Short Course _Thermodynmics
Transcript of Short Course _Thermodynmics
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To be a partner of choice in corro
research
orrosion Thermodynamics
5/27/2014 1Corrosion Electrochemistry :
Thermodynamics
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Thermodynamics deals with energy and itschanges in reactions.
Reactions are viewed in terms of changes in freeenergy.
According to the first law of thermodynamicsenergy can be neither created nor destroyed.
The second law states that the free energy isreleased from the system to surroundings in allspontaneous changes.
Corrosion reactions are spontaneous and aregoverned by the laws of thermodynamics.
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Tendency to corrode and rate of corrosion
Tendency to corrodeis determined by the free energy difference
between a metal and its corrosion product.
Consider the reaction:
Rate of corrosionis determined by the size of the energy barrier, the free energy of
activation,
G
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For the forward reaction the temperature dependence is given by theArrhenius equation.
A= undefined constant; R= gas constant, T= absolute temperature
The relationship between free energy and equilibrium constant for a
reaction:
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For a spontaneous transition from high energy(initial
state, Gi) to low energy (final state, Gf), the free energy
change
G = Gf- Gi< 0 (negative)
For a spontaneous reaction to occur, Gmust benegative !
.Mg + H2O + O2 = Mg(OH)2 (
G
o
= -596 kJ/mol)Cu + H2O + O2 = Cu(OH)2 (G o= -119 kJ/mol)
Au + 3/2H2O + O2 = Au(OH)3 (G o= +66 kJ/mol)
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Nernst equation
The free energy change is related to electrical potential by
Faraday's law:
G=
nFE and G=
nFE
Where, F=Faraday constant, 96,500 coulombs/mole; n
number of electrons transferred in the corrosion reaction and
E is the measured potential in volts.
Under standard conditions, G= -nFEo
From G = G+ (RT)ln{[C][D]/{[A][B]},
we have
- nFE= -nFEo+ (RT)ln{[C][D]}/{[A][B],the Nernst equation:
E= Eo- (RT/nF)ln{[[C][D]}/{[A][B]}
This is one of the most fundamental equations incorrosion science and engineering.
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Under standard conditions: T=298k, R=8.3143 J (mol k)-1,
ln X = 2.3 logX
Nernst equation can be written as
E= Eo- (0.059/z)log{[Products]/[Reactants]}
E is the non equilibrium potential generated by the corrosionreaction;
[reactants] = concentration of reactants and
[products] = concentration of products.
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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 ElectrodeThe half-cell in which the hydrogen reaction takes place is called the
standard hydrogen electrode, often abbreviated SHE
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Electrode Standard Electrode Potential,EoV
(SHE)
Au3++ 3e-= Au + 1.50
Cl2 + 2e- = 2Cl- + 1.360
O2+ 2H++ 2e = H2O + 1.228Br2+ 2e = 2Br
_ + 1.065
Ag++ e = Ag + 0.799
Hg22++ 2e = 2Hg + 0.789
Fe2++ e_= Fe3+ + 0.771
I2+ 2e = 2I-
+ 0.536Cu++ e = Cu + 0.520
Cu2++ 2e = Cu + 0.337
2H++ 2e = H2 0.000 (by definition)
Pb2++ 2e = Pb 0.126
Sn2++ 2e = Sn 0.136
Ni2++ 2e = Ni 0.250
Fe2++ 2e = Fe 0.440
Cr3++ 3e = Cr 0.740
Zn2++ 2e = Zn 0.763
Al3++ 3e = Al 1.663
Mg2++ 2e = Mg
2.370+ + =
The logical choice
is for all reactants
in their standard
states and
potentials for this
condition are
described asstandard
electrode
potent ials, Eo.
Noble
or
Cathodic
Activeor
Anodic
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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
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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=-nFE < 0, spontaneous reaction.
Eo= + 0.44 V, dG is negative, iron corrodes spontaneously in acid.
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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)
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(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
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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 bethought of in
terms of the simple corrosion cell.
Separation of anode and cathode
permanent separation
differential aeration
random distribution
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Cell Potential
Dry batteries and Daniell cell
Symbolism of a corrosion cell: Zn | Zn2+|| Cu2+| Cu
Zn electrode on the left, immersed in Zn2+ions; Cu electrode on
the right, immersed in Cu2+
ions. Ionic species are separated by || .Zinc is the anode, copper is the cathode.
The two half-cell reactions are:
Zn = Zn2++ 2e-
Cu2++ 2e- = Cu
The overall reaction Zn + Cu2+= Zn2++ Cu
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To find the theoretical cell potential, using Nernst equation for each
of the half-cell reactions:
E(Zn/Zn2+)=Eo(Zn/Zn2+) - (0.059/2)lg[Zn2+], andE(Cu2+/Cu)=Eo(Cu2+/Cu) - (0.059/2)lg{1/[Cu2+]}.
So the cell potential
E(cell)=Eo(cell) - (0.059/2)lg{[Zn2+]/[Cu2+]}
Convention:
do not use oxidation potential
do use reduction potential
Eo(cell) = Eo(cathode) + Eo(anode)