ELECTROCHEMICAL THERMODYNAMICS
AND ELECTRODE POTENTIAL
Reading Material: Chapter 2 in
“Principles and Prevention of Corrosion, Denny Jones, Prentice-Hall, 1996”.
Dr. Ramazan Kahraman
Chemical Engineering Department
King Fahd University of Petroleum & Minerals
Dhahran, Saudi Arabia
1
ELECTROCHEMISTRY
● Deals with chemical changes produced by an electric
current and with the production of electricity by chemical
reactions
● All electrochemical reactions involve transfer of electrons
and are redox (oxidation-reduction) reactions
● Electrochemical reactions take place in electrochemical
cell (an apparatus that allows a reaction to occur through
an external conductor)
2
ELECTROCHEMICAL CELLS
Two types:
1. Electrolytic cells: - these are cells in which an external
electrical source forces a nonspontaneous reaction to
occur
(one common process is called electrolysis)
(not to be covered here)
2. Voltaic cells: - also called galvanic cells. In these
cells spontaneous chemical reactions generate
electrical energy and supply it to an external circuit
(corrosion and fuel cells are examples)
3
ELECTRODES
(as also covered earlier)
● Electric current enters and exits the cell by electrodes electrodes are surfaces upon which oxidation or reduction
half-reactions occur
● Two kinds of electrodes:
▬
Cathode: - electrode at which reduction occurs
(electrons are gained by a species)
▬
Anode: - electrode at which oxidation occurs (as
electrons are lost by some species)
4
GALVANIC CELLS
● Cells in which spontaneous reactions produces electrical
energy
● The two half-cells are separated so that electron transfer
occurs through an external circuit
● Each half-cell contains the oxidized and reduced forms of
species in contact with each other
● Half-cells linked by a piece of wire and a salt bridge
5
+
6
A salt bridge has three functions
1. It allows electrical contact between the two halfcells
2. It prevents mixing of the electrode solutions
3. It maintains electrical neutrality in each half-cell
as ions flow into and out of the salt bridge
Note that: No electron flow between the two electrodes in separate solutions
without the salt bridge.
7
In galvanic cells, voltage/potential difference between the
electrodes drops to 0 as the reaction proceeds
Salt bridge
Zn Zn 2+ (1.0 M) Cu
2+ (1.0
M) Cu
Electrode
Species (with
concentrations) in
contact with electrodes
8
The Silver-Copper Cell
• Composed of two half-cells:
1. A strip of copper immersed in 1 M CuSO4
2. A strip of silver immersed in 1 M AgNO3
• Experimentally we see:
- Initial voltage (potential difference between electrodes
when not connected yet) is 0.46 volts
- When electrodes connected
- The mass of the copper electrode decreases
- The mass of the silver electrode increases
- [Cu2+] increases and [Ag+] decreases
9
The Silver-Copper Cell (cont.)
Cu → Cu2+ + 2e-
(oxidation, anode)
2(Ag+ + e- → Ag)
(reduction, cathode)
2Ag+ + Cu → Cu2+ + Ag
(overall cell reaction)
Cu |Cu2+(1.0 M) ||Ag+(1.0 M) | Ag
•
Notice that in this case the copper electrode is the anode
10
FREE ENERGY and ELECTRODE POTENTIAL
• Associated with each galvanic cell is a potential difference between
the electrodes called the equilibrium cell potential, Ecell (also called
reversible or electrochemical potential or electromotive force – emf)
E = Ecathode - Eanode
The electrodes are considered not connected to each other so no
reaction takes place on them yet, i.e. each electrode is at equilibrium
with its environment
• E measures the spontaneity of the cell’s redox reaction
∆G = -nFEcell
• Higher (more positive) cell potentials (more negative ∆G) indicate a
greater driving force for the reaction as written (in forward direction,
i.e. from left to right)
11
STANDARD ELECTRODE POTENTIAL
• The standard equilibrium cell potential (E°cell) is for the cell operating
under standard state conditions
• For an electrochemical cell, standard conditions are:
- solutes at 1 M concentrations
- gases at 1 atm partial pressure
- solids and liquids in pure form
- all at some specified temperature, usually 298 K
12
The electrode potentials are commonly measured versus the Standard
Hydrogen Electrode (SHE): E° = 0.00 V
EoZn= -0.762 vs SHE
Zn and Pt electrodes are at equilibrium with their environments at standard
conditions and they are are not connected to eachother by a conductor
so no current flows between Zn and Pt. The potential measured (EoZn) is the
equilbrium potential at standard conditions vs SHE.
13
EQUILIBRIUM POTENTIAL
When we immerse a metal in solution, there will be a tendency for the
metal to react with the solution, either with metal atoms dissolving as
cations or cations already in the solution depositing as metal atoms:
Zn → Zn2+ + 2eZn2+ + 2e- → Zn
As a result of these reactions, the metal will tend to accumulate a
negative or positive charge. The build-up of this charge on the metal will
change its potential in such a way as to inhibit the reaction generating
the charge until the potential reaches a value at which the rates of the
two reactions are equal and opposite. This is known as the equilibrium
potential, and is the potential the metal will adopt in the solution in the
absence of any other reactions.
14
15
STANDARD ELECTRODE POTENTIAL (cont.)
• For covenience all half-cell reactions are written in the table in
reduction form and the electrode potentials are invariant
e.g. Zn=Zn2++2e- and Zn2++2e-=Zn are identical and represent zinc
in equilibrium with its ions with a potential of -0.762 V vs SHE
• The more positive the E° value for a half-reaction the greater the
tendency for the reaction to proceed as written (in reduction-cathodic
form) at standard conditions
• The more negative the E° 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 the course of a reaction, in that case E°cell and
the actual Ecell are related by the Nernst Equation
16
THE NERNST EQUATION
E = E° - (RT/nF) lnK
E = equilibrium potential under the nonstandard conditions
E° = equilibrium potential at standard state (all reactants and products at unit activity)
R = gas constant, 8.314 J/mol.K
T = absolute temperature
n = number of equivalents (moles of electrons) transferred
F = Faraday’s constant, 96,485 J/V ⋅ equivalent ⋅ mol (= C/equivalent ⋅ mol)
K= ∏
activities of products raised to the power of their coefficients
activities of reactants raised to the power of their coefficients
p.s. If the equation is to be used for determining the equilibrium potential for a half
cell rxn, then the half cell rxn is to be written in reduction form or the (-) sign in
the equation is to be replaced by (+) sign in case the half cell rxn is written in
oxidation form.
17
Notation: E for the overall cell rxn, e for the half-cell rxns
p.s. If the equation is to be used for determining the equilibrium potential for
a half cell rxn, then the half cell rxn is to be written in reduction form or
the (-) sign in the equation is to be replaced by (+) sign in case
the half cell rxn is written in oxidation form.
18
ACTIVITY
• Activity of a dissolved species A (A) is equal to its 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.
• Activity of a gas is approximated at ordinary pressures by its partial
pressure in atmospheres (atm).
• The activities of pure solids and water are set equal to unity in
aquaeous solutions.
• At 25°C, 2.303RT/F = 0.0592 V ⋅ equivalent
19
ACTIVITY (Cont.)
Activity Coefficients of Strong Electrolytes (M=molality)
[“Corrosion and Corrosion Control”, H. H. Uhlig and R. W. Revie, John Wiley & Sons, 1985.]
20
ACTIVITY (Cont.)
Activity Coefficients of Strong Electrolytes (M=molality)
[“Corrosion and Corrosion Control”, H. H. Uhlig and R. W. Revie, John Wiley & Sons, 1985.]
21
ACTIVITY (Cont.)
Activity Coefficients of Strong Electrolytes (M=molality)
[“Corrosion and Corrosion Control”, H. H. Uhlig and R. W. Revie, John Wiley & Sons, 1985.]
22
PREDICTION OF SPONTANEITY
1. First write the half-cell rxn with the more positive (less
negative) E° for the reduction-cathodic along with its half
cell electrode potential
2. Write the other half-cell rxn as an oxidation-anodic and
include its half cell electrode potential
3. Balance the electron transfer
4. Obtain the cell rxn by adding the reduction and oxidation
half-cell rxns.
5. Determine the overall cell potential, Ecell, from
Ecell= Ecathode- Eanode
23
Note that half-cell reaction potentials are the same
regardless of the species’ stoichiometric coefficient in
the balanced equation.
Ecell > 0
Forward reaction (left-to-right) is
spontaneous
Ecell < 0
Backward reaction (right-to-left) is
spontaneous
24
EXAMPLE PROBLEM 1
Notation: e for the half-cell rxns, E for the overall cell rxn.
25
EXAMPLE BROBLEM 2
Notation: E for the overall cell rxn, e for the half-cell rxns
26
27
EXAMPLE PROBLEM 3
Notation: e for the half-cell rxns, E for the overall cell rxn.
Assuming standard states for all reactants and products, determine the
spontaneous direction of the following reactions by calculating the cell potential:
CuCl2 + H2 = Cu + 2HCl
CuCl2 + H2 Æ Cu + 2HCl
28
A NOTE FOR THE SIGN NOTATION
Note that the sign notation used in the text is different and
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.
29
SECONDARY REFERENCE ELECTRODES
• If a reference electrode other than SHE is used to
measure the equilibrium potential of a reaction, 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)
30
Potential Values for Common Secondary Reference Electrodes
(Standard Hydrogen Electrode included for reference)
31
EXAMPLE PROBLEM
(Prb. 2.15 in “Principles and Prevention of Corrosion”, Denny Jones, 1996)
A corrosion potential of -0.229 V versus SCE was measured for a corroding alloy. What
is the potential versus (a) SHE, (b) Ag/AgCl (saturated), (c) Cu/saturated CuSO4?
(a)
(b)
(c) Home Exercise
32
Potential
The Pourbaix (Equilibrium Potential-pH)
Diagram
2.0
1.6
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
O2 is stable
2H2O = O2 + 4H+ + 4e-
H2O is stable
2H+ + 2e- = H2
H2 is stable
0
7
14
33
Potential
Pourbaix Diagram for Zinc
2.0
1.6
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
Equilibrium for
Zn2+ + 2OH- ⇔ Zn(OH)2
ZnO22stable in
solution
Zn(OH)2
stable
solid
Zn2+ stable
in solution
Equilibrium for
Zn + 2OH- ⇔ Zn(OH)2 + 2eEquilibrium for
Zn + 4OH- ⇔ ZnO22- + 2H2O + 2e-
Zn metal
stable
0
Equilibrium for
Zn(OH)2 + 2OH- ⇔ ZnO22- + 2H2O
7
14
Equilibrium for
Zn ⇔ Zn2+ + 2e-
34
Pourbaix Diagram for Zinc
Potential
Corrosion requires strong oxidising agent
2.0
1.6
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
Corrosion requires strong
oxidising agent
Corrosion
Passivity
Corrosion
Corrosion is possible,
but likely to be stifled by
solid corrosion product
Corrosion possible with
oxygen reduction
Corrosion possible with
hydrogen evolution
Immunity
0
7
14
Corrosion is thermodynamically impossible
35
2H+ + 2e- → H2
hydrogen evolution in acids
2H2O + 2e- → H2 + 2OH-
hydrogen evolution in water/bases
(the above two reactions are equivalent reactions)
O2 + 2H2O + 4e- → 4OH- oxygen reduction in water/bases
O2 + 4H+ + 4e- → 2H2O
oxygen reduction in acids
(the above two reactions are equivalent reactions)
36
Potential
Pourbaix Diagram for Gold
2.0
1.6 C
Passivity
1.2
0.8
0.4
Gold metal stable
0.0
-0.4
Immunity region
-0.8
-1.2
-1.6
0
7
Gold can’t corrode
with oxygen reduction
or hydrogen evolution
C
14
37
2.0
1.6
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
Cu oxides
stable
Cu2+ stable
in solution
CuO22- stable in soln.
Potential
Pourbaix Diagram for Copper
No - hydrogen evolution
only occurs below the
potential for copper
corrosion
Cu metal stable
0
7
Will copper corrode
in deaerated acid?
14
38
2.0
1.6
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
Cu oxides
stable
Cu2+ stable
in solution
CuO22- stable in soln.
Potential
Pourbaix Diagram for Copper (Cont.)
Usually it will just passivate,
but corrosion can occur in
slightly acid solutions
Cu metal stable
0
7
Will copper corrode
in neutral aerated
waters?
14
39
Potential
Pourbaix Diagram for Iron
2.0
1.6
1.2
Fe3+
0.8
0.4
Fe oxides
stable
0.0
-0.4 Fe2+ stable
-0.8
Fe metal stable
-1.2
-1.6
0
7
Will iron corrode in
deaerated acid?
14
Yes - there is a reasonably
wide range of potentials
where hydrogen can be
evolved and iron dissolved
40
Potential
Pourbaix Diagram for Iron
2.0
1.6
1.2
Fe3+
0.8
0.4
Fe oxides
stable
0.0
-0.4 Fe2+ stable
-0.8
Fe metal stable
-1.2
-1.6
0
7
Will iron corrode in
neutral waters?
Yes - although iron can form
an oxide in neutral solution, it
tends not to form directly on
the metal, as the potential is
too low, therefore it is not
protective.
14
41
Potential
Pourbaix Diagram for Iron
2.0
1.6
1.2
Fe3+
0.8
0.4
Fe oxides
stable
0.0
-0.4 Fe2+ stable
-0.8
Fe metal stable
-1.2
-1.6
0
7
Will iron corrode in
alkaline solutions?
No - iron forms a solid oxide
at all potentials, and will
passivate
14
42
Potential
Pourbaix Diagram for Aluminum
1.2
0.8
0.4
0.0
-0.4
-0.8
-1.2
-1.6
-2.0
-2.4
Al3+
Al2O3
AlO2-
Al
0
7
14
43
Limitations of Pourbaix Diagrams
z
Tell us what can happen, not necessarily
what will happen
z
No information on rate of reaction
z
Can only be plotted for pure metals and
simple solutions, not for alloys
44
EXAMPLE PROBLEM
(Prb. 2.10 in “Principles and Prevention of Corrosion”, Denny Jones, 1996)
Using the Pourbaix diagram for nickel, give the anodic and cathodic reactions on Ni
in water for the following conditions, assuming activity of 10-6 for all soluble species:
(a) deaerated pH 2, (b) deaerated pH 10, (c) aerated pH 2, aerated pH 10.
45
46
Home Work Problems
z
Determine whether silver will corrode with hydrogen evolution (PH2=1
atm) in deaerated KCN solution, pH=9, when CN- activity=1.0 and
Ag(CN)2 activity=0.001. What is the cell potential in volts?
Ag(CN)2- + e- ↔ Ag + 2CN-
z
eo=-0.31 V
Prbs. 3, 5, 6 of Chapter 2
in “Principles and Prevention of Corrosion”, Denny Jones, Prentice-Hall, 1996.
47
Home Exercise Problems
z
Prbs. 7, 8, 11 and 12 of Chapter 2
in “Principles and Prevention of Corrosion”, Denny Jones, Prentice-Hall, 1996.
48
References
z
“Principles and Prevention of Corrosion”, Denny Jones, Prentice-Hall, 1996.
z
“Corrosion Engineering”, Mars Fontana, McGraw-Hill, 1986.
z
“Corrosion and Corrosion Control”, H. H. Uhlig and R. W. Revie, John Wiley &
Sons, 1985.
z
Web Site of Dr. R. A. (Bob) Cottis.
z
Web Site of Dr. Floyd Beckford
49