Corrosion Mechanisms

Randhir Kumar Singh

Asst Professor

OPJIT

Chemical vs. Electrochemical Reactions

Chemical reactions are those in which elements are

added or removed from a chemical species.

Electrochemical reactions are chemical reactions in

which not only elements may be added or removed

from a chemical species but at least one of the

species undergoes a change in the number of

valence electron.

Corrosion processes are electrochemical in nature.

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Simplest Example: Dry Cell Battery

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Faraday’s laws of electrolysis, in chemistry,

quantitative laws used to express magnitudes of

electrolytic effects, first described by the English scientist

Michael Faraday in 1833.

The laws state that

(1) the amount of chemical change produced by current at

an electrode-electrolyte boundary is proportional to the

quantity of electricity used, and

(2) the amounts of chemical changes produced by the same

quantity of electricity in different substances are

proportional to their equivalent weights.

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Contd…

In electrolytic reactions, the equivalent weight of a

substance is the gram formula weight associated with a

unit gain or loss of electron. The quantity of electricity

that will cause a chemical change of one equivalent

weight unit has been designated a faraday. It is

equivalent to 9.6485309 × 104 coulombs of electricity.

Thus, in the electrolysis of fused magnesium chloride,

MgCl2, one faraday of electricity will deposit 24.312/2

grams of magnesium at the negative electrode and

liberate 35.453 grams of chlorine at the positive

electrode.

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Faraday’s Law

The mass of an element discharged at an

electrode is directly proportional to the

amount of electrical charge passed through

the electrode

weight of metal reacting = kIt

where I = Current Intensity

t = time of current passage

k = Constant

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What Happens if the Battery is Not in Use?

There will be some “local action current”

generated by “local action cells” because of

other metallic impurities in zinc

Shelf life of an ordinary zinc-carbon rod battery

is limited

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Local Action Cell

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Anode & Cathode

Anode

Loss of electron in oxidation

Oxidation always occurs at the anode

Cathode

Gain of electron in reduction

Reduction always occurs at the cathode

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Difference between Electrochemical and Electrolytic Cell

fes/fe2+

aq //sn2+aq/sns Both types of cells consist of two electrodes

connected to an electrolyte (an ionically conducting phase).

Electrode reactions then take place at the electrode-solution

surfaces. The change from electronic current to ionic current and

visa versa are always accompanied by oxidation/reduction

reactions.

An electrochemical cell is simply a device that converts chemical

energy into electrical energy when a chemical reaction is occurring

in a cell. An electrolytic cell converts electrical energy into chemical

energy.

In an electrochemical cell the reaction occurs spontaneously at the

electrodes, while an electrolytic cell reaction is not spontaneous at

the electrodes - the reaction has to be forced by applying an

external electrical current.

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In an electrochemical cell the cathode is positive and the anode is

negative. In an electrolytic cell the cathode is negative and the anode is

positive (does this mean that the electrons are going against their

gradient here?).

In a spontaneous chemical reaction electrons are passed directly from

one element to another. In an electrochemical cell these simultaneous

redox reactions are "spatially separated" - i.e. happen at different places.

The resultant ions then combine to form a new product. During this

process electrons are conducted from the anode to the cathode through

an outside electrical current which can be used. This action can be

reversed in a electrolytic cell.

Electrochemical cells are used usually as batteries, while electrolytic

cells are used for electroplating metals. Also, the recharging of a

rechargeable battery is an electrolytic reaction.

Corrosion Cells

Galvanic cell (Dissimilar electrode cell) – dissimilar metals

Salt concentration cell – difference in composition of aqueous environment

Differential aeration cell – difference in oxygen concentration

Differential temperature cell – difference in temperature distribution over the body of the metallic material

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Dissimilar Electrode Cell

When a cell is produced

due to two dissimilar metals

it is called dissimilar

electrode cell

Dry cell

Local action cell

A brass fitting connected to a

steel pipe

A bronze propeller in contact

with the steel hull of a ship

Zn anode

HCl Solution

Cu cathode

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Salt Concentration Cell

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Differential Aeration Cell

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Corrosion at the bottom of the electrical poles

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Differential Temperature Cell

This is the type of cell form when two identical electrodes are immersed in same electrolyte, but the electrodes are immersed into solution of two different temperatures

This type of cell formation takes place in the heat exchanger equipment where temperature difference exists at the same metal component exposed to same environment

For example for CuSO4 electrolyte & Cu electrode the electrode in contact with hot solution acts as cathode.

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Factors affecting choice of an engineering material

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Factors affecting corrosion resistance of a metal

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Corrosion Rate Expressions

mm/y – millimeters penetration per year

gmd – grams per square meter per day

ipy – inches penetration per year

mpy – mils penetration per year (1000 mil =

1 inch)

mcd – milligrams per square centimeter per

day

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Corrosion Rate Expressions

The most used expression for Corrosion Rate in the US is the mpy (Mils per

year).

Mils per year (mpy) = 534W/DAT

To convert corrosion rate (corrosion rate conversion) between the mpy and

the equivalent in metric unit mm/y (millimeter per year):

1 mpy = 0.0254 mm/y = 25.4 micron/y

To calculate the corrosion rate from metal loss:

mm /y = 87.6 x (W / DAT)

W = weight loss in mg

D = density of specimen material in g/cm3

A = area in cm2

T= exposure time in hours

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Classification of metallic materials

according to their rate of uniform attack

A. <0.005 ipy (<0.15 mm/y) – Metals in this

category have good corrosion resistance

and can be used for critical parts

B. 0.005 to 0.05 ipy (0.15 mm/y to 1.5 mm/y) –

Metals in this group are satisfactory if a

higher rate of corrosion can be tolerated

C. >0.05 ipy (>1.5 mm/y) – Usually not

satisfactory

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Electrochemical Aspects

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Electrochemical Reactions

Electrochemical Reactions

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The above concept is illustrated in the Fig.1

Fig.1 Electrochemical reactions

occurring during corrosion of zinc

in air-free hydrochloric acid

Electrochemical Reactions

25 Fig.2 Electrochemical reactions occurring during corrosion

of zinc in aerated hydrochloric acid

Let us look at a Zn-H Cell

(a)

(b)

The Zn electrode moves away from equilibrium by the removal of negative charges from the Zn plate and positive ions are released from the Zn plate to the liquid (a)

Zn is dissolved at the same rate as electrons are transported to the Pt plate, where they are consumed in the hydrogen reaction

The same cell process can be totally obtained on a Zn plate submerged in a solution containing hydrogen ions and Zn ions (b)

The reactions are accompanied by the same changes in free enthalpy and have the same equilibrium potentials as before

However, there is a higher resistance against the hydrogen reaction on the Zn plate than on Pt, and thus the reaction rate will be lower on the Zn surface

So We Also Need to Know …

Electrode kinetics to predict the corrosion

rates for the actual conditions

Single and mixed electrodes

Whenever only one electrode reaction takes place on a

metal surface in a given solution, that system is called a

single electrode.

This is the case for copper immersed into de-aerated and

slightly acidic copper sulfate solution:

Cu2+ + 2 e = Cu -- (1)

The open circuit potential (ocp) is the potential set up

spontaneously by an electrode in the absence of an

external current. For a single electrode, the open circuit

potential is equal to the equilibrium potential, Erev.

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Single and mixed electrodes Most often, several electrode reactions take place simultaneously at a metal

electrolyte interface. Such systems are referred to as mixed electrodes. If, in the

previous example, the copper sulfate solution is aerated, two electrode reactions

(called partial reactions) are observed at the open circuit potential; the oxidation of

copper

Cu→Cu2+ + 2 e -- (2)

and the reduction of oxygen:

½ O2 + 2H+ + 2 e → H2O -- (3)

The corresponding overall reaction is

Cu + ½ O2 + 2H+ → Cu 2+ H2O -- (4)

The copper thus corrodes without any external current. The open circuit potential of

a mixed electrode undergoing corrosion, is called the corrosion potential (in the

literature it is sometimes also called the free corrosion potential).

The corrosion potential has a value that lies in between the equilibrium potentials

of the partial electrode reactions. In contrast to the equilibrium potential, which is a

thermodynamic quantity, the corrosion potential is determined by kinetics; its value

depends on the rates of both the anodic and the cathodic partial reactions present.

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Single and mixed electrodes

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Figure A copper electrode in contact with itw own ions (single electrode)

and with an aerated solution (mixed electrode).

Polarization and Overpotential Polarization

Electrode reactions are assumed to induce deviations from equilibrium due to the passage of an electrical current through an electrochemical cell causing a change in the electrode potential. This electrochemical phenomenon is referred to as polarization.

The polarization ζ expresses the difference between the potential of a mixed

electrode subjected to anodic or cathodic polarization and its corrosion potential.

ζ = E – Ecor

A polarization of ζ > 0 indicates an anodic and a polarization of ζ < 0 a cathodic

current flow

Overpotential

The deviation from equilibrium causes an electrical potential difference between the polarized and the equilibrium (unpolarized) electrode potential known as overpotential

η = E – Erev

A positive overpotential indicates that an anodic current is crossing the interface; a

negative one means that the current is cathodic.

Polarization and Overpotential

Equilibrium potential for cathodic reaction = Eoc

Equilibrium potential for anodic reaction = Eoa

Real potential = E

Cathodic Overpotential ηc = E – Eoc < 0

anodic Overpotential ηa = E – Eoa > 0

The Polarized Cell

Exchange Current Density

At the equilibrium potential of a reaction, a reduction and an oxidation reaction occur, both at the same rate.

For example, on the Zn electrode, Zn ions are released from the metal and discharged on the metal at the same rate

The reaction rate in each direction can also be expressed by the transport rate of electric charges, i.e. by current or current density, called, respectively, exchange current, Io, and (more frequently used) exchange current density, io.

The net reaction rate and net current density are zero

How Polarization is Measured

Causes of Polarization

Depending on the type of resistance that

limits the reaction rate, we are talking about

three different kinds of polarization activation polarization

concentration polarization and

resistance (ohmic) polarization or IR Drop

Activation Polarization

When current flows through the anode and the cathode electrodes, their shift in potential is partly because of activation polarization

An electrochemical process that is controlled by reaction sequence at

the metal-electrolyte interface.

This is easily illustrated by considering hydrogen evolution reaction on

zinc during corrosion in acid solution.

An electrochemical reaction may consist of several steps

The slowest step determines the rate of the reaction which requires activation energy to proceed

Subsequent shift in potential or polarization is termed activation polarization

Activation polarization usually is the controlling factor during corrosion

in media containing a high concentration of active species(e.g.

concentrated acids).

Most important example is that of hydrogen ion reduction at a cathode, H+ + e- → ½ H2, the polarization is termed as hydrogen overpotential

Activation Polarization

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Fig. Hydrogen-reduction reaction under activation control(simplified).

Hydrogen Overpotential

Hydrogen evaluation at

a platinum electrode:

H+ + e- → Hads

2Hads → H2

Step 2 is rate limiting

step and its rate

determines the value of

hydrogen overpotential

on platinum

Tafel Equation

Activation polarization (η) increases with

current density in accord with Tafel equation:

The Tafel constant is given by:

oi

ilog

nF

RTβ

α

3.2

Overpotential

Values

Concentration Polarization

It refers to electrochemical reactions that are controlled

by the diffusion in electrolyte.

Sometimes the mass transport within the solution may

be rate determining – in such cases we have

concentration polarization

Concentration polarization implies either there is a

shortage of reactants at the electrode or that an

accumulation of reaction product occurs

Concentration polarization generally predominates when

the concentration of the reducible species is small(e.g.

dilute acids, aerated salt solutions).

Concentration Polarization

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Fig.3 Concentration polarization during hydrogen reduction.

Concentration Polarization: reduction of

oxygen O2H4e4HO 22

IR Drop

When polarization is measured with a potentiometer and a reference electrode-Luggin probe combination, the measured potential includes the potential drop due to the electrolyte resistance and possible film formation on the electrode surface

The drop in potential between the electrode and the tip of Luggin probe equals iR.

If l is the length of the electrode path of cross sectional area s, k is the specific conductivity, and i is the current density then resistance

iR drop in volts = k

lR

k

il

k

il

Combined Polarization

Total polarization of an electrode is the sum

of the individual contributions,

If neglect IR drop or resistance polarization is

neglected then:

rcaT ηηηη

caT ηηη

Combined Polarization

Effect of temperature, concentration and velocity of the aqueous

environment on combined polarization is shown in the figure

Passivity

Passivity refers to the loss of chemical

reactivity experienced by certain metals and

alloys under particular environmental

conditions.

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Fig. Corrosion rate of a metal as a

function of solution oxidizing

power (electrode potential).

Passivity

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Fig. Corrosion characteristics of an active-passive metal as a function

of solution oxidizing power (electrode potential).

Passivity in Iron Chromium Alloys

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Dissolution of an active metal (active dissolution), involves a charge transfer at

the metal-electrolyte interface. Soluble ions, either hydrated or complexed, are

formed and dissolve into the electrolyte, while the liberated electrons either flow

to the cathode or are taken up by an oxidizing agent.

When a passive metal dissolves (passive dissolution), cations are formed by a

charge transfer reaction at the metal-film interface. They migrate across the

passive film to the film-electrolyte interface, where they dissolve into solution as

hydrated or complexed ions.

Cont… Because of the presence of an oxide film, the dissolution rate of a passive

metal at a given potential is much lower than that of an active metal. It

depends mostly on the properties of the passive film and its solubility in the

electrolyte.

During passivation, which is a term used to describe the transition from the

active to the passive state, the rate of dissolution therefore decreases

abruptly.

The polarization curve of a stainless steel in sulfuric acid, given in Figure,

illustrates this phenomenon. In this electrolyte, the corrosion potential of the

alloy is close to –0.3 V.

Anodic polarization leads to active dissolution up to about –0.15 V, where the

current density reaches a maximum. Beyond this point, the current density,

and hence the dissolution rate, drops sharply.

It then shows little further variation with potential up to about 1.1 V. Above

that value the current density increases again because transpassive

dissolution and oxidation of water to oxygen becomes possible.

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Fig. Polarization curve of Fe-17Cr stainless steel in 0.5 M H2SO4. Sweep

rate is 0.02 V/min.

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Figure. Variation of partial anodic current density with potential for a

passivating metal (Evans Diagram).

Generally speaking, we can distinguish three potential regions in the polarization curve of a

passivating metal (Figure ):

• the active region;

• the passive region;

• the transpassive region.

In contrast to the active and passive regions, the surface state of the metal in the transpassive

region is not well defined and an oxide may or may not cover the surface.

The current density measured during a polarization experiment is the

sum of all anodic and cathodic partial current densities. Figure .

Schematically shows the variation of the anodic partial current density of

a passivating metal as a function of the potential. It allows us to define a

number of quantities that describe the polarization behavior of

passivating metals.

The passivation potential Ep separates the active from the passive

potential region.

The corresponding current density at the maximum is the passivation

current density, ip.

The passive current density ipp characterizes the dissolution behavior of

the metal in the passive potential region.

The transpassivation potential Eb marks the end of the passive

potential region and the transition from passive to transpassive behavior.

Beyond this point the anodic partial current density increases markedly

with increasing potential due one of the following processes: uniform

transpassive dissolution resulting from oxidation of the passive film,

dissolution by pitting resulting from local film breakdown, oxygen

evolution due to water oxidation.

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When dissolution by pitting is the dominating reaction, the potential Eb is

called pitting potential or critical pitting potential. Often, Eb is also

referred to as film breakdown potential, indicating that pitting is initiated

by passive film breakdown.

Depending on conditions, the value of Eb can be either above or below

the reversible potential of the oxygen electrode, Erev,O2. For sufficiently

stable passive films with good electronic onductivity, oxygen evolution

rather than transpassive dissolution may therefore account for the

observed current at high anodic potentials.

The pitting potential, plays an important role for the corrosion resistance

of passive metals and alloys. Generally speaking, to have a good

corrosion resistance an alloy should exhibit a low value of Ep and a high

value of Eb.

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Figure Variation of passivation potential for Fe-Cr alloys with pH in

sulfate solutions

Generally, the oxides of less noble metals exhibit a lower standard

potential of formation. These metals passivate

spontaneously in the presence of protons. Furthermore, many oxides

exhibit good chemical stability in acidic environments. This explains

the higher corrosion resistance of metals such as titanium, tantalum

and chromium.

According to the relation,

the value of Erev,oxide decreases by 59 mV per pH unit, regardless of

the stoichiometry of the oxide formed; this is because the number of

charges does not appear in the equation.

The condition Ep ≥ Erev,oxide, suggests that the passivation potential

also decreases with increasing pH.

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Figure confirms this behavior: When the passivation potentials of iron, chromium

and their alloys in sulfuric acid are plotted as a function of the pH, a straight line

is obtained with a slope of –59 mV/pH. Similar results have been found for other

metals.

In this figure, for zero pH (pH = 0; Erev,oxide = E°), the passivation potentials of

chromium and iron do not match the standard potentials of the oxides Cr2O3 and

Fe2O3 listed in Table 6.8.

One explanation is that kinetic limitations lead to a higher passivation potential

than predicted by thermodynamics. Another reason could be that the listed

standard potentials were measured on bulk samples. The extreme thinness of

passive films could influence their thermodynamic properties.

In addition, their composition does not always correspond to a simple

stoichiometry.

For example, chromium-iron alloys form passive films containing both iron and

chromium cations and their passivation lie between those of iron and chromium.

Then again, they exhibit the same pH dependence as the pure metals.

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Figure 6 Anodic polarization curves of Fe, Ni and Cr in 0.5 M H2SO4

In acidic media, the metals iron, nickel and chromium

have passivation current densities that increase in the

order Cr < Ni < Fe.

In Figure 6, the anodic polarization curves for the three

metals in 0.5 M sulfuric acid (25°C)are compared.

Chromium has lower values of both ip and Ep than the

other two metals.

By alloying increasing amounts of chromium to steel one

therefore improves the corrosion resistance.

Experience shows that above a chromium concentration

of 12 to 13%, a steel passivates spontaneously in

contact with aerated water. It becomes "stainless“,

meaning it does not rust easily

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Figure Corrosion potentials of Fe-Cr alloys in aerated 4% NaCl as a function of

their chromium content. The measured average rate of corrosion in salt spray

tests is also shown

Corrosion potentials of iron-chromium alloys in an

aerated solution of 4% NaCl together with the corrosion

rate measured in a salt spray test.

Spontaneous passivation above a chromium content of 8

to 12%, leads to a rise in corrosion potential and to a

drop of corrosion rate.

The magnitude of the passivation current density

depends on different factors:

the kinetics of active dissolution;

the mass transport of the dissolution products;

the pH of the electrolyte;

the water content of the electrolyte.

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THANK YOU

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