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Lecture 3: The Theory for Aqueous Corrosion (I)


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4.1 Thermodynamic Aspects of Corrosion Reactions

For most metals, the chemically combined states are preferred by nature:

      • oxides
      • sulphides
      • carbonates
      • other complex compounds

Ores are in low energy states (stable states)

Nature is always concerned with minimizing energy.

Metals extracted from their ores and are in high energy states, thermodynamically, they are not stable. Corrosion is a natural occurring process, predicted by thermodynamics laws.

high energy states (metal) = low energy states (metals compounds)

It is this tendency of metals to recombine with elements presents in the environments that leads to the phenomenon known as Corrosion.

Corrosion is the degradation of a metal by an electrochemical reaction with its environment.

Points to note:

The low energy corrosion product is not the same as an ore (may be similar), the energies of the ore and the corrosion product may well be comparable.

Tendency to corrode and rate of corrosion

Tendency to corrode is determined by the free energy difference between a metal and its corrosion product.

Rate of corrosion is determined by the size of the energy barrier (the free energy of activation, G*)

The rate constant

kcorr =A exp (-G*/RT)

A= undefined constant; R= gas constant, T= absolute temperature

Rate of corrosion reaction = kcorr [reactants]

Points to note about G*:

G* is the minimum amount of energy required to drive the molecules/atoms over the activation energy barrier, so that appreciable reaction can take place.

Free energy is the single factor to determine the possibility of a corrosion reaction.

All interactions between elements and compounds are governed by the free energy changes available to them.

For a spontaneous transition from high energy (initial state, Gi) to low energy (final state, Gf), the free energy change

= Gf - Gi < 0 (negative)

For a spontaneous reaction to occur, must be negative !

Points to note:

At room temperature, most chemical compounds of metals have lower energy than the uncombined metals.

Most metals have an inherent tendency to corrode.

Why do gold, platinum and other precious metals not corrode ?

  • the energetics may not be favorable
  • the size of activation energy may be too great (the rate is extremely slow)

Thermodynamic data alone tells us that copper and magnesium are expected to corrode naturally In wet, aerated atmosphere.

Exceptions:

Iron objects preserved after centuries of immersion at the bottom of peat bogs. There may always be special circumstances why corrosion does not occur when expected.

For a reaction: A + B = C = D, the free energy change, , is given by thermodynamic equation

, where J is the reaction quotient.

At equilibrium condition,

J=K, where K is the equilibrium constant.

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Nernst equation

The free energy change is related to electrical potential by Faraday's law:

Where, F=Faraday constant, 96,500 coulombs/mole; z number of electrons transferred in the corrosion reaction and E is the measured potential in volts.

Under standard conditions,  DGo = -zFEo

From DG = DGo + RTln{[C][D]/{[A][B]}, we have

- zFE= -zFEo + RTln{[C][D]}/{[A][B], the Nernst equation:

E= Eo - (RT/zF)ln{[[C][D]}/{[A][B]}

This is one of the most fundamental equations in corrosion science and engineering.

Under standard conditions: T=298k, R=8.3143 J (mol k)-1, Nernst equation can be written as

E= Eo - (0.059/z)lg{[Products]/[Reactants]}

E is the non equilibrium potential generated by the corrosion reaction; [reactants] = concentration of reactants and [products] = concentration of products.

At equilibrium conditions, E=0

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4.2 The Basic Wet Corrosion Cell

Four essential components of a corrosion cell

  • The Anode
  • The Cathode
  • The Ionic Conductor (electrolyte)
  • The Metallic Conductor (electrical connection)

(1) The anode

The anode corrodes by loss of electrons: M= Mz+ + ze-

  • Anodic reaction
  • Oxidation reaction
  • Electron generation

(2) The cathode

The cathode does not corrode. Most important cathodic reactions:

  • (i) pH < 7 2H+ + 2e- =H2
  • (ii) pH > = 7 2H2O + O2 + 4e- = 4OH-

Other cathodic reactions are possible Depending on the environment.

Cathode

  • Cathodic reaction
  • Reduction reaction
  • Electron consumption

(3) An electrolyte (ionic conductor)

a solution conducting electricity

(4) Electrical connection

the anode and cathode in a corrosion cell must be in electrical contact. Difference in free energies between the anode and the cathode produces electrical potential which is the driving force for corrosion reaction.

Current: flow of electrons; Corrosion current: corrosion rate

Points to note: all aqueous corrosion reactions can be thought of in terms of the simple corrosion cell.

Separation of anode and cathode

  • permanent separation
  • differential aeration
  • random distribution

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4.3 Standard Electrode Potentials

The potential difference between the anode and cathode can be measured by voltage measuring device. The absolute potential of the anode and cathode cannot be measured directly.

To define a standard electrode, all other potential measurements are made against the standard electrode. If the standard electrode potential is arbitrarily set to zero, the potential difference measured can be considered as the absolute potential.

Standard Hydrogen Electrode

The half-cell in which the hydrogen reaction takes place is called the standard hydrogen electrode, often abbreviated SHE

Standard Electrode Potential

The potential difference measured between metal, M, and the hydrogen electrode, under well defined conditions.

Example:

Iron corrodes in a solution of H+

(a) iron dissolves: half-cell reaction: Fe = Fe2+ + 2e-

(b) hydrogen gas formed: half-cell reaction: 2H+ + 2e- = H2

(c) overall reaction: corrosion reaction: Fe + 2H+ = Fe2+ + H2

Substituting into Nernst equation:

E=Eo - (0.059/2)lg{[Fe2+][H2]/[Fe][H+]2}

The terms [H+] and [H2] have been made to 1, [Fe] can be approximated as unity, so under standard conditions, [Fe2+] = 1 M,

E= Eo

The measured potential difference is the electrode potential of the iron under standard conditions.

Points to note: if the measured potential is positive

dG=-zFE < 0, spontaneous reaction.

Eo = + 0.44 V, dG is negative, iron corrodes spontaneously in acid.

The standard electrode potential of other metals can be measured in the same way as for iron:

Points to note

  • Standard reduction potential
  • Standard redox potential

are used in stead of standard oxidation potential

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Summary

In this lecture, we introduced one of the most important equations in corrosion science and engineering, the Nernst Equation. We also analysed the four essential components of a corrosion cell, i.e., the ANODE, CATHODE, ELECTROLYTE and METALLIC (ELECTRON) CONDUCTOR. The Standard electrode potential of a metal in equilibrium condition is measured with reference to the Standard Hydrogen Electrode.

Reading Assignments

To reinforce learnings in this lecture read pages 69- 85 (textbook)
To prepare yourself for the
next lecture read pages 85-95 (textbook)

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