3 1 Thermodynamics of Corrosion
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CH.CH.22. . TTHERMODYNAMICSHERMODYNAMICS OFOF CCORROSIONORROSION
1Image Source: Corrosion Doctors, www.corrosion-doctors.org
Why Do Metal Corrode?Why Do Metal Corrode?} Any spontaneous reaction in the universe is associated with a lowering in the free
energy of the system, i.e., a negative free energy change.} All metals except some noble metals have free energies greater than those of their
yy
} All metals except some noble metals have free energies greater than those of their compounds. Therefore, they tend to become their compounds through the process of corrosion.
} As an example, take the following examples:
kJ/mol119GCu(OH)O1
OHCu
kJ/mol596G,2Mg(OH)2O21
O2HMg
=++
J/molk66G,3Au(OH)2O43
O2H23
Au
kJ/mol119G,2Cu(OH)2O2O2HCu
+=++
=++
} It is obvious from G that Mg and Cu will corrode.} Au has a positive G and does not corrode.
T d t i M C A
2
} Tendency to corrosion: Mg > Cu >> Au
Electrochemical Nature of CorrosionElectrochemical Nature of Corrosion
} All metallic corrosion are electrochemical reactions, i.e., metal is converted to its ions and compounds with a transfer of electrons.converted to its ions and compounds with a transfer of electrons.
} The overall reaction may be split into oxidation (anodic) and reduction (cathodic) partial reactionsreduction (cathodic) partial reactions.
} Electrochemical reactions take place in electrochemical cell.
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Electrochemical CellElectrochemical Cell
Metallic PathMetallic Path
ee--
A C+ ions
A C+ ions
A C- ions
Electrolytic Path
A C ions
Electrolytic PathElectrolytic Path
Conventional Current Flow
Electrolytic Path
Conventional Current Flow
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Conventional Current FlowConventional Current Flow
Electrochemical CellsElectrochemical Cells
Two types:
1. Electrolytic Cells - these are cells in which an external electrical source forces a non-spontaneous reaction to occur.
(One common process is called electrolysis.)(not to be covered in this course)
2. Galvanic Cells also called voltaic cells. In these cells, ,spontaneous chemical reactions generate electrical energy and supply it to an external circuit.
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(e.g., corrosion, battery and fuel cell, etc.)
2e gainedper Cu2+ ionreduced2e
lostper Zn atom
e
Cu2+ eCuZn2+Zn
Voltmetere e
oxidized
Salt bridgeNa+Zn CuSO42
Anode()
Cathode(+)
Zn2+ Cu2+
Oxidation half-reaction
Reduction half-reaction
Zn(s)
Zn Cu
Zn2+(aq) + 2e
Overall (cell) reactionZn(s) + Cu2+(aq)
Cu2+(aq) + 2e Cu(s)
Zn2+(aq) + Cu(s)Image Source: M.S. Silberberg, Chemistry:
The Molecular Nature of Matter and Change, 2nd ed.
Cell NotationCell Notation
Zn (s) | Zn2+ (aq 1 0M) || Cu2+ (aq 1 0M) | Cu (s)Zn (s) | Zn (aq, 1.0M) || Cu (aq, 1.0M) | Cu (s)
C ll N i} Cell Notation Write components in sequence Separate phases with a single vertical line | A salt bridge or membrane is represented by a double vertical line || Included a specification of the species concentration
} In a galvanic cell, voltage/potential difference between the electrodes drops to 0 as the reaction proceed (to be covered in Chapter 3).
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Corrosion ProcessCorrosion ProcessConsider what happens when a metal electrode M is dipped into a solution containing the metal ions Mn+. The concentration of Mn+ in the metal is much larger than in the
l ti hi h t h i l d i i f th t t d t li th t tisolution, which creates a chemical driving force that tends to equalize the concentration. A charge separation occurs as a result of the metals tendency to oxidize.
M Mn+ (s) + ne- (M) : Dynamic equilibrium.M M (s) ne (M) : Dynamic equilibrium.
e-+
++++
++++
++
CuZn++
+
++++
++++++ + ++
++
++
WATERM t l
8
water+
WATER
Forces of ionization
Metal
Image source: Prof. H.S. Kwons Lecture Note (corrosion.kaist.re.kr)
Solvation Shells (Hydration Sheaths) of a CationSolvation Shells (Hydration Sheaths) of a Cation
Schematic of the primary and secondary solvent molecules for a cation in water.
Primary shell with completely oriented water moves as and where the ion moves. Secondary shell with partly-oriented water is a region of disturbance
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Secondary shell with partly oriented water is a region of disturbance..
Structure of Electrified InterfacesStructure of Electrified InterfacesThe metal electrode - solution interface is also hydrated because of the charge effects.
Schematic of a charged interface and the location of cations at the electrode surface.
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Structure of Electrified InterfacesStructure of Electrified InterfacesBecause of adsorbed H2O on surface and primary hydration shell, ions can only approach to within a fixed distance from the electrode: the OUTER HELMHOLTZ PLANE (OHP). The resulting double layer of charge is often treated theoretically as a capacitor..
--
-+++
--
---
Scattered ions
++ -
-- -
--
'Stuck' ions
M S
M
x
S
St d l f EDL
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Simplified double layer at a metal-solution interface.
Stern model for EDL
Structure of Electrified InterfacesStructure of Electrified InterfacesThe electric double layer (EDL; ) at the surface acts as a barrier to electron transfer.
A potential drop develops at the metal-solution interface(At equilibrium, this potential drop is representative of the reversible (electrode) potential )potential.)
This causes the surface to polarize, i.e., the potential has to change in order for the electrons to overcome the barrier.
This is called ACTIVATION POLARIZATION (CH.3).
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Corrosion ThermodynamicsCorrosion ThermodynamicsThe free energy for any species, G:
G = G0 + RT ln a
yy
Consider the generic chemical reaction:
aA + bB = cC + dD
The free energy change, G, which is a measure of driving force for the reaction to proceed, is given by the difference of the free energy of the products and reactants.
++==
BAdC
tstanreacproducts
)bGaG()dGcG(
GGG
+=
++++++=
bB
aA
dD
cC0
B0BA
0AD
0DC
0C
BAdC
aaaa
lnRTG
)alnbRTbGalnaRTaG()alndRTdGalncRTcG(
)()(
In generalized form:
+=
ktstanreac
jproducts0
)a()a(
lnRTGG
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IF G
Gibbs Function and WorkGibbs Function and Work} Start with the First Law of Thermodynamics and some standard thermodynamic
relations. We find
Gibbs Function and WorkGibbs Function and Work
dU = dq + dwdq = T dSdw = PdV + dwelectricalelectrical
dHP = dUP + PdVdU = T dS PdV + dwelectrical
P PdGT = dHT T dS
= dUT ,P + PdV T dS= T dS PdV + dwelectrical + PdV T dS
dGT ,P = dwelectrical
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Thus the Gibbs function is at the heart of electrochemistry, for it identifies the amount of work we can extract electrically from a system.
Free Energy and Electrode PotentialFree Energy and Electrode Potential} Here we can easily see how this Gibbs function relates to a
potential.
gygy
welectrical = V Qsince Q = n F
} By convention we identify work which is being done by the system
= n F E
} By convention, we identify work which is being done by the system on the surroundings has a negative sign. And negative free energy change is identified as defining a spontaneous process.
GT ,P = welectrical = n FE
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} Note how a measurement of a cell potential directly calculates the Gibbs free energy change for the process.
Free Energy and Electrode PotentialFree Energy and Electrode Potential
} Associated with each electrochemical (galvanic) cell is a potential difference between the electrodes called the equilibrium cell potential,
( l ll d ibl l h i l i l l i
gygy
Ecell (also called reversible or electrochemical potential or electromotive force emf)
E = E E = (e +e )Ecell = Ecathode Eanode = (ea+ec)
} The electrodes are considered not connected to each other so no reaction takes place on them yet, i.e., each electrode is at equilibriumreaction takes place on them yet, i.e., each electrode is at equilibriumwith its environment.
} E measures the spontaneity of the cells redox reaction
G = -nFEcell
} Higher cell potentials (more negative G) indicate a great driving force f th ti itt (i f d di ti i f l ft t i ht)
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for the reaction as written (in forward direction, i.e., from left to right).
HalfHalf--cell potentialcell potential
} Synonyms (Single) electrode potential
pp
(Single) electrode potential Redox (reduction/oxidation) potential
Cathode
Cathode Potential DifferenceElectrolyte
Ecell
A d P t ti l Diff
Cat ode ote t al e e ce
Anode
Anode Potential Difference
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Standard state298K, at unity activity
Selection of hydrogen evolution reaction as having a standard half-cell potential of 0 000Vcell potential of 0.000V.
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HalfHalf--cell Potentialcell Potential
We do not have a means of measuring the potential difference between the electrode and the solution.
?
Cathode
?
Electrolyte?
S l Ecell
Anode
Solution:
Adopt an arbitrary reference electrode.
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?
HalfHalf--cell (electrode) Potentialcell (electrode) Potential( )( )
Measurement of voltage with electrical circuitMeasurement of voltage with electrical circuitConnection for an electrochemical measurement
Source: P.J. Moran, E. Gileadi. J. Chem. Educ. 66 (1989) 912.
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Inclusion of second interface allows combined voltage of the two interfaces to be measured with a voltmeter
Standard Hydrogen ElectrodeStandard Hydrogen Electrodey gy g
} The convention is to select a particular electrode and assign its d d d i i l h l f 0 0000V Thi l d istandard reduction potential the value of 0.0000V. This electrode is
the Standard Hydrogen Electrode (SHE).
2H+(aq) + 2e H2(g)H2
V000.0e0H/H 2=+
H2
PtThe standard aspect to this cell is that the activity of H2(g) and that of H+(aq) are both 1. This means that the pressure of H2 is 1 atm
d h i f H i 1M i h
H+
and the concentration of H+ is 1M, given that these are our standard reference states.
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H+
HalfHalf--Cell Potential MeasurementCell Potential Measurement
SHE0
Zn/ZnV762.0e 2 =+
} Zn and Pt are at equilibrium with their environments at standard conditions.} The measured potential is the equilibrium potential at standard conditions
vs SHE
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vs. SHE.
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Standard Electrode PotentialStandard Electrode Potential
} For convenience, all half-cell reactions are written in the table in reduction form. Zn=Zn2++2e- and Zn2++2e-=Zn are identical and represent zinc in
equilibrium with its ions with a potential of -0.762V vs. SHE.
Th iti th 0 l f h lf ll ti th t} The more positive the e0 value for a half-cell reaction the greater the tendency for the reaction to proceed as written (in reduction-cathodic form) at standard conditions.
} The more negative the e0 value, the more likely is the reverse of the reaction as written (in oxidation-anodic form) at standard conditions.
} However, usually concentrations of reactants differ from one another and also change during a reaction, in that case E0cell (or e0half-cell) and
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the actual Ecell (ehalf-cell) are related by the Nernst Equation.
Nernst EquationNernst EquationConsider the following reaction involving electron transfer
b d
qq
aA + bB + cC + dD +
nFEaaaa
lnRTGG bB
aA
dD
cC0 =
+=
And for the following half-cell reaction in reduction form
KlognF
RT303.2E
aaaa
lognF
RT303.2EKln
nFRT
Eaaaa
lnnFRT
EE 0bB
aA
dD
cC00
bB
aA
dD
cC0 =
==
=
g
aA + bB + + ne- cC + dD +
dcdc aaRT3032aaRT
O1 + O2 + + ne- R1 + R2 +
=
= b
BaA
DC0bB
aA
DC0
aaaa
lognF
RT303.2e
aaaa
lnnFRT
ee
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[ ][ ]
+=
RO
lnnFRT
ee 0
ActivityActivity
} Activity of a dissolved species A (aA or (A) in textbook) is equal to its concentration in moles per 1000 grams of water (molality) multiplied
yy
concentration in moles per 1000 grams of water (molality) multiplied by the activity coefficient, f.
} Activity coefficients are extensively tabulated in numerous chemical and electrochemical handbooks.a d e ect oc e ca a dboo s
} Activity of a gas is approximated at ordinary pressures by its partial pressure in atmospheres (atm)pressure in atmospheres (atm).
} The activities of pure solids and water are set equal to unity in aqueous solutions.
} At 25C 2 303RT/F 0 0592 V equivalent} At 25 C, 2.303RT/F 0.0592 Vequivalent
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Activity CoefficientActivity Coefficientyy
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Source: Corrosion and Corrosion Control, H.H. Uhlig and R.W. Revie, John Wiley and Sons, 1985.
Prediction of SpontaneityPrediction of Spontaneity
1. First, write the half-cell reaction with the more positive (less negative) e0 for the reduction-cathodic along with its half-cell
p yp y
negative) e for the reduction cathodic along with its half cell electrode potential.
2. Write the other half-cell reaction as an oxidation-anodic and include its half-cell electrode potentialts a ce e ect ode pote t a
3. Balance the electron transfer4. Obtain the cell reaction by adding the reduction and oxidation half-
cell reactionscell reactions5. Determine the overall cell potential, Ecell from
EEcell = ec - ea
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Example Example 11(Notation: e for the half(Notation: e for the half cell rxn E for the overall cell rxn )cell rxn E for the overall cell rxn )
For the reaction 2H+ + 2e- = H2 at 25oC and 1 atm.
(Notation: e for the half(Notation: e for the half--cell rxn, E for the overall cell rxn.)cell rxn, E for the overall cell rxn.)
220
220
)H()H(
log2
0592.0e
)H()H(
lognF
RT303.2ee ++ ==
2H
)H(
Plog
20592.0
0e 2+=
[ ]205920 [ ]2H )Hlog(Plog20592.0e 2 +=[ ])Hlog(2Plog0592.0e += [ ])Hlog(2Plog
2e
2H=
[ ]pH2Plog2
0592.0e
2H+= )Hlog(pH +=
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2
pH0592.0e = when PH2 = 1 atm
Example 2Example 2Lets consider the emf (E) of the following Cu-Zn cell at 250C.
pp
Zn Cu
Zn2+ Cu2+
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Example 2Example 2Reduction reaction for Cu: Cu2+ + 2e- Cu e0 = 0.337 V
and
pp
)Culog(0592.0
3370e 2++=and
Reduction reaction for Zn: Zn2+ + 2e- Zn e0 = -0.763 V
and
)Culog(2
337.0e +=
)Znlog(2
0592.0763.0e 2++=
Lets assume that copper corrodes (i.e., Cu is anode, Zn is cathode).
Overall reaction = [reduction rxn at cathode] [oxidation rxn at anode]
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Cu + Zn2+ Cu2+ + Zn
emf (E) of the cell is then;
)Zn(05920 2+
Lets assume that the activities of Cu & Zn to be the same, then the emf, E = -1.100V (vs. SHE).
Negative E (i.e., positive G) means that the reaction is not spontaneous as written, but goes instead in the opposite direction
)Cu()Zn(
log2
0592.0100.1eeeeE 2CuZnanodecathode ++===
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in the opposite direction.In other words, the Zn electrode corrodes (anode), and Cu2+ plated out on the Cu electrode (cathode).
Example Example 3 3 ((PrbPrb. . 22--33a)a)Assuming standard states for all reactants and products, determine the spontaneous direction of the following reactions by calculating the cell potential:
pp (( ))
Cu + 2HCl = CuCl2 + H2
Anodic Cu2+ + 2e- = Cu ea = ea0 = -0.342 V vs. SHECathodic 2H+ + 2e- = H2 ec = ea0 = 0.000 V vs. SHE
E = ec ea = 0.342 V > 0
The reaction is spontaneous as written ().
NOTEA spontaneous reaction does not necessarily mean a fast reaction !!!
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A spontaneous reaction does not necessarily mean a fast reaction !!!
A Note for The Sign NotationA Note for The Sign Notation
Note that the sign notation used in the text is different and confusing
gg
confusing
The text uses
E = ec+ ea
which is actually identical to the notation presented above since the text takes ea as the negative value of the anode half-cell electrode potential.
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Secondary Reference ElectrodesSecondary Reference Electrodes
} If a reference electrode other than SHE is used to measure the equilibrium potential of a rxn, the potential of the reference
yy
equilibrium potential of a rxn, the potential of the reference electrode relative to SHE should be added to the measured potential if one wants to determine the equilibrium potential of the reaction relative to SHE.
} The reference electrodes are also used to measure the corrosion potential of a corrosion cell (Ecorr) (to be covered later).p ( corr) ( )
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Potential Values for Common Secondary Reference Electrodes Potential Values for Common Secondary Reference Electrodes (Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)(Standard Hydrogen Electrode included for reference)
Image source: A J Bard L R Faulkner Electrochemical Methods 2nd ed John Wiley & Sons Inc 2001
Name Half-Cell Reaction Potential, V vs. SHE
Mercury-Mercurous Sulfate HgSO4 + 2e- = Hg + SO42- +0.615
Image source: A.J. Bard, L.R. Faulkner, Electrochemical Methods, 2nd ed., John Wiley & Sons, Inc., 2001.
Copper-Copper Sulfate CuSO4 + 2e- = Cu + SO42- +0.318
Saturated Calomel Hg2Cl2 + 2e- = 2Hg + 2Cl- +0.241
Silver Silver Chloride (saturated) AgCl + e- = Ag + Cl- +0 222
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Silver-Silver Chloride (saturated) AgCl + e = Ag + Cl +0.222
Standard Hydrogen 2H+ + 2e- = H2 +0.000
Reference Electrode Conversion ScaleReference Electrode Conversion Scale
Image source: R.G. Kelly et al., Electrochemical Techniques in Corrosion Science and Engineering, Marcel Dekker Inc., 2002.
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Example ProblemsExample Problems
A corrosion potential of -0.229 V versus SCE was measured for a corroding metal. What is the potential versus (a) SHE, (b) Ag/AgCl (saturated), (c)
pp
Cu/saturated CuSO4?
(a) Ecorr vs. ref 1 = Ecorr vs. ref 2 + eref 2 vs. ref 1Ecorr = Ecorr vs SCE + eSCE vs. SHEeSCE = 0.241 V vs. SHEEcorr = -0.229 + 0.241 = 0.012 V vs. SHE
(b) & (c)Home Exercise
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Homework ProblemsHomework Problems
} Problems 3, 5, 6, 12 of Chapter 2 in textbook.
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