the Faraday number, C is the bulk concentration of Η

ions in the solution,

and x is the thickness of the diffusion layer adjacent to the electrode surface

through which the concentration of the reacting species (H

ions) changes

from C in the bulk to zero at the electrode surface.

A mathematical expression for concentration polarization involves i

and

is given by:

=−













log

(2.17)

where η

is overvoltage due to concentration polarization (in volts). A graphi-

cal representation of the equation is shown in Figure 2.7.

It can be seen from the graph in Figure 2.7 as well as from Equation

2.17 that as i approaches i

, η

tends to innity. As evident from Equation

2.16, factors such as increasing velocity (smaller x), increasing temperature

(higher D), and increasing concentrations will increase the value of i

, that

is, a shift in the vertical position of the curve in Figure 2.7 more toward

the right.

There is no question of concentration polarization when the supply of

reacting species is abundant. Hence, in metal dissolution reactions, its effect

is negligible as the supply of metal atoms for dissolution is unlimited. On

the other hand, for a hydrogen evolution reaction, concentration polarization

becomes signicant in the solutions of low Η

concentration. More often, the

reduction process is controlled by a combined polarization — that is, activa-

tion polarization at lower reaction rates and concentration polarization at

higher reaction rates — as i approaches i

. A graphical representation of such

combined polarization is shown in Figure 2.8.

Increasing velocity

Increasing temperature

Increasing concentration

Log iLog i

(b)(a)

–

FigurE 2.7

(a) Concentration polarization curve for reduction process, and (b) effect of environmental

variations on concentration polarization curve.