Electrochemistry:Introduction – cells – types - representation of galvanic cell - electrode potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from single electrode potential - reference electrode: construction, working and applications (Determination of potential of the unknown electrode and pH of the unknown electrode) of standard hydrogen electrode, standard calomel electrode - glass electrode – EMF series and its applications - potentiometric titrations (redox) - conductometric titrations - mixture of weak and strong acid vs strong base.

1. DEPARTMENT OF CHEMISTRY
WELCOMES YOUALL
22CYT12 &Chemistry for Computer Systems
(Electrochemical series and its Applications)
Prepared By
Krishnaveni K
Assistant Professor
Department of Chemistry
Kongu Engineering College,
Perundurai, Erode-638060

2. ELECTROCHEMISTRY
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3. Introduction – cells – types - representation of galvanic cell - electrode
potential - Nernst equation (derivation of cell EMF) - calculation of cell EMF from
single electrode potential - reference electrode: construction, working and
applications of standard hydrogen electrode, standard calomel electrode - glass
electrode – EMF series and its applications - potentiometric titrations (redox) -
conductometric titrations - mixture of weak and strong acid vs strong base.
UNIT-I
ELECTROCHEMISTRY

4. History of Electrochemistry
 16 th Century - William Gilbert –Father of Magnetism
 18 th Century – William Nicholson & Wilhelm Ritter – Decomposition of water – Electrolysis
 Svante Arrhenius - Dissociation of electrolytes
 Walther Hermann Nernst – Theory of Electromotive Force
 Conductance? Ability to conduct current , mho

5. ELECTROCHEMISTRY
INTRODUCTION
 It is a branch of chemistry
 The study of process involved in the interconversion of
chemical and electrical energy.
KEY TERMS IN ELECTROCHEMISTRY
 Conductor: Material which conduct electric current
 Non conductor: Material which do not conduct electric current
 Current: The flow of electrons through a wire or any conductor
 Oxidation: Loss of electrons
 Reduction: Gain of electrons
 Redox reaction: oxidation and reduction reactions occur simultaneously
 Reducing agent: A reactant in which donates an electron to the reduced species. (The reducing agent
is oxidized)
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6.  Oxidizing agent: A reactant in which accepts an electron from the oxidized species. (The oxidizing agent
is reduced)
 Anode: The electrode at which oxidation occurs
 Cathode: The electrode at which reduction occurs
 Electrolyte: A water soluble substance and conduct an electric current
 Half cell: A single electrode immersed in an electrolytic solution and developing a definite potential
difference.
 Cell: Two half cells are connected through one wire
 Oxidation Potential : It is the tendency of an electrode to loss electrons
 Reduction potential: It is the tendency of an electrode to gain electrons
 Electrode Potential: It is the tendency of an electrode to loss or gain electrons
 Single Electrode Potential: It is the tendency of an electrode to loss or gain electrons when it is dipped in
its own salt solution. (Standard- 1M concentration at 250C).
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7. 17-Dec-22
LEOGER BOARD

9. ELECTROCHEMICAL CELL
Introduction
An electrochemical cell is a device in
which a redox reaction is utilized to get
electrical energy.
An electrochemical cell is also commonly
referred to as voltaic or galvanic cell.
The electrode where reduction occurs is
called cathode.
The electrode where oxidation occurs is
called anode.
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10. Construction
 Electrochemical Cells are made up of two half-cells, each consisting of an electrode
which is dipped in an electrolyte. The same electrolyte can be used for both half cells.
These half cells are connected by a salt bridge which provides the platform for ionic
contact between them. A salt bridge minimizes or eliminates the liquid junction
potential.
 The practical application of an electrochemical or galvanic cell is the Daniel cell.
 It consists of a Zn electrode dipping in ZnSO4 solution and a Cu electrode dipping in
CuSO4 solution.
EMF= Eoxi + E Red
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11. Cell reaction
 Anode : Zn → Zn2+ + 2e- (Oxidation) {0.76V}
 Cathode : Cu2+ + 2e- → Cu (reduction) {0.34V}
 Overall : Zn + Cu2+ → Zn2+ + Cu (Redox)
 Representation of Daniel cell : Zn / Zn2+ || Cu2+ / Cu
 Zn / ZnSO4 (1M) // CuSO4 (1M) / Cu
 Cell EMF : 1.1 V
 EMF= Eoxi + E Red
= EZn + Ecu = 0.76+0.34
 CuSO4 - Cu2+ + SO4
2-

12. Electrolytic cells
 Electrical Energy -- Chemical Energy
 Anode  positive Charge - oxidation --- 2Cl-  Cl2 + 2e-
 Cathode  negative charge  reduction --- 2Na+ + 2e-  Na
 Overall reaction --- 2Na+ + 2Cl-  2NaCl

13. Representation of Galvanic Cell
Anode : Zn Cathode : Cu
Zn Zn2+ ZnSO4 CuSO4 Cu2+ Cu
Metal and the electrolyte or metal ion can be separated by , / ;
Zn / Zn2+ Cu2+ / Cu
Concentration of the electrolyte should be in ()
Zn / ZnSO4 (1M) CuSO4 (1M) / Cu
Zn , Zn2+ Cu2+ , Cu
Salt bridge can be represented by ||
Zn / Zn2+ || or // Cu2+ / Cu

14. Electrochemical Series
 The standard electrode potentials of a number of electrodes are arranged in the
increasing order of reduction potential at 25°C is referred to as emf or electrochemical
series.
Characteristics of electrochemical series:
 Lithium is the first member of the series.
 Highly reactive metal systems are at the top of the series.
 In other words, good reducing agents are at the top of the series, having the negative sign and act as
anode.
 All good oxidizing agents are at the bottom of the series , having the positive sign and act as cathode.
 Hydrogen system is at the middle of the series. All the elements which displace hydrogen from dilute
acids are placed above it.

17. Applications of Electrochemical Series
 To Find Reactivity of Metals
 As we move down in the electrochemical series reactivity of metal
decreases
 Alkali metals and alkaline earth metals at the top are highly reactive.
They can react with cold water and evolve hydrogen. They dissolve in
acids forming salts.
 Metals like Fe, Pb, Sn, Ni and Co which lie a little down in the series,
do not react with cold water but react with steam and evolve hydrogen.
 Metals like Cu, Ag and Au which lie below the hydrogen are less
reactive and do not evolve hydrogen from water.

18. For Studying displacement reaction
 Elements having higher reduction potential will gain electrons and that having lower
reduction potential will lose electrons. Hence element higher in electrochemical series
can displace an element placed lower in electrochemical series from its salt solution.
Example
Can zinc displaces copper from its salt solution?
Zn displaces Cu from CuSO4, because, zinc is placed higher in electrochemical series
while Cu is placed lower in electrochemical series. Hence zinc can easily displace
copper from CuSO4.
Zn+CuSO4 --------> ZnSO4 + Cu E0
Zn = -0.76 volts
Cu+ZnSO4 --------> No recation E0
Cu = +0.34 volts

19. For choosing elements as Oxidizing Agents
 The elements which have more electron-accepting tendency are oxidizing agents. The
strength of an oxidizing agent increases as the value of reduction potential becomes more
and more positive. Elements at the bottom of the electrochemical series have higher (+ve)
reduction potential. So they are good oxidizing agents. Thus, oxidizing power increases
from top to bottom in the series.
Example- F2 is a stronger oxidant than Cl2, Br2 and I2.
Cl2 is a stronger oxidant than Br2 and I2.

20. For choosing elements as Reducing Agents
The elements which have more electron losing tendency are reducing agents. The
power of reducing agent increases as the value of reduction potential becomes more and
more negative. Elements at the top of the electrochemical series have higher (-ve)
reduction potential. So they are good reducing agents. Thus, reducing power decreases
from top to bottom in the series.
Example-
The element like Zn, K, Na, Fe, etc. are good reducing agent.

21. Displacement of hydrogen from dilute acids by metals
 The metal which can provide electrons to H+ ions present in dilute acids for reduction evolve hydrogen
from dilute acids. The metal having negative values of reduction potential possesses the property of
losing an electron or electrons.
 Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from
dilute acids and on descending in the series, tendency to liberate hydrogen gas from dilute acids
decreases.
 The metals which are below hydrogen in the electrochemical series like Cu, Hg, Au and Pt do not evolve
hydrogen from dilute acids.
Example
Zinc reacts with dil.H2SO4 to give H2 but Ag does not. Why?
Zn+H2SO4 --------> ZnSO4 + H2 ; E0
Zn = -0.76 volts
Ag+H2SO4 --------> No reaction; E0
Ag = +0.80 volts
The metal with a positive reduction potential will not displace hydrogen from an acid solution.

22.  Displacement of hydrogen from water
 Iron and the metals above iron are capable of liberating hydrogen from water. The tendency
decreases from top to bottom in the electrochemical series.
 Alkali metals and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe
liberate hydrogen from hot water or steam.
 For Calculation of Standard emf of the cell
Standard reduction potential values are given in emf series. From the values E0
cell is calculated
using formula
E0
cell or standard emf of a cell = E0
oxi(cathode) - E0
red(anode)

23. Calculation of standard EMF of the cell
 EMF= Eoxi + E Red
 Zn & Cu Couple
 EMF= Eoxi + E Red
= EZn + E Cu
= 0.76+ + 0.34
= 1.1V
 Fe & H2
 EMF= Eoxi + E Red
 EMF= EFe + E H2
 = 0.441+ 0
0.441V

24.  Ni & Hg Couple
 Ni – Anode
 Hg - Cathode
 EMF= Eoxi + E Red
 = ENi + E Hg
= 0.236 + 0.61= 0.846V
 EMF = Standard reduction potential of R.H.S electrode- Standard reduction potential of L.H.S
electrode
 E0 = E0
RHS - E0
LHS
 = E0
Hg- E0
Ni
 = 0.61 – (-0.236)
 = 0.61+0.236 = 0.846V
= ENi + E Hg
= 0.236+0.61
= 0.846V

25. Cr & Sn Couple
Cr – Anode
Sn - Cathode
EMF= Eoxi + E Red
EMF= ECr + E Sn
= -0.74+(-0.14)
= 0.60V
EMF = Standard reduction potential of R.H.S electrode-
Standard reduction potential of L.H.S electrode
E0 = E0
RHS - E0
LHS
= E0
Sn - E0
Cr
= – 0.14 -(-0.74)
= -0.14+0.74 = 0.60V

26. For predicting spontaneity of the cell reaction
E0
cell > 0 cell reaction is spontaneous
E0
cell < 0 cell reaction is non-spontaneous
E0
cell = 0 cell reaction is in equilibrium
For determination of equilibrium constant for a reaction
We know that
-∆G0 = RTlnK
= 2.303RT logK
log K =
log K = (-∆G0 = nFE0)
Thus, from the value of E0 for a cell reaction, its equilibrium constant can be calculated.

27. REFERENCES:
 1.Palanisamy P.N., Manikandan P., Geetha A.& Manjula Rani K, “Applied
Chemistry”, 6th Edition, Tata McGraw Hill Education Private Limited, New
Delhi, 2019.
 2 .Paya Payal B.Joshi, Shashank Deep., “Engineering Chemistry”, Oxford
University Press, New Delhi, 2019.
 3.Palanna O., “Engineering Chemistry”, McGraw Hill Education, New
Delhi, 2017.
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28. 17-Dec-22
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