1

ENVIORNMENTAL DEGRADATION OF MATERIALS 2013

A compilation from literature by Dr. Balakrishna Palanki

palankibalakrishna@yahoo.com

Mighty ships upon the ocean

suffer from severe corrosion.

Even those that stay at dockside

are rapidly becoming oxide

Alas, that piling in the sea

is mostly Fe2O3.

And when the ocean meets the shore

You´ll find there´s Fe3O4,

´cause when the wind is salt and gusty

things are getting awful rusty.

- T.R.B. WATSON

1. Basics of Corrosion Science and Engineering

Metals are valuable to man due to their useful properties such as strength and

formability. Unfortunately, metals are unstable and easily combine with other

elements to form compounds. For this reason, in nature, most metals occur in the

form of compounds as minerals. For example, the metal Al occurs in the form of

Al2O3. Cu exists as carbonate or sulphide. Just as work has to be done to pump water

to a higher level, work has to be done to extract a metal from its mineral. On the other

hand, just as water flows naturally from a higher to a lower level, the corrosion of a

metal is a spontaneous process. The metal would like to go back to its combined form

by way of corrosion. Fe rusts by forming oxide, silver stains by reacting with sulphur

in the environment and copper statues form patina of green carbonate. Hence

corrosion is extractive metallurgy in reverse. Corrosion may be defined as the

degradation of a material (metal, ceramic, polymer or composite) caused by its

unintentional chemical or electrochemical interaction with environment, starting at

its surface. Corrosion is a progressive oxidation reaction balanced by a

corresponding reduction reaction. Combustion of a metal powder is not considered as

corrosion. Corrosion increases with exposure time, temperature and environmental

and several other factors. Only iron and steel rust. Other metals corrode!

1.1. Corrosion of iron

The importance of iron is seen from the fact that the quantity of iron used in the world

is 7 times that of all other metals put together. The corrosion (oxidation) of iron in

acids yields soluble salts and hydrogen gas and can be expressed as half reactions:

Fe → Fe++

+ 2e

2 H+ + 2 e → H2

In the above the oxidation of Fe to Fe++

is balanced by the reduction of H+

to H2 gas.

2

The term OILRIG indicates that Oxidation is Loss of electrons and Reduction is Gain

of Electrons.

The half reactions may be combined as below:

Fe + 2 H+ = Fe

++ + H2

The corrosion of metals can also occur in fresh water, seawater, salt solutions, and

alkaline or basic media. In almost all of these environments, corrosion occurs

importantly only if dissolved oxygen is also present. Water solutions rapidly

dissolve oxygen from the air, and this is the source of the oxygen required in the

corrosion process. The most familiar corrosion of this type is the rusting of iron when

exposed to a moist atmosphere.

Fe → Fe++

+ 2e

½ O2 + H2O + 2 e → 2 OH-

The half reaction above may be combined to get

4 Fe + 6 H2O + 3 O2 → 4 Fe(OH)2 ↓

In this equation, iron combines with water and oxygen to produce an insoluble

reddish-brown corrosion product that falls out of the solution, as shown by the

downward pointing arrow.

During rusting in the atmosphere, there is an opportunity for drying, and this ferric

hydroxide dehydrates and forms the familiar red-brown ferric oxide (rust) or Fe2O3, as

shown below:

2 Fe(OH)3 → Fe2O3 + 3 H2O

Experiment to demonstrate the importance of water and oxygen in corrosion:

Take three test tubes A, B and C. Keep one clean nail in each. Leave A sealed with a

cork only with air, using anhydrous calcium chloride at the bottom as dessicant to

keep air dry. Pour water in B to partly submerge the nail and seal In C, submerge the

nail fully in distilled water, pour oil (which floats on the water) to prevent ingress of

air into water and seal the tube. After a few days, only the nail in B is found to show

rust. Corrosion of iron or steel can be prevented by removing contact with air and

moisture.

1.2 Metal electrodes

When a metal is placed in its solution, a small part of the metal dissolves or ionises

and the metal ions (which no longer have metallic properties) join the solution. This

process of ionization leaves a negative charge on the metal. The left over negative

charge in the metal gives the metal a potential.

The transfer of the metal from the metal into its solution is called oxidation or

corrosion. Oxidation liberates electrons.

3

M → M+ + e

or

M → M++

+ 2e

The number of electrons released into the metal depends on its valency. It is usually

denoted by the letter z. z is 1, 2 or 3 etc. The place where oxidation or corrosion takes

place is called the anode.

Soon the reverse reaction begins.

M+ + e → M

The reverse reaction s called reduction. Reduction consumes electrons. The place

where reduction takes place is called cathode. The rate of oxidation and the rate of

reduction become equal. We say that dynamic equilibrium has been reached. It is the

lowest free energy state of the reaction. At equilibrium, the system is stable and there

is no driving force for any change from that state. When there is no net reaction, that

is, when the equilibrium is reached, the potential displayed by the metal is called

equilibrium electrode potential. If you increase the potential by some means, more

oxidation will take place. If you decrease the potential by some means, more

reduction will take place.

The tendency to oxidise is not the same for different metals. Some metals tend to

oxidise readily, while some are reluctant to oxidise. The metals that oxidise readily

are called base metals. For example, you can not hold a piece of sodium in your hand.

It burns instantly. Some other metals like calcium or magnesium readily react with

water. Some other metals such as iron or zinc react with acid. On the other hand, there

are metals like gold which do not react readily. Such metals are called Noble metals.

Noble metals as well as base metals all exhibit a potential when in contact with their

solution, though the tendency to oxidise are different. Based on their reactivity, the

metals can be arranged in an order.

The equilibrium potential exhibited by a metal in contact with its solution when the

conditions are standard, is called standard electrode potential. By standard conditions,

we mean a temperature of 25 C (298 K), a pressure of one atm and a pH of zero or a

solution strength of one molarity. Metals arranged in the order of the standard

electrode potentials are called the electromotive or emf series. The series has base

metals at one end and noble metals at the other end.

A potential can not be measured in absolute terms. It can be measured as a potential

difference. Just as you take mean sea level as zero and measure heights relative to it,

we take the hydrogen gas – hydrogen ion equilibrium under standard conditions to be

zero volts and measure potentials of all other metals relative to it.

2 H+ + 2 e → H2 Zero volts.

The above equation is the same as

4

H+ + e → (1/2) H2 Zero volts.

pH is defined as

pH = - log [H+]

When [H+] is one mole per litre of the solution, log [H

+] is zero.

The reference electrode, which is used for potential measurements, is called the

standard hydrogen electrode or SHE, the potential of which has been taken to be 0 V.

Zn → Zn+ + e has a standard electrode potential of minus 0.76 volts.

Cu → Cu++

+ 2e has a standard electrode potential of plus 0.334 volts.

Suppose one wishes to measure the heights of poles in a field, and he is not able to see

the ground level, because of tall grass. Still, he can measure the heights by taking the

tip of one pole to be at a height of zero meters. Then he measures the heights of the

tips of other poles relative to this. Then the poles shorter than the standard pole will

have negative heights. The poles longer than the standard pole will have positive

heights.

The standard hydrogen electrode is analogous to the standard pole above.

Reactions such as Zn → Zn++

+ 2e or

2 H+ + 2 e → H2 are called Half reactions.

In contrast, a reaction such as

Zn + 2 H+ → Zn

++ + H2 is a full reaction, the half reactions of which are as above.

The full reaction is obtained by adding the half reactions. The potential of the full

reaction is also the addition of the potentials of the half reactions.

In real situations, there are no half reactions. All are full reactions. One type of half

reaction is balanced by another type of half reaction. For example, oxidation half

reaction has to be balanced by a reduction half reaction. This is because electric

charge has to be conserved.

1.3 Nernst Equation

The potential of a metal in contact with its solution can be changed by changing the

concentration, pressure or temperature. The way a potential changes when these are

changed is expressed as Nernst equation.

To determine the potential of a system in which the reactants are not at unit activity,

Nernst Equation is employed. For the chemical equation represented by

A + B = C + D

BA

DC

Haa

aa

zF

RTEE ln0

5

which is the same as

DC

BAH

aa

aa

zF

RTEE ln0

where EH is the cell potential, E0 the standard cell potential, R is the gas constant, T is

the absolute temperature, z is the number of electrons transferred, F is the Faraday

Constant, aA , aB etc are the activities (concentrations) of the species. Both E0 and EH

are equilibrium potentials.

Converting log to the base 10,

BA

DC

Haa

aa

zF

RTEE log3.20

At 25°C, and substituting F by 96485 coulombs, R by 8.314 J.mol-1

.K-1

BA

DC

Haa

aa

zEE log

0592.00

For the half reaction Zn → Zn+ + e, z = 1

For the half reaction, Cu → Cu++

+ 2e, z = 2

We have seen how the potential can by increased or decreased by changing

temperature, pressure or concentration. There are other important ways of changing

the potential.

One way is to connect the metal to a more noble metal. For example, you can connect

zinc electrode with a copper electrode. The potential of the zinc electrode increases

from minus 0.76 volts to a greater value. The potential of the copper decreases from

plus 0.334 volts to a lesser value.

Since the potential of zinc has been increased (by connecting to copper),

oxidation rate of zinc increases. Since the potential of copper has been decreased (by

connecting to zinc), the reduction rate of copper increases. When the copper and zinc

remain connected electrolytically, the zinc corrodes rapidly. The oxidation half

reaction of zinc is balanced by the reduction half reaction of copper.

1.4 Electrolyte and Electrolytic cell

An electrolyte is a solution containing positive and negatively charged ions. For

example, an electrolytic solution of NaCl consists of Na+ and Cl

- ions, in water. The

electrolyte is usually an aqueous solution. Molten salts not containing water are also

electrolytes. An electrolytic cell is a combination of an anode and a cathode, both

immersed in an electrolyte. The anode and cathode may not be connected to each

other, or may be connected to a load or may be shorted. On this basis, three types of

electrolytic cells are described below.

6

When the electrodes are not in contact with each other, the electrolytic cell is

called battery ( a dc power source) and can be used for supplying electrical energy to

a load, such as a torch light or engine starter in an automobile. When a dc power

source is connected to an electrolytic cell, electrolysis or electroplating takes place.

Electro forming and elctrowinning are also achieved by connecting a dc power

source to the electrolytic cell. Electrolysis of water yields hydrogen gas at the cathode

and oxygen gas at the anode. When the electrodes of electrolytic cells are joined

together by a conducting wire, it becomes a corrosion cell. A short circuited

electrolytic cell is a corrosion cell. Zinc and copper in their solutions and connected

by a conducting wire form a corrosion cell, where the zinc undergoes oxidation and

copper undergoes reduction. The electrons released by zinc changing to zinc ions are

taken up by copper ions to become copper metal. The oxidation reaction of zinc is

balanced by the reduction reaction of copper. In other words, zinc corrodes and

protects the copper.

It is possible to change the potential of the copper and zinc simply by connecting

them. It is also possible to change the potential by connecting to the pole of a battery

or dc power source. If you connect the metal to the positive pole, the potential of the

metal increases and oxidation of the metal increases. If you connect it to the negative

pole, reduction of the metal increases.

When the potential is changed, we say that there is polarisation. Polarisation is the

difference between actual applied voltage and the equilibrium potential. When the

potential is at equilibrium value, there is no net oxidation or net reduction. When the

potential is changed, there is an increase in oxidation or reduction rates. The excess

potential or the overvoltage is related to the consequent current by Tafel Law.

1.5 Corrosion reactions

The corrosion of a metal which is an oxidation process taking place at the anode may

be represented by the half reaction:

Zn → Zn++

+ 2e or

Fe → Fe++

+ 2e

Cu → Cu++

+ 2e

For corrosion to be possible, the above half reactions taking place at the anode need to

be necessarily balanced by one or more reduction half reactions taking place at the

cathode. Electric charge can not be accumulated and has to be conserved. In other

words, the electrons released by the anodic oxidation half reaction travel through the

conducting wire, reach the cathode and get consumed in the cathodic reduction half

reaction. Note that ions travel through the electrolyte, constituting an ionic current,

while electrons travel through the metal constituting an electronic current.

Some common cathodic reduction half reactions responsible for causing corrosion of

a metal are given below:

In acidic solutions, with no oxygen present 2 H+ + 2 e → H2

7

In acidic solutions, with oxygen present ½ O2 + 2 H+ + 2 e → H2O

In neutral or alkaline solutions ½ O2 + H2O + 2 e → 2 OH-

The reduction of another metal ion may also serve as a balancing half reaction:

Cu++

+ 2e → Cu

2. Thermodynamics of Corrosion

2.1 Energy can neither be created nor destroyed.

All spontaneous changes occur with a release of free energy from the system to the

surroundings at constant temperature and pressure. In nature, water runs downhill, hot

coffee goes cold and spinning tops slow down. Work has to be performed for

achieving changes that are not spontaneous. Corrosion is a spontaneous

electrochemical process, where free energy is released and the metal returned to its

stable (combined) state. The driving force for corrosion reaction is chemical energy,

energy stored in chemical bonds of substances or internal energy. Free energy is the

portion of internal energy available for powering engines or causing corrosion

reaction. We can force an electrochemical reaction to proceed in the reverse direction

by connecting a source of current to the electrodes, as in electroplating or electro-

winning.

2.2 Electrolysis of water

In the electrolysis of water, hydrogen gas is evolved at the cathode and oxygen gas is

evolved at the anode. Electrolysis of water takes place when the applied voltage

exceeds 1.2 V as per the following equations:

At the anode: H2O = ½ O2 + 2 H +

+ 2 e (E0 = 1.23 V)

At the cathode: 2 H+ + 2 e = H2 (pH < 7) (E0 = 0 V)

The electrons released by the anodic reaction are consumed by the cathodic reaction.

Note that other anodic reactions are possible, with other values of equilibrium

potentials. Correspondingly, the electrolytic voltage requirement also changes.

The potential of a metal in contact with its solution M = M++

+ 2e is given by the

Nernst equation:

E = E0 - 2.303 [RT/(nF)] log (M+/M)

In the case of hydrogen gas and hydrogen ions, for the equilibrium represented by

2 H+ + 2 e = H2 (pH < 7) (E0 = 0 V)

the Nernst equation may be written as

E = E0 - 2.303 [RT/(nF)] log [1/(H+)]

8

At a temperature of 25 C, E0 has the value zero and 2.303 [RT/(zF) has the value

0.059

E = 0.059 log (H+)

By definition, pH = - log (H+)

E = - 0.059 pH

This is the equation of a straight line, with negative slope.

If the potential is smaller than that indicated in the above expression, hydrogen gas

evolves.

In the case of oxygen gas and oxygen ions,

½ O2 + 2 H +

+ 2 e = H2O (E0 = 1.23 V)

the equation becomes

E = 1.23 + 0.059 log (H+)

E = 1.23 - 0.059 pH

This is the equation of a straight line, parallel to the hydrogen gas / hydrogen ion

equilibrium line.

If the potential is larger than that indicated in the above expression, oxygen gas

evolves. Thus, for the electrolysis of water, the potential difference should therefore

be at least 1.23 volts.

Note that the equilibrium potential becomes the standard equilibrium potential when

pH = 0.

2.3 Pourbaix Diagrams:

The application of thermodynamics to corrosion phenomena have been generalized by

means of potential – pH plots. These are called Pourbaix diagrams, after Dr. Marcel

Pourbaix who first suggested their use. Such diagrams are constructed from

calculations based on the Nernst equation and solubility data for various compounds.

In the diagram for iron, for example, it is possible to delineate areas in which iron,

iron hydroxide, ferrous ions, etc., are thermodynamically stable. That is, these forms

represent states of lowest free energy.

The main uses of these diagrams are (1) predicting the spontaneous direction of

reactions, (2) estimating the composition of corrosion products, and (3) predicting

environmental changes that will prevent or reduce corrosive attack. For example, the

large region labeled Fe indicates that iron is inert under these conditions of potential

and pH.

9

The electromotive series only takes into account, those electrochemical equilibria

involving metals and their simple cations (Mz+

), in other words, the only type of

reaction considered is

M = Mz+

+ 2 e

which only contains ‘e’ terms in addition to the metal and metal ion and so is only

potential dependent. There are several other relevant reactions such as

M + z H2O = M(OH)z + z H+ + 2 e

in which the metal hydroxide can represent a protective film if it is closely packed,

tightly adherent and has a sufficiently low solubility. Another reaction is

M + 2z (OH) = MO2 z-

+ z H2O + 2 e

which represents corrosion of the metal in an alkaline medium. These equilibria are

potential and pH dependent due to the presence of e, H+ or OH

-. Reactions dependent

only on pH can also occur:

Mz+

+ z (OH-) = M(OH)z

MOzz-

+ z H+ = M(OH)z

The line a-b in the figure represents the hydrogen-hydrogen ion equilibrium. EH = E0

= 0 when pH = 1. If a point is below the line a-b, the reaction proceeds from left to

right i.e., H2 evolves. If it is above the line, H2 ionises till equilibrium is reached.

Similarly when the point (EH, pH) is above c-d, O2 evolution is possible.

If a potential difference of 1.2 volts is maintained between the electrodes, at the anode

H2O = ½ O2 + 2 H+ + 2 e and at the cathode 2 H

+ + 2 e = H2.

All metals corrode, though the rates and extents vary. As corrosion proceeds, the

concentration of metal ions increases. If the equilibrium concentration of the ions is

small (10-6

gram ions/litre), the metal is taken to be immune to corrosion and the lines

in the EH - pH diagram are drawn for this concentration.

1. Fe = Fe++

+ 2e

2. Fe++

= Fe+++

+ e

3. Fe++

+ 3 (OH-) = Fe(OH)3 + e

4. Fe+++

+ 3 H2O = Fe(OH)3 + 3H+

5. Fe + 3 H2O = Fe(OH)3 + 3H+ + 3e

6. Fe + 2 H2O = FeO2H- + 3H

+ + 2e

7. FeO2H- + H2O = Fe(OH)3 + e

10

Limitations: The thermodynamic approach excludes kinetic data. Only those

equilibria concerned with metal in conjunction with water are used in constructing EH

- pH diagrams so that only OH- is considered as the precipitating or complexing ion.

In practical corrosion, there are Cl-, SO4

=, PO4

= etc.

2.4 Passivity

Metals that normally fall victim to corrosion will sometimes exhibit a passivity to

corrosion. Passivity is a phenomenon in which a metal (or alloy) exhibits a much

higher corrosion resistance than its position in the electrochemical series would

indicate. This happens when the initial corrosion product, if tenacious and nonporous,

actually shields the metal against further corrosion. Iron corrodes (dissolves) in dilute

nitric acid but not in concentrated acid due to passivation. Aluminium, chromium and

zirconium which are at the base end of the electrochemical series, display excellent

corrosion resistance because the initial oxide films formed (Al2O3, Cr2O3, ZrO2)

protect the metal from further attack. Chromium used as alloying element in SS

behaves the same way. Passivity is the characteristic of a metal exhibited when the

metal does not become active in the corrosion reaction. Passivity is caused by the

build up of a stable tenacious layer of metal oxide on the surface of the metal. This

oxide layer is formed by corrosion on a clean metal surface, where the corrosion

products are insoluble in the particular environment to which the metal is exposed.

Once the layer, or film, is formed, it acts as a barrier separating the metal surface from

the environment. For further corrosion to occur, the reactants must diffuse through

the oxide film. Such diffusion is very slow or non-existent, thus corrosion either

decreases markedly or stops. Metals such as zirconium, chromium, aluminum and the

stainless steel form thin, tenacious oxide films when exposed to the atmosphere or to

pure water at room temperature. In some cases, the film is extremely thin and may be

invisible to the unaided eye, but it is still very effective in giving these metals a

marked passivity.

2.4 The Pilling – Bedworth Ratio

If the volume of the scale produced is greater than the volume of metal ionised in

producing it, the scale tends to be formed in lateral compression, and can be compact

and protective – assuming the compressive forces are not so great to disrupt it. On the

other hand, if the volume of scale is less than the volume of metal ionised in

producing it, the scale tends to be formed in lateral tension, and since most scales are

not mechanically strong in tension it will probably be cracked and crazed, thus

affording incomplete coverage of the metal and therefore be nonprotective.

Molecular weight of the scale > Molecular weight of the metal

Density of the scale Density of the metal

(M/D) > (m / d) Or (Md / Dm) > 1 for scale to form in compression

For Al – Al2O3 the ratio is 1.275 so the scale will be formed in compression. In

general, metals with volume ratios of less than 1 form nonprotective oxides, as do

those with very high volume ratios (2 to 3), but this ratio was intended only as an

11

empirical criterion and does not encompass other properties more important in

determining oxidation resistance. To be protective to oxygen reaction, an oxide must

possess good adherence, a high melting point, a low vapour pressure, good high

temperature plasticity to resist fracture, and low electrical conductivity or low

diffusion coefficients for metal ions and oxygen. For cyclic temperatures, the metal

and oxide should possess similar coefficients of expansion.

So far it has been assumed that when a corrosion reaction is balanced by a reduction

reaction both the anode and the cathode are polarised to the same potential Ecorr. In

practice, there is IR drop due to the resistance of the metal & electrolyte.

oxygen evolution

and acidification

hydrogen evolution

and alkalization

water thermodynamically stable

12

13

3. Kinetics of Corrosion

3.1 The Faraday

From Faraday’s laws of electrolysis it is known that the mass of a substance liberated,

deposited or dissolved in any electrochemical reaction is proportional to the charge or

current multiplied by time.

dm = I dt

where dm is the mass involved, I is the current and dt is the time.

Valency is a measure of the combining power of an atom. Gram equivalent for acids

means the amount of material capable of releasing one gram ion hydrogen ions. For

example for HCl, it is the same as molecular weight, 36.5 g and for H2SO4; it is half

the molecule weight i.e. 49 g.

Consider M M++

+ 2 e

6 x 1023

correspond to

Atoms of M 2 x 6 x 1023

electrons

or

1 gm atom 2 x 6 x 1023

x 1.6 x 10-19

coulombs

( or 2 gm equivalents in this case)

1 gm equivalent 6 x 1023

x 1.6 x 10-19

coulombs

96487 coulombs (1 faraday)

3.2 Corrosion current as a measure of corrosion rate

Corrosion rate may be measured as thickness loss or weight loss per unit time per unit

area. It may also be measured using current density.

14

From Faraday’s first law of electrolysis it is known that the mass of a substance

liberated, deposited or dissolved in any electrochemical reaction is proportional to the

charge or current multiplied by time.

dm I dt

where dm is the mass involved, I is the current and dt is the time.

According to the first law of Faraday, the rate of a reaction can be measured by the

current density. Current density is denoted by the small i.

Mass per unit time (dm/dt) current ( I )

Mass per unit area per unit time current density ( i )

According to Faraday’s second law, for the same charge, the mass deposited is

proportional to the equivalent weight of the element.

3.3 Dissimilar Metal Coupling Effect:

The more active metal will become the anode of the couple in a wet corrosion cell.

The more noble metal will become the cathode of the couple in a wet corrosion cell.

The corrosion rate of the more active metal will be accelerated, while that of the more

noble metal is retarded (the cathode may still corrode, depending upon the magnitude

of the cathodic polarization induced by coupling of the more reactive metal).

3.4 Corrosion Rate, Temperature Effect

The order of galvanic series may reverse under certain circumstances, for example,

Zn and Fe at elevated temperatures in some potable water systems. Like most other

chemical reactions, corrosion rates increase as temperature increases. Temperature

and pressure of the medium govern the solubilities of the corrosive species in the

fluid, such as oxygen (O2), carbon dioxide (C02), chlorides, and hydroxides. A rule of

thumb is that the reaction rate doubles with a 20°F to 50°F temperature rise. This

linear increase with temperature does not continue indefinitely due, in part, to a

change in the oxide film. The environment plays an important role in a corrosion

process. Corrosion rates in tropical countries and seaside places are higher than in

cold countries and inland places.

3.5 Corrosion rate, other effects

Cathode to anode area ratio: Greater area ratio will result in a greater accelerating

factor in the rate of galvanic corrosion

Velocity effect: When water velocity is extremely high, the impact of the water tends

to remove the protective oxide layer and some of the metal under it (erosion), thus,

exposing more metal to corrosion. Water velocities of 30 to 40 ft per second are

usually considered to cause erosion.

15

Composition Increasing concentration of corrosive species (e.g. H+ ions) usually

increases corrosion rate (except in passivation).

Microstructure Cold worked regions of a metal are more susceptible to corrosion

than the annealed regions.

Alloying Alloys tend to have higher corrosion rates than their pure metals (except

when passive films form – stainless steels).

3.6 Passivity

With passive metals, the immunity to corrosion depends upon the properties of the

oxide film such as:

(a) thickness and impermeability to media

(b) adherence of the film to the base metal

(c) resistance of the film to chemical attack

(d) mechanical strength

(e) the ability to repair defects which may arise in the film.

Such oxide films are produced even on exposure to air, and although invisible while

in contact with the base metal, have been separated and studied. The film produced on

iron in many environments does not fulfil the above conditions, consequently

corrosion occurs. Chloride and sulphate ions are especially liable to cause film

breakdown, but chromates and phosphates promote repair.

On the other hand, aluminium and chromium form films of the impervious type and

this property is usually conferred on their alloys, so long as reducing conditions do

not prevail. Zirconium and titanium also exhibit passivity.

Passivation due to Chromium

Chromium is the chief alloying element in iron and steel for inhibiting corrosion. This

resistance is not due to the inertness of the chromium, for it combines with oxygen

with extreme rapidity, but the oxide so formed is very thin and stable, continuous and

impervious to further attack. This property is, fortunately, conferred upon its solid

solution in iron, becoming very marked as the amount exceeds 12 percent in low

carbon steels. Thus in oxidising environments, such as nitric acid, the high chromium

is initially attacked at the same rate as ordinary plain steel, but it rapidly builds up an

oxide film, known as a self-healing passive film, which efficiently protects the

underlying metal.

In the corrosion of some metals, there is a danger of break-away corrosion, which

means that after a period of a low rate of corrosion, the film of oxide becomes

discontinuous and flakes away from the surface and the corrosion rate accelerates,

e.g., zirconium in water at 400 C.

3.7 Tafel Equation

16

When there is no applied voltage, the net current density is zero. When the voltage is

increased, there is net oxidation current density. When the voltage is decreased, there

is net (reduction) current density. The relationship between the applied voltage and

the current is called Tafel Law.

Thus the rate of dissolution of a metal can be represented by an electric current.

Reaction rate per unit area of electrode current density ( i ). At equilibrium there is

no net current. Deviation of the applied potential from the equilibrium value is

responsible for driving the current. The deviation, given by E - E0 or E - EH is called

overpotential or over voltage and represented by the Greek letter

= a + b log i (Tafel equation)

3.8 Exchange Current Density

For a metal M, in equilibrium with its own ions M = M+ + e, the rate of forward

reaction (oxidation M = M+ + e) is exactly balanced by the rate of the reverse reaction

(reduction M+ + e = M), i.e.

ia = ic = io

where i is the rate of reaction expressed in terms of current density and subscripts, a

and c, represent the anodic (oxidation) and cathodic (reduction) processes. io is called

the exchange current. It is to be noted that exchange current is always associated with

an equilibrium condition, that is, the rate of forward (anodic) reaction equals that of

the reverse (cathodic) reaction. The net current is zero.

It has been possible to measure the exchange current density of a half reaction. The

values of exchange current density and standard electrode potential can be shown as a

point on a graph. Then the Tafel lines can be drawn beginning with this point for

M/M+, H2/2H+ etc.

3.9 Evan’s Diagram

When a metal is in contact with its solution, it exhibits a potential. A metal in contact

with its solution and exhibiting a potential is called a ‘metal electrode’. When the

metal and its ions are in equilibrium, the metal exhibits an equilibrium potential.

Under standard conditions, the equilibrium potential is E0. Under nonstandard

conditions, the equilibrium potential is EH and given by Nernst equation. By

equilibrium, we mean that both oxidation and reduction are taking place at equal

rates, and there is no net reaction. When the potential is changed deliberately away

from EH, the equilibrium is disturbed and a net reaction takes place. If the potential is

increased, there is net oxidation. If the potential is decreased, there is net reduction.

The potential of the metal may be increased by connecting it to the positive pole of a

battery or dc source or by connecting to a more noble metal. The potential of the

metal may be decreased by connecting to the negative pole of a dc source or by

connecting to a less noble metal. The increase in potential is the basis of anodic

protection (in the case of metals that can form an impervious adherent protective

oxide film layer). The decrease in potential is the basis for cathodic protection. The

increase or decrease of potential, away from equilibrium potential, is called

‘overpotential’ or ‘polarization’. When an electrode is polarized, there is a net current.

The current is a result of the polarization. Yet, usually, the polarization voltage is

17

shown on the Y axis while the current density is shown on the X axis. The graph

shows a point corresponding to standard electrode potential and the corresponding

exchange current density for an oxidation - reduction equilibrium. Beginning at this

point, one line, sloping upwards indicates oxidation. Another line, sloping downwards

indicates reduction. The excess over the standard potential is the over-potential. It is

positive for oxidation and negative for reduction.

A copper electrode exhibits a standard potential of +0.334 V and a zinc electrode

exhibits a standard potential of -0.76 V. When a copper electrode is joined to a zinc

electrode (shorted), the potential reaches an intermediate value. Then it is called a

‘mixed electrode’. The difference between the new value and the previous value is

polarization. The oxidation of zinc is balanced by the reduction of copper, both

processes taking place at the new potential. Here the oxidation and reduction are

taking place in different metals. Zinc will continue to dissolve or corrode to the last

atom.

You have to distinguish between corrosion current and current density. Current

densities enable comparison of different corrosion rates. However, since there can be

no accumulation of charge, it is the currents involved (in oxidation and in reduction)

that are to be balanced and not the current density. That is, sum of anodic currents =

sum of cathodic currents.

A mixed electrode is an electrode or metal sample that is in contact with two or more

oxidation reduction systems. An example is zinc immersed in hydrochloric acid.

Single electrode (zinc in acid) reactions:

Zn = Zn2+

+ 2 e- anodic reaction

2 H+ + 2 e

- = H2 cathodic reaction

Each has its own half-cell electrode potential and exchange current density.

The two half-cell electrode potentials cannot coexist separately on the same

electrically conductive surface (metal).

Each must polarize or change potential to a common intermediate value, the mixture

of the half-cell electrode potentials, most often called the mixed potential. The

common potential is called Corrosion potential (Ecorr) and the common current density

is called corrosion current density (icorr).

18

4. Types of Corrosion

4.1 Direct Chemical Attack

This form of corrosion is essentially ordinary chemical attack. The corrosive agent

such as an acid pickling solution used to clean steel surfaces, dissolves the surface

uniformly without the formation of protective layers and without selective attack of a

certain phase or component. The attack continues at an almost constant rate.

Fe + 2H+ Fe

++ + H2 (pickling of steel)

This reaction differs from the indirect process of electrochemical corrosion in which

iron is dissolved as the anodic reaction and hydrogen takes part in a separate cathodic

reaction.

It is uniform chemical reaction across entire metal surface. Some areas are anodic

and some are cathodic, but these change with time giving uniform overall corrosion,

usually producing a scale or deposit. Ex. General rusting of steel, tarnishing of silver.

Direct Chemical Attack is also fundamentally electrochemical in nature. However, no

current flow is detectable, nor are there any definite anodic or cathodic areas. The

theoretical rate of a chemical attack can be affected by the formation of a protective

film on the metal surface, through secondary reactions involving the products of

corrosion, and the mechanical removal of protective films, such as by erosion, flexing

of the metal surface, or by temperature changes.

In (indirect) electrochemical corrosion, there are electric currents through parts of the

metal in which corrosion is not taking place. In direct corrosion, on the other hand,

the electrical effects involved occur in the immediate vicinity of the reaction.

4.2 Galvanic corrosion:

19

Galvanic corrosion occurs when two dissimilar metals are in electrical contact with

each other and are exposed to an electrolyte. A less noble metal, say zinc, will form

the anode and dissolve, whereas the more noble metal, say copper will act as the

cathode. Depending on the nature of the corrosive environment, the cathodic reaction

may occur either by hydrogen evolution or oxygen absorption.

Differential Couples:

From the structural point of view it is usually undesirable to connect copper and iron

together where there is going to be any moisture present. A similar generalisation

applies to carbon and iron in any sort of aqueous environment – particularly a low pH

one – not only from the thermodynamic stand point of the relative nobility of the

carbon but also because of the kinetic fact that most forms of carbon have low

hydrogen over-potential and therefore act as very efficient cathodes. A buried ferrous

pipe line such as a gas main, lying in any form of carbonaceous backfill, especially in

wet conditions will have a poor chance of a long life. Intergranular corrosion is

another example. Differences in concentration, oxygenation, stress or even in

temperatures will result in corrosion cell.

4.3 Differential Aeration (Concentration Cell Corrosion)

This occurs when one part of a metal is exposed to a different air concentration from

the other parts. This causes a difference in potentials between differently aerated

areas. Areas on a metal surface where the oxygen concentration is low are anodic.

Areas where the oxygen concentration is high are cathodic. Different aeration

accounts for the corrosion of metals partially immersed in a solution just below the

water line. Similarly areas on iron covered by droplets are screened from access of

oxygen and become anodic in respect to the other areas freely exposed to air. Metals

exposed to aqueous media corrode under blocks of wood or pieces of glass which

screen that portion of the metal from oxygen access. The presence of crevices as

formed by improper gasket fit can cause corrosion.

4.4 Crevice corrosion: Crevice Corrosion is the attack on the surface of a metal

partly shielded from contact with the corroding fluid, usually by non-metallic

materials. Typically, this is a concentration cell effect, the shielded area being anodic.

Concentration cells form due to differences in metal ion concentration in the

electrolyte between two regions of (same) metal piece. Metal in contact with more

concentrated electrolyte becomes “Cathode”; metal in contact with more dilute

solution becomes anode and corrodes. Electrons flow from the low-oxygen area on

the metal which acts as the anode to the high-oxygen area on the metal which acts as

the cathode.

To reduce crevice corrosion, use welded rather than bolted or riveted joints. Use non-

absorbing gaskets. Remove accumulated deposits. Design vessels with complete

drainage without stagnant areas.

4.5 Uniform Corrosion: Uniform unlocalized attack is the name given to the

corrosion attack which affects the whole surface of a metal simultaneously.

20

4.6 Pitting: Pitting is a localized accelerated attack resulting in the formation of

cavities or pin holes around which the metal is relatively unattacked – Pitting is

usually the result of the breakdown or cracking of the protective film on a metal at

specific points. This gives rise to the formation of small anodic and large cathodic

areas. Breakdown of the protective film may be caused by mechanical factors or

certain chemicals. Stainless Steel and Aluminium show pitting in halide solutions.

Pitting corrosion is a localized form of corrosion by which cavities or "holes" are

produced in the material. Pitting is considered to be more dangerous than uniform

corrosion damage because it is more difficult to detect, predict and design against.

Corrosion products often cover the pits. A small, narrow pit with minimal overall

metal loss can lead to the failure of an entire engineering system. Pitting corrosion,

which, for example, is almost a common denominator of all types of localized

corrosion attack, may assume different shapes.

Pitting is initiated by:

a. Localized chemical or mechanical damage to the protective oxide film; water

chemistry factors which can cause breakdown of a passive film are acidity,

low dissolved oxygen concentrations (which tend to render a protective oxide

film less stable) and high concentrations of chloride (as in seawater)

b. Localized damage to, or poor application of, a protective coating

c. The presence of non-uniformity’s in the metal structure of the component,

e.g. nonmetallic inclusions.

Pitting corrosion can produce pits with their mouth open (uncovered) or covered with

a semi-permeable membrane of corrosion products. Pits can be either hemispherical

or cup-shaped.

Apart from the localized loss of thickness, corrosion pits can also be harmful by

acting as stress raisers. Fatigue and stress corrosion cracking may initiate at the base

of corrosion pits.

4.7 Intergranular Corrosion Precipitation of certain compounds at grain boundaries

leaves the solid solution adjacent to the grain boundary impoverished in one

constituent. The depleted zone is anodic with respect to the precipitated compound,

as well as center of the grain, so that it will be preferentially attacked by the corrosive

environment (CuAl2 ppt in Al-Cu alloys, weld decay in austenitic stainless steel).

4.8 Erosion – Corrosion: Erosion-corrosion is the attack caused by the combined

effect of corrosion and wear produced by the turbulent flow of gases, vapours and

liquids, and also by the rubbing action of solids over a metal surface. Erosive action

removes the protective films. It is frequently encountered in condenser tubes, piping,

agitators, or vessels in which streams of liquids or gasses emerge from an opening and

strike the side wall.

Erosion Corrosion, as the name implies, occurs when the corrosion products which

would normally afford a protective film are scoured off by moving fluids, particularly

21

if the fluids contain abrasive materials. The erosion will expose clean metal, and

develop a physical pattern so obviously a result of erosion that the corrosive factor

may not be recognized.

Cavitation or Impingement Attack is a process which is very similar to erosion. In

cavitation, collapsing gas bubbles in regions of turbulence and local pressure

fluctuations may activate serious corrosion. Condenser tubes and pump impellers are

subject to this type of attack.

4.9 Stress Corrosion: Stress corrosion or stress corrosion cracking is a term

describing the combined effect of static tensile stresses and the corrosive environment

or a metal. The stresses can be either residual or applied. The residual stresses may

result from unequal cooling, from precipitation and phase transformation, from poor

design, or from cold working and welding. Stress corrosion is characterised by a

highly localised attack occurring when over-all or uniform corrosion is very low.

Season cracking is a term applied to stress corrosion of copper alloys, mainly brasses,

which when highly stressed under go cracking in atmosphere containing traces of

ammonia.

Stainless steels may show cracking on acid chloride solutions, sea water and H2S.

Stress Corrosion Cracking results when even a very small pit forms in a metal under

stress. The concentrated stress may either deepen and extend the pit, or crack any

protective film which might tend to form. Under continued exposure to the corrosive

medium and stress, the crack extends by alternate corrosion and stress failure.

4.10 Corrosion Fatigue Corrosion fatigue may be described as the reduction of

fatigue strength by corrosive environment. In a corrosive environment the corrosion

fatigue strength is much below the fatigue limit. The damage ratio (i.e corrosion

fatigue strength divided by air fatigue strength) due, for example, to a salt water

environment is about 0.4 for Aluminium alloys, 0.5 for Stainless Steels, 0.2 for plain

carbon steels and 1.0 for Cu.

Corrosion Fatigue occurs when a metal is subjected to alternating stress and relief in a

corrosive environment. Metal failure occurs much more rapidly under the alternating

stress than under either stress or relief alone. Continuous removal of protective films,

and the repeated exposures of clean metal by small stress failures cause corrosion

fatigue.

4.11 Hydrogen Embrittlement and hydrogen attack results when atomic hydrogen,

contained in chemical and refinery processes or produced electrolytically by the

process, penetrates onto the grain boundaries of steel producing micro cracks,

blistering and loss of ductility. The atomic hydrogen combines into molecules, cannot

escape, resulting in blistering and laminations.

4.12 Caustic Embrittlement is usually the result of steam escaping through a

crevice, such as between boiler plates or pipe flanges. The escaping water, usually in

22

a fissure of steam, becomes highly concentrated, and the increased alkalinity of the

concentrated water causes failure by stress corrosion.

Caustic embrittlement is a stress corrosion phenomenon occurring on mild steel

exposed to alkaline solutions at high temperatures and stresses. In steam boilers,

water of high alkalinity attacks the steel plate at the crevices near rivets.

4.13 Dry Oxidation and Tarnish result when the clean surfaces of metals are

exposed to air or other gasses to form oxides or other compounds on the surface of the

metals. Many of these films are invisible at room temperature, but at higher

temperatures these films may reach considerable thickness. The rate of film growth is

usually greater at higher temperatures. At higher temperatures, and more particularly

under changes of temperature, the film may crack or spall to expose fresh metal to

attack. Bending or stressing the metal can induce or increase the spalling. The

presence of sulfur-bearing gasses may greatly increase the rate of attack. And the

presence of moisture will accelerate attack, and complicate it by permitting

electrolytic corrosion.

4.14 Dezincification One constituent of an alloy, such as the zinc in brass or bronze,

may be selectively removed, leaving a porous replica of the original part. Often the

whole alloy is initially dissolved, with one element redeposited in spongy form.

4.15 Graphitization is most often seen as the electrolytic corrosion of cast iron, and

often takes a form very similar to the dezincification of brass. Iron is removed

selectively, leaving a replica composed of carbon or graphite.

4.16 Fretting Corrosion. Metal surfaces in close physical contact, in a corrosive

environment, and subjected to vibration, can accelerate corrosion attack by the

continuous removal of protective films. Machine parts with small relative motion and

high unit loads are subject to fretting corrosion.