Yeungnam University
Basic Concepts in Electrochemistry
Mosab Kaseem
Plasticity Control and Mechanical Modeling Laboratory
School of Materials Science and Engineering, Yeungnam University, Korea
Contents
 redox reactions
 Balancing redox reactions
 electrochemical cells
• electrode processes
• construction
• notation
 cell potential and Go
 standard reduction potentials (Eo)
 non-equilibrium conditions (Q)
Yeungnam University
School of Materials Science and Engineering
Introduction
1)
Redox Reaction involves transfer of electrons from one species to another.
Redox reactions:
electron transfer processes
Oxidation: loss e- ………causes reduction “reducing agent”
Reduction: gain e- ……… causes oxidation “ oxidizing agent”
Zn metal
With time, Cu plates out onto Zn metal
strip, and Zn strip “disappears.”
Cu+2 ions
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Redox Reaction
• Zn is oxidized
reducing agent
Zn(s)  Zn2+(aq) + 2e• Cu2+ is reduced
oxidizing agent
Cu2+(aq) + 2e-  Cu(s)
wire
elect rons
 Electrons travel thrugh external wire.
 Salt bridge allows anions and cations
to move between electrode compartments.
 This maintains electrical neutrality.
Zn
Zn2+ ions
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salt
bridge
Cu
Cu2+ ions
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The first two reactions are known as “1/2 cell reactions”
Include electrons in their equation
3.) The net reaction is known as the total cell reaction
No free electrons in its equation
2.)
½ cell reactions:
Net Reaction:
4.) In order for a redox reaction to occur, both reduction of one
compound and oxidation of another must take place simultaneously
Total number of electrons is constant
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Balancing redox reactions in acidic solution
Cr2O72-(aq) + I-(aq) → Cr3+(aq) + I2(s)
Step 1: Divide the reaction into half-reactions.
Cr2O72- → Cr3+
I→ I2
Step 2: Balance the atoms and charges in each half-reaction.
For the Cr2O72-/Cr3+ half-reaction:
Balance atoms other than O and H:
Cr2O72- → 2Cr3+
Balance O atoms by adding H2O molecules:
Cr2O72- → 2Cr3+ + 7H2O
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Balancing redox reactions in acidic solution
Balance H atoms by adding H+ ions:
14H+ + Cr2O72- → 2Cr3+ + 7H2O
Balance charges by adding electrons:
6e- + 14H+ + Cr2O72- → 2Cr3+ + 7H2O
For the I-/I2 half-reaction:
Balance atoms other than O and H:
2I- → I2
There are no O or H atoms, so we balance charges by adding electrons:
2I- → I2 + 2eStep 3: try to make the number of e- lost in the oxidation equals the number of egained in the reduction.
6e- + 14H+ + Cr2O72- → 2Cr3+ + 7H2O
6I- → 3I2 + 6e6I-(aq) + 14H+(aq) + Cr2O72-(aq) → 3I2(s) + 7H2O(l) + 2Cr3+(aq)
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School of Materials Science and Engineering
Balancing redox reactions in acidic solution
Example:
MnO4-(aq) + C2O4-2 (aq) → Mn+2 (aq) + CO2 (g)
ClO3-(aq) + SO2 (g) → SO4-2 (aq) + Cl- (aq)
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School of Materials Science and Engineering
Balancing redox reactions in basic solution
MnO4-(aq) + Br- (aq) → MnO2 (s) + BrO3- (aq)
Se + Cr(OH)3 → Cr + SeO3-2
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School of Materials Science and Engineering
Electrical parameters
1.)
Electric Charge (q)
Measured in coulombs (C)
Charge of a single electron is 1.602x10-19C
Faraday constant (F) : 9.649x104C is the charge of a mole of
electrons
Relation between charge
and moles
Coulombs
2.)
moles
Electric current
Quantity of charge flowing each second
through a circuit
-Ampere: unit of current (C/sec)
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3.)
Electric Potential (E)
Measured in volts (V)
Work (energy) needed when moving an electric
charge from one point to another
-Measure of force pushing on electrons

 Relation between free energy, work and
voltage:
Joules
Volts
Higher potential difference
Coulombs
Higher potential difference requires more work to
lift water (electrons) to higher trough
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4.)
Electric Potential (E)
Combining definition of electrical charge and potential
Describes the voltage that can be generated by a chemical reaction
5.)
Ohm’s Law
Current (I) is directly proportional to the potential difference (voltage)
across a circuit and inversely proportional to the resistance (R)
-Ohms (W) - units of resistance
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6.)
Power (P)
Work done per unit time
-Units: joules per second J/sec or watts (W)
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Electrochemical Cells
There are two types of electrochemical cells
 Galvanic cell: ones that spontaneous produce electrical energy.
 Electrolytic cell: ones that consume electrical energy.
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Electrochemical cells
1.) Galvanic or Voltaic cell

Spontaneous chemical reaction to generate electricity
-One reagent oxidized the other reduced
-two reagents cannot be in contact

Electrons flow from reducing agent to oxidizing agent
-Flow through external circuit to go from one reagent to the other
Reduction:
Oxidation:
Net Reaction:
AgCl(s) is reduced to Ag(s)
Ag deposited on electrode and Clgoes into solution
Cd(s) is oxidized to Cd2+
Cd2+ goes into solution
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School of Materials Science and Engineering
Galvanic or Voltaic cell
 Daniell cell
Salt Bridge
Connects & separates two
half-cell reactions
Prevents charge build-up and
allows counter-ion migration

-0.76 V
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
A potential difference between two electrodes represents a tendency for the reaction to o
ccur.
• Chemical energy  electrical energy
• Spontaneous: (so Ecell is positive)
Example:
 Primary (non-rechargeable)
• Le Clanche (dry cell)
 Secondary (rechargeable)
• Lead storage battery
 Hydrogen-Oxygen Fuel Cell
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Exercise1
Consider a simple galvanic cell consisting of two beakers connected by a salt bridge.
One beaker contains a solution of MnO4− in dilute sulfuric acid and has a Pt electrode. The other beaker
contains a solution of Sn+2 in dilute sulfuric acid, also with a Pt electrode. When the two electrodes are
connected by a wire, current flows and a spontaneous reaction occurs that is described by the following
balanced chemical equation:
2MnO4−(aq) + 5Sn+2 (aq) + 16H+(aq) → 2Mn+2 (aq) + 5Sn+4 (aq) + 8H2O(l)
For this galvanic cell,
a. write the half-reaction that occurs at each electrode.
b. indicate which electrode is the cathode and which is the anode.
c. indicate which electrode is positive and which is negative.
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Applications of Galvanic Cells- Lead-Acid Battery
Reactions
PbO2 + 4H+ + 2e- + SO4-2  PbSO4 + 2H2O
Pb(s) + SO4-2  PbSO4(s) + 2e-
E° = 1.685V
E° = 0.356 V
Total reaction
PbO2(s) + Pb(s) + 2H2SO4  2 PbSO4(s) + 2 H2O E° = 2 V
Lead-acid batteries
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Electrolytic cell
 Electrolytic cell: is a cell in which current flows, power is consumed, and the cell
reaction being driven is the reverse of the spontaneous cell reaction.
Cathode (-)
Anode (+)
2 H2O
O2(g) + 4
H+
+
4 H2O + 4e-
4e-
2H2 + 4 OH-
• Electrical energy  chemical energy
• Non-spontaneous: (Ecell is negative)
Example:
• Electrolysis of water
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 Comparison between electrolytic and galvanic cells
Energy
Energy
• Energy is released from spontaneous redox reaction
• Anode is negative
• Energy is absorbed to drive nonspontaneous redox reaction.
• Anode is positive
In both voltaic and electrolytic cells, electrons flow from the anode to the cathode in the external
circuit
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Schematic Representation of Electrochemical Cells
 The anode is always on the left
: boundaries
: salt bridge
For copper-zinc cell:
Other conditions like concentration are listed just after each species
 The cell reaction
H2 (g) + Cu2+ (aq)  2 H+ (aq) + Cu (s)
Pt H2 (g) H+ (aq) Cu2+ (aq) Cu (s) Pt
Yeungnam University
School of Materials Science and Engineering
Cell potential, Eo
For Zn/Cu, voltage is 1.10 V at 25°C and when
[Zn2+] and [Cu2+] = 1.0 M.
 This is the
Standard cell potential, Eo
 Eo is a quantitative measure of the tendency of reactants to proceed
to products when all are in their standard states at 25 °C.
Eo and Go
Eo is related to Go, the free energy change for the reaction.
Go = - n F Eo
F = Faraday constant = 9.6485 x 104 J/V•mol
n = the number of moles of electrons transferred.
Yeungnam University
School of Materials Science and Engineering
Standard cell potential, Eo
• Can’t measure half- reaction Eo directly. Therefore, measure it relative
to a standard half cell: the Standard Hydrogen Electrode (SHE)
2 H+(aq, 1 M) + 2e-
H2(g, 1 atm)
Eo = 0.0 V
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School of Materials Science and Engineering
Standard Potentials

When combining two ½ cell reaction together to get a complete net reaction, the total cell
potential (Ecell) is given by:
E0cell = E0cathode - E0anode
 Predict voltage observed when two half-cells are connected
• Reactions always written as reduction
Potentials measured at standard conditions
• All concentrations (or activities) = 1M
• 25oC, 1 atm pressure
Best Oxidizing agent?
Best Reducing agent ?
As Eo increases, the more favorable the reaction and the more easily the compound is reduced
(better oxidizing agent).
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 E0cell > 0
G0 < 0
Spontaneous reaction
proceeds to right (products)
 E0cell < 0
G0 > 0
Non-spontaneous reaction
proceeds to left (reactants)
 E0cell = 0
G0 = 0
Equilibrium reaction
Exercise?
Br2(aq)+2V3+ +2H2O(l)  2VO2+(aq)+ 4H+(aq)+ 2Br-(aq)
Given:
E0cell = +1.39 V
E0Br2 = +1.07 V
What is E0V3+ and is the reaction spontaneous?
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School of Materials Science and Engineering
• Eo only applies to [ ] = 1 M for all aqueous species
• at other concentrations, the cell potential differs
• Ecell can be predicted by Nernst equation
E = Eo -
RT
nF
ln (Q)
n = # e- transferred
F = Faraday’s constant
= 9.6485 x 104 J/V•mol
Zn / Zn2+ (0.5 M) // Cu2+ (2.0 M) / Cu
Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)
[Zn2+]
Q=
[Cu2+]
E = 1.10 - (0.0257) ln ( [Zn2+]/[Cu2+] )
2
E = 1.10 - (-0.018) = 1.118 V
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School of Materials Science and Engineering
Electrodes
1) Working electrode
 In corrosion testing, the working electrode is a sample of the corroding metal.
 The working electrode can be bare or coated metal.
2) Auxiliary (counter) electrode
 The auxiliary electrode is a conductor that completes the cell circuit.
 Such as an inert conductor like platinum or graphite
 The current that flows into the solution via the working electrode leaves the solution via the
auxiliary electrode.
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3) Reference electrode
A reference electrode is used in measuring the working electrode potential. A reference electrode
should have a constant electrochemical potential as long as no current flows through it.
 Standard hydrogen electrode (SHE)
The standard reference electrode which is often referred to is the standard hydrogen electrode
(SHE). Al other reference electrodes can be expressed in terms of some constant deviation
from SHE. The SHE reaction is based upon the reaction:
2H   2e   H 2 g 
Eo =0.000VSHE
Where Eo is the reference state at standard temperature and pressure. It should be noted that
this potential is independent of temperature.
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 Reference electrode
 Saturated calomel electrode(SCE)
The saturated calomel electrode consists of pure mercury covering a platinum wire which
passes through a sealed glass tube. The mercury is covered with mercurous chloride and imme
rse in saturated potassium chloride. The temperature coefficient is -6.6x 10 -4 V/oC
Hg 2Cl2  2e   2 Hg  2Cl 
Eo =0.268VSHE
 Silver-silver chloride reference electrode:
Ag-AgCl electrode consits of a silver, the surface which has been chloridized (transformed to silve
r chloride), typically in dilute hydrochloric acid,. The temperature coefficient is –4.3x 10 -4 V/oC .
Eo = 0.222 VSHE
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Potentials of common reference electrodes
Common Name
Electrode
V vs NHE
Saturated Calomel Electrode (SCE)
Hg/Hg2Cl2/sat. KCl
+0.241
Calomel
Hg/Hg2Cl2/1M KCl
+0.280
Mercurous sulphate
Hg/Hg2SO4/sat. K2SO4
+0.640
Mercurous oxide
Hg/HgO/1M NaOH
+0.098
Silver chloride
Ag/AgCl/sat. KCl
+0.197
Copper sulphate
Cu/sat. CuSO 4
+0.316
Zinc in seawater
Zn/seawater
~ -0.8
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References
(1) Modern electrochemistry- 2nd edition-John O’M. Bocker and Amulya K.N. Reddy
(2) Electrochemical methods ; fundamental and applications- Allen J. Bard and Larry R. Faulkner
Yeungnam University
School of Materials Science and Engineering