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BASICS OF CORROSION

Dr. Ramazan Kahraman

Chemical Engineering DepartmentKing Fahd University of Petroleum & Minerals

Dhahran, Saudi Arabia

Reading Material: Chapter 1 in“Principles and Prevention of Corrosion, Denny Jones, Prentice-Hall, 1996”.

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What is Corrosion?

Reaction of a metal with its environment

♦ Aqueous corrosion− reaction with water (usually containing

dissolved ions)♦ High temperature oxidation

− reaction with oxygen at high temperature♦ High temperature corrosion

− reaction with other gases

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Examples of Corrosion

Rusting of steel–corrosion product (rust) is solid

but not protectiveReaction of aluminium with water–corrosion product is insoluble in

water, so may be protectiveBurning of magnesium in air–high temperature oxidation

Rust on an iron surface

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Corrosion Science and Engineering

Corrosion Science– Study of the chemical and metallurgical

processes that occur during corrosion.

Corrosion Engineering– Design and application of methods to

prevent corrosion.

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Why is Corrosion Happening?

Because metals want to go back to their stable states.

For Example, Fe is stable when it reacts with oxygen.

So, in the presence of a corrosive environment, Fe tends to separate (decompose) from steel and reacts with oxygen

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Nature of Corrosion

Formation of cell is essential for corrosion

Corrosion cell comprises of the following

–Anode (supplies e- - oxidation reaction)–Cathode (consumes e- - reduction reaction)–Electrolyte–Conductor (electron path)

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Electrodes

Electrodes are pieces of metal on which an electrochemical reaction is occurring

An anode is an electrode on which an anodic or oxidation reaction is occurring

A cathode is an electrode on which a cathodic or reduction reaction is occurring

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Electrochemical CellElectrochemical Cell

A C

HCl

ANODE

CATHODE

ELECTROLYTE

electrons

ELECTRON PATH

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Electrochemical Cell (cont.)Electrochemical Cell (cont.)

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MetalM

e -e - H+

H+ H+

H+

M+n

Cl-Cl-

H2

H+

HCl solution

Anodic Rxn M M+n + n e-

Cathodic Rxn nH+ + n e- n/2 H2

Corrosion of a Metal in AcidCorrosion of a Metal in Acid

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MetalM

e - H2O

M+n

OH-O2

Aerated H2O or Basic Solution

Anodic Rxn M M+n + n e-

Cathodic Rxn (n/2)H2O + (n/4)O2 + ne- n OH-

Corrosion of a Metal in Aerated Water or Corrosion of a Metal in Aerated Water or Aerated Basic SolutionsAerated Basic Solutions

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Acids and Bases

An acid is a substance that produces excess hydrogen ions (H+) when dissolved in water–examples are HCl, H2SO4

A base (alkali) is a substance that produces excess hydroxyl ions (OH-) when dissolved in water–examples are NaOH, KOH

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Note that H+ and OH- are in equilibrium in water:H2O ⇔ H+ + OH-

The product of [H+] times [OH-] is 10-14, so in pure water both [H+] and [OH-] are 10-7. This leads to the concept of pH, which is defined as -log[H+]

Hence pH = 0 is strong acid, 7 is neutral, and 14 is strong alkali

Acids and Bases (cont.)

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Corrosion of Zinc in Acid

●● Zinc known as a base or active metalZinc known as a base or active metal

● Zinc dissolves with hydrogen evolutionZn + 2HCl → ZnCl2 + H2

But we can separate metal dissolution and hydrogen evolution

Zn → Zn2+ + 2e-

2H+ + 2e- → H2

These are known as electrochemical reactions

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Corrosion of Platinum in Acid

Platinum does not react with acids

Platinum is known as a noble metal

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Zinc and Platinum in Acid – Not Connected

Zn Pt

HCl

Zinc and platinum not connected, no reaction

on platinum

Zn + 2HCl → ZnCl2 + H2metal + acid → salt + hydrogen

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Connection of Platinum to Zinc(This is galvanic corrosion which will be studied in detail later)

Zn

A

Pt

C

HCl

Zinc and platinum connected, current flows and hydrogen is evolved

on platinum

Zn → Zn2+ + 2e-

metal → metal ions + electrons(negligible cathodic rxn on Zn relative to that on Pt)

2H+ + 2e- → H2hydrogen ions + electrons → hydrogen gas

electrons

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External Current Applied to Platinum in Acid

Pt Pt

HCl

+-Oxygen evolved on positive electrode

2H2O → O2 + 4H+ + 4e-

Hydrogen evolved on negative electrode

2H+ + 2e- → H2

Overall reaction

2H2O → 2H2 + O2

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External Current Applied to Platinum in Alkali

Pt Pt

NaOH

+-Oxygen evolved on positive electrode

4OH- → O2 + 2H2O + 4e-

Hydrogen evolved on negative electrode

2H2O + 2e- →H2 + 2OH-

Overall reaction

2H2O → 2H2 + O2

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External Current Applied to Platinum

Hydrogen evolution at one electrode2H+ + 2e- → H2 (acids)

or 2H2O + 2e- → H2 + 2OH- (alkalis)

A piece of metal in the solution

Oxygen evolution at the other electrode2H2O → O2 + 4H+ + 4e- (acids)

or 4OH- → O2 + 2H2O + 4e- (alkalis)

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Anodic ReactionsOxidation reactionsProduce electronsExamples

Zn → Zn2+ + 2e- zinc corrosionFe→ Fe2+ + 2e- iron corrosionAl→ Al3+ + 3e- aluminium corrosionFe2+ → Fe3+ + e- ferrous ion oxidationH2 → 2H+ + 2e- hydrogen oxidation in acids

H2 + 2OH- → 2H2O + 2e- hydrogen oxidation in water or bases2H2O → O2 + 4H+ + 4e- oxygen evolution in acids4OH- → O2 + 2H2O + 4e- oxygen evolution in water or bases

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Cathodic Reactions

Reduction reactionsConsume electronsExamplesO2 + 2H2O + 4e-→ 4OH- oxygen reduction in water/basesO2 + 4H+ + 4e- → 2H2O oxygen reduction in acids2H2O + 2e-→ H2 + 2OH- hydrogen evolution in water/bases2H+ + 2e- → H2 hydrogen evolution in acidsCu2+ + 2e- → Cu copper platingFe3+ + e- → Fe2+ ferric ion reductionSn4+ + 2e- → Sn2+

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Cathodic Rxns in Acidic & Basic Solns

Deaerated Acidic Solutions2H+ + 2e- → H2

Aerated Acidic Solutions2H+ + 2e- → H2

O2 + 4H+ + 4e- → 2H2O(presence of O2 further increases corrosion)

Deaerated Neutral or Basic Solutions2H2O + 2e- → H2 + 2OH-

Aerated Neutral or Basic SolutionsO2 + 2H2O + 4e- → 4OH-

(this reaction causes higher corr. rate)

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Corrosion Rate

Simplest and most useful technique for corrosion rate determination is the Weight Loss Technique

Corrosion Rate = mass / exposed surface area . timeor

Corrosion Rate = avg. corrosion penetration depth / time( = mass / density . surface area . time )

Common Corrosion Rate Units– gmd (grams of metal lost per square meter per day)– mm/y (average millimeters penetration per year)– mpy (avg. mils penetration per year, 1 mil = 0.001 in)

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Example

A carbon steel test specimen of dimensions 2-in × 3-in × 0.125-in with a 0.25-in hole for suspending in solution is exposed for 120 hours in an acid solution and loses 150 milligrams. Calculate the corosion rate in mpy and mm/y.

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Home Exercise Problems

Prbs. 1, 4, 8, 10 and 11 of Chapter 1

in “Principles and Prevention of Corrosion”, Denny Jones, Prentice-Hall, 1996.

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Faraday’s Law

Charge is related to mass of material reacted in an electrochemical reaction:

M → Mn+ + ne-

One metal ion

Reacts

To produce one mol of metal ion and

n mols of electrons

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Faraday’s Constant

One mole of metal (MW g) contains Avogadro’s number (6×1023) of metal atoms

Hence each mole of metal will produce n times that many number of electrons

Charge on the electron is 1.6 × 10-19 C (coulomb)

Hence one mole of metal will produce a charge of n ×96500 C

96500 C/equivalent is known as Faraday’s constant(also in units of J/V⋅equivalent)

Conversions: 1 A (ampere) = 1 C/s, 1 J = 1 C⋅V

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Faraday’s Law

(g/mole) metal of weight (atomic)molecular (g) oxidized metal of mass

metal of molper ed transferrelectrons) of (mols sequivalent ofnumber

nt)C/equivale (96500constant s Faraday'C) (coulomb, charge where

==

===

=

Mm

nFQ

MnFQm

So, if Q is known, mass loss by corrosion can be determined.

The details of corrosion rate determination by electrochemical techniques will be covered later.

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References

“Principles and Prevention of Corrosion”, Denny Jones, Prentice-Hall, 1996.

Web Site of Dr. R. A. (Bob) Cottis.