1. The document discusses different types of electrochemical cells including galvanic/voltaic cells and electrolytic cells. 2. Galvanic cells are further classified as primary cells, which cannot be recharged, and secondary cells, which are rechargeable batteries. 3. The Nernst equation is derived, which relates the electrode potential to the standard electrode potential and the concentrations of the metal ions involved in the electrochemical cell reaction.

1. MODULE – I
I/II SEM BE , ENGINEERING CHEMISTRY
BY
RASHMI M
ASSISTANT PROFESSOR
SKIT
BANGALORE

3. MODULE I
ELECTRODE POTENTIAL & CELLS
Electrochemistry: It is a branch of chemistry which deals with the study of
transformation of chemical energy into electrical energy and vice versa”
Electrochemical cell and Classification with examples.
An electrochemical cell is a device, which is used to convert chemical energy
into electrical energy and vice versa.
These electrochemical cells are classified into two types as follows’
1) Galvanic or Voltaic cells: These are the electrochemical cells, which
converts chemical energy into electrical energy.
Ex. Daniel cell, Dry cell, etc
2) Electrolytic cells-are devices which convert electrical energy into
chemical energy.
Example: Electrolysis of molten NaCl, Recharge process of lead acid battery

4. Galvanic or Voltaic cells:
Galvanic or Voltaic cells are again classified into three types as follows
a) Primary cells: These are the cells which serve as a source of energy only
as long as the active chemical species are present in the cell. The cell
reactions are irreversible. These are designed for only single discharge and
cannot be charged again.
Ex: Dry Cell, Zn – Hgo cell, Zn-Ag2o cell etc.
b) Secondary cells: These cells are chargeable and can be used again and
again. The cell reactions are reversible and are often called as reversible
cells. During discharging the cells acts like voltaic cell converting chemical
energy into electrical energy. During charging the cell acts like electrolytic
cell by converting electric energy into chemical energy, hence these batteries
are called as storage battery.
Ex: Lead acid Battery, Ni-cd cells. Lithium ion cells etc.
c) Concentration of cells:
These are the electrochemical cells consisting of same metal electrodes
dipped in same metal ionic solution in both the half cells but are different in
the concentration of the metal ions.
Ex: Cu/Cu2+ (M1) || Cu2+ (M2)/Cu
Ex: copper concentration cell, Zinc concentration cell

5. Oxidation: A species loses one or more electrons resulting in the increase in its
oxidation number.
Reduction: A species gain one or more electrons resulting in a decreasing in its
oxidation number.
Oxidation should accompanied by reduction, because if one losses electrons another
must ready to accept electrons. This reaction is called redox reaction.
Single electrode Potential:
It is defined as the potential developed at the interphase between the metal and the
solution, when a metal is dipped in a solution containing its own ions. It is
represented as E
Standard reduction potential (Eo) :
It is defined as potential developed at the interface between the metal and the
solution, when a metal is dipped in a solution containing its own ions of unit
concentration at 298K. [If the electrodes involve gases then it is one atmospheric
pressure] It is denoted as E0.
Electromotive force (EMF):
It is defined as the potential difference between the two electrodes of a galvanic cell
which causes the flow of current from an electrode with higher reduction potential to
the electrode with lower reduction potential.
It is denoted as E cell.
E cell = E right –E left.
E cell = E cathode – E anode.

6. Electrochemical series:
The arrangement of elements in the increasing order of their standard reduction
potential is referred to as emf or electrochemical series. Such a arrangement of few
elements given in the table
Mn+/M Eo (volts) Mn+/M Eo (volts)
Li+/Li -3.05 H+/H2 0.00
K+/K -2.93 Sn4+/ Sn2+ +0.15
Mg+/Mg -2.37 Cu2+/Cu +0.16
Al3+/Al -1.66 Cu2+/Cu +0.34
Zn2+/Zn -0.76 Cu+/Cu +0.52
Cr3+/Cr -0.74 I2/I- +0.54
Fe2+/Fe -0.44 Fe3+/Fe2+ +0.77
Cr3+/Cr2+ -0.41 Hg2+/Hg+ +0.79
Cd2+/Cd -0.40 Ag+/Ag +0.80
Ni2+/Ni -0.25 Hg2+/Hg +0.85
Sn2+/Sn -0.14 Pt2+/Pt +1.20
Pb2+/Pb -0.13 Cr7+/Cr3+ +1.31
Fe3+/Fe2+ -0.041 Cl2/2Cl- +1.36
Au3+/Au +1.50

7. • A negative value indicates oxidation tendency and a
positive value indicates reduction tendency with
respect to hydrogen.
• The metal with lower electrode potential is more
reactive and as the electrode potential increases, the
reactivity decreases, and metals with higher electrode
potentials are nobler.
• Metals with lower electrode potentials have the
tendency to replace metals with higher electrode
potential from their solutions for example, Zn
displaces Cu, and Cu displaces Ag
• Metals with negative electrode potentials can liberate
hydrogen from acidic solutions

8. Derivation of Nernst Equation for Electrode
potential
In 1889 Nernst derived a quantitative relationship between
the electrode potential and the concentrations of metal ions are
involved. The maximum work available from a reversible
chemical process is equal to the maximum amount of electrical
energy that can be obtained; it shows decrease in free energy.
Wmax = – ∆G--------[1]
And
Wmax = difference in potential between two electrode x
total quantity of charge flowing
through the cell
Total quantity of charge flowing through the cell = (No. of
moles of electrons) x (Faradays
constant)
So Wmax = nFEcell --------[2]

9. c
0
0
c
KlnGG
,isotherm'reactionhoffvant'abyrelatedareGandG,K
]n[M
[M]
c
K
aswrittenbecan
c
Kconstantequlibrium,reactionabovefor the
M-nenM
reaction,electrodereversibleaconsider
0nFE-0Gion,std.conditunder
[3]-----nFE-G
[2]&[1]eqnequate
RT







10. ]6[
]n[M
1
logRT0EE
1[M]condition,standardunder
]n[M
[M]
lnRT0EE
nF-bysidesboth thedivide
]n[M
[M]
ln
0
-nFEnFE-
equation,abovetheto
0
GandG,cKofvaluesthesubstitute








RT

11. Where,
E = Electrode potential
E0 = standard electrode potential
n = no. of electrons
[Mn+] = Concentration of metal ions
R = Universal gas constant = 8.314J K-1 mol-1
T = Temperature (In Kelvin) = 298K
]nlog[M
n
0.05910EE
eqn[6],thetovaluesthesubstitute


12. Concentration Cells
• Concentration cell is an electrochemical cell that has two
equivalent half-cells of the same material differing only in
concentrations. One can calculate the potential developed by
such a cell using the Nernst Equation similar to an
electrochemical cell.
• In concentration cell the voltage comes from its attempts to
reach equilibrium, which will stop when the concentration in
both half-cells are equal.
Types of concentration cells:
The concentration differences could be affected in the electrode
material or in the electrolyte.
1. Electrode concentration cells
2. Electrolyte concentration cells

13. Electrolyte concentrations cells
“A concentration cell is an electrochemical cell in which
electrode materials and electrolytes of two half cells are
composed of same material but the concentration of two
solutions are different”
Ex- Cu/Cu2+(M1)|| Cu2+(M2)/Cu
A concentration cell consists of two same metal electrode
dipped into their own ionic solutions of two different
concentration. Thus in a concentration cell, the electrode with
lower electrolyte concentration acts as anode and the one with
higher electrolyte concentration acts as cathode. The
concentration of ions at anode increases and at cathode
decreases, when the cell is in operation.

14. Consider two copper rods are dipped into their own ionic
solutions of M1 and M2 and it is represented as
Cu/Cu2+
(M1) Cu2+
(M2)/Cu
By electrochemical conventions, if M2 > M1 then, we have the
following reactions.
At anode:
At cathode:



 eMCu
s
Cu 2)1(
2
)(
)(
2)2(
2
s
CueMCu 




15. 1
2
log
2
0591.0
1log
2
0591.002log
2
0591.00
bewillcellionconcentrattheofemfThe
M
M
cell
E
MEME
cell
E
anodecathodecell EEE







The emf of the cell is + ve only if M2 > M1
The following characteristics of concentration cell can be
noted:
• When M2 = M1, the concentration cell does not generate
electrical energy.
• When M2 > M1, the Ecell is + ve.
• When M2<M1, Ecell is – ve.
• Higher the ratio of M2/M1, greater is the cell potential.

16. 


 eCd
M
Cd 2
2
)1(
)2(
2
2
M
CdeCd 



Electrode concentration cells
In these cells, the potential difference is developed between two electrodes
at different concentrations dipped in the same solution of the electrolyte.
For example,
Cd-Hg (M1) | CdSO4 | Cd-Hg (M2)
Two Cd-Hg electrodes of different concentration immersed in a CdSO4 solution.
Electrode containing high concentration of metal acts as a anode , where oxidation
occurs
Electrode containing low concentration of metal acts as a cathode, where reduction
occurs
M2.M1only whenveisEcell
]2[
]1[
log
2
0591.0
]1log[
2
0591.00]2log[
2
0591.00
bewillcellionconcentrattheofemfThe








M
M
cell
E
MEME
cell
E
anodecathodecell EEE

17. Types of electrodes:
• Metal-Metal ion electrode: An electrode of this type consists of a metal
dipped in a solution containing its ions.
Ex- Zn/Zn2+, Cu/Cu2+ etc.
• Metal-Metal salt ion electrode: These electrodes consist of a metal is in
contact with a sparingly soluble salt of the same metal dipped in a solution
containing anion of the salt.
Example-Calomel (Hg|Hg2Cl2|Cl- , Silver- Silver salt electrode (Ag| AgCl |Cl-
• Gas electrode: Gas electrode consists of a gas bubbling about an inert
metal wire, immersed in solution containing ions to which the gas is
reversible. The metal provides electrical contact and facilitates the
establishment of equilibrium between the gas and its ions.
Example-Hydrogen electrode (Pt|H2|H+) , Chlorine electrode (Pt|Cl2|Cl-)

18. • Oxidation-Reduction electrode: An oxidation-reduction
electrode is a one in which the electrode potential arises from
the presence of oxidized and reduced forms of the same
substance in solution. The potential arises from the tendency of
one form changes into the other more stable form. The
potential developed is picked up by an inert electrode like
platinum.
Example-Pt|Fe2+, Fe3+ Pt|Ce3+, Ce4+
• Ion selective electrode: In ion selective electrode, a
membrane is in contact with a solution, with which it can
exchange ions.
Example- Glass electrode.
• amalgam electrode: Metal- amalgum is in contact with a
solution containing its own ions.
Example: lead amalgum electrode ( Pb – Hg/Pb2+)

19. Reference Electrodes:
• Reference electrode: “Reference electrode are the electrode with reference to
those, the electrode potential of any electrode can be measured.” It can acts both as
an anode or cathode depending upon the nature of other electrode.
The Reference Electrodes can be classified in to two types
i) Primary reference electrodes Ex: Standard hydrogen electrode
ii) Secondary reference electrodes Ex: Calomel and Ag/Agcl electrodes
Construction and working of Standard Calomel
electrode (SCE):
1. Calomel electrode is a metal-metal salt Ion electrode.
2. It consists of mercury, mercurous Chloride and a solution of KCl. Mercury is
placed at the bottom of a glass tube.
3. A paste of mercury and mercurous chloride Is placed above the mercury. The
space above the paste is filled with a KCl solution of known concentration.
4. A platinum wire is kept immersed into the mercury to obtain electrical contact.
5. Calomel electrode can be represented as,
Hg | Hg2Cl2 | KCl(sat)

20. • The calomel electrode can acts as anode or cathode
depending on the nature of the other electrode of the
cell.
Saturated KCl
Mercury
Calomel paste
Pt wire
Porous disc

21. Therefore electrode potential of calomel electrode is
depending upon the concentration of KCl. The electrode is
reversible with chloride ions.
 
 
  298Katlog0591.0
2nwhere,log
303.20
2
log
303.20
-2Cl2Hg(l)-2e)(2Cl2Hg
is,reactionelectrodereversiblecellnetThe
0 




ClEE
Cl
F
RT
EE
Cl
nF
RT
EE
s

22. Concentration
of KCl
E0
[v]
saturated KCl 0.241
1M KCl 0.281
0.1M KCl 0.33
4
MEASUREMENT OF ELECTRODE POTENTIAL USING CALOMEL ELECTRODE:
Electrode potential of a given electrode can be measured by using calomel electrode as a
reference electrode.
Example: To measure the electrode potential of zinc, Zinc electrode is coupled with SCE. So zinc
acts as anode and SCE acts as cathode
cellZnZn
cellcathodeanode
anodecathodecell
EE
EEE
EEE



 241.02
/
Applications:
1. It is used as secondary reference electrode in the measurement of single electrode.
2. It is used as reference electrode in all potentiometer determinations and
to measure pH of the given solution

23. Construction and working of Silver- Silver Chloride
electrode:
• Silver-Silver chloride is also a metal-metal salt ion electrode.
• Silver and its sparingly soluble salt silver chlorides are in
contact with a solution of chloride ions. Generally a silver wire
is coated with AgCl and dipped in a solution of KCl .
• Cell representation is as follows
Ag |AgCl | Cl-
Concentration of KCl E0 [v]
saturated KCl 0.241
1M KCl 0.281
0.1M KCl 0.334

24.  
  298Katlog0591.0
1nwhere,log
303.20
-ClAg-2e)(AgCl
is,reactionelectrodereversiblecellnetThe
0 



ClEE
Cl
nF
RT
EE
s
Therefore electrode potential of calomel electrode is
depending upon the concentration of KCl. The
electrode is reversible with chloride ions.
Applications:
•Used as secondary reference electrode in ion selective elctrode.

25. Construction and working of Ion selective electrode
(ISE) :
“Ion selective electrode is one which selectively responds to a specific
ion in a mixture and the potential developed at the electrode is a
function of the concentration of that ion in the solution”
Glass electrode: A glass electrode is an ion selective electrode where
potential depends upon the pH of the medium.
1. The glass electrode consists of a glass bulb made up of special type of
glass (sodium silicate type of glass) with high electrical conductance.
2. The glass bulb is filled with a solution of constant pH (0.1MHCl) and
insert with a Ag-AgCl electrode, which is the Internal reference
electrode and also serves for the external electrical contact.
3. The electrode dipped in a solution containing H+ ions as shown in the
figure.
The electrode representation is,
Glass | 0.1M HCl | Ag-AgCl

27. WORKING:
• The glass electrode works on the principle that when a thin
glass membrane is placed between two different concentration
of a solution, a boundary potential Eb is developed at layers of
the glass membrane. This potential arises due to difference in
the concentration of H+ ion inside and outside the membrane.
External Solution glass membrane Internal solution
C2=[H+] C1=[CONSTANT]=k
E2 Eb E1
• Boundary potential, Eb = E2 – E1

28. 12
1
0
2
0
21
21
2
0
2
log0591.0log0591.0
1nwherelog
0591.0
log
0591.0
getwe(1),eqnintoE&EofvaluestheSubstitute
membrane.outerandinnerions+HofionconcentratareCandCWhere,
log
0591.0
1Clog
n
0.05910E1E
CCE
C
n
EC
n
EE
C
n
EE
b
b










 
10.0591logC-=K=ConstWhere
log
0591.0
(H+))=(C2)&constantisC1constant,issolutioninnertheofionconcentratthe(Since

 H
n
ConstEb

29. asyAgClAgG
GG
asyAgClAgG
asyAgClAgbG
EEKEwhere
pHEE
EEpHKE
EEEE





/
0
0
/
/
b
0591.0
0591.0
potential.EasycalledexistsispotentialadditionalsmallaHowever,C2.=C1for0=Eblly,Theoritica
Easy.potentialAsymmetryiii)
andEAg/AgClpotentialelectrodeAgCl-Agii)
Eb,potentialboundaryi)
theofsumispotentialelectrodeglassThe
pH
n
0.0591
KE

31. Determination of pH using glass electrode:
0591.0
0591.0
0591.0
evaluated.becanglassE
cell,theemfknowingisESince
anodeEcathodeEcellE
0
0
0
SCE
SCEcellG
SCEG
GG
SCEGcell
EEE
pH
EpHEEcell
pHEE
EEE







32. Advantages
• This electrode can be used to determine PH in the range 0-
9, with special type of glass even up to 12 can be
calculated.
• It can be used even in the case of strong oxidizing agents.
• The equilibrium is reached quickly.
• It is simple to operate, hence extensively used in various
laboratories.
Limitations
• The glass membrane though it is very thin, it offers high
resistance. Therefore ordinary potentiometers cannot be
used; hence it is necessary to use electronic
potentiometers.
• This electrode cannot be used to determine the PH above 12

33. 1. Two Copper electrodes placed in CuSO4 solutions of equal concentration are
connected to form o concentration cell.
a) What is the cell voltage?
b) If one of the solutions is diluted until the concentration of Cu2+ ions is 1/5th of its
original value. What will be the cell voltage after dilution?
Solution:
• The cell potential of concentration cell is given as
• When the concentration of the species are gqual (C2 = C1) the cell voltage is zero.
• When one of the solution is diluted to of its original value, C2 =1M & C1
1
2
log
0591.0
C
C
ncell
E 
VE
E
E
E
cell
cell
cell
cell
0206.0
699.002955.0
5log
2
0591.0
5/1
1
log
2
0591.0
C1
C2
log
n
0.0591
cellE






34. 2.Calculate the emf of the given concentration cell at 298K.
Ag(s) /AgNO 3 (0.018M) AgNO3 (1.2M) /Ag.
Solution :
V
cell
E
cell
E
cell
E
C
C
ncell
E
1708.0
66.66log
1
0591.0
018.0
2.1
log
1
0591.0
1
2
log
0591.0



