Electrochemistry

1. CHAPTER 1: ELECTROCHEMISTRY
CHEMISTRY- MODULE I- B.Tech I year
Dr. B. Divya
Assistant Professor
Department of Chemistry
Institute of Aeronautical Engineering (IARE)
Hyderabad

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ELECTROCHEMISTRY
TOPICS
 Introduction to Electrochemistry
 Conductors: Types
 Electrochemical cells
 Electrode potential
 Standard electrode potential
 Determination of Standard electrode potential
 Prediction of spontaneity of reaction

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ELECTROCHEMISTRY
Introduction
 Electrochemistry is the branch of chemistry which deals with the transformation of electrical
energy into chemical energy and vice versa.
 The laws of electrochemistry form the basis of electrolysis and
electrosynthesis.
 The knowledge of electrochemistry is of immense importance to study
about the causes of destruction of materials due to corrosion.
 In electro-chemistry, there are two processes: electrolysis and
electromotive process. Both these processes are interrelated.
Electrical energy causing chemical reaction Electrolysis (Electrolytic cell)
Chemical reactions producing electrical reaction Electromotive (Galvanic
cell)

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ELECTROCHEMISTRY
Conductors
 The substances which allow the passage of
electric current are called conductors. Metals
such as copper and silver are good conductors
of electricity.
Electrical conductors are of two types:
1. Metallic conductors
2. Electrolytic conductors.

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ELECTROCHEMISTRY
Electrolytic Conductors: Substances which allow the electricity to pass through them in their
molten states (ionic) or in the form of their aqueous solutions are called electrolytic conductors
is known as electrolytic conductance.
Eg: sodium chloride, potassium chloride, acids
etc.
Metallic Conductors: The metals which conduct electricity through electrons are called metallic
conductors

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ELECTROCHEMISTRY
Metallic Conductor Electrolytic Conductor
Conductance due to the migration of
Eg: Metals, graphite
Conductance due to the migration of
solution of fused electrolyte
Eg: Acids and bases
Passage of current due to electron
chemical reaction takes place
Passage of current due to movement
Some chemical reactions takes place
Free electrons are responsible for
conduction
Free ions are responsible for electrical
Mass is not transferred Mass is transferred
With increase in temperature
and conductance decreases
With increase in temperature
and conductance increases

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ELECTROCHEMISTRY
Electrolytes are classified into two types:
Strong electrolytes: The electrolytes which completely dissociates in solution at all
in solution at all concentrations. Their conductance is very high.
Eg. NaCl, HCl, NaOH.
Weak Electrolytes: The electrolyte which partially dissociates at moderate concentration. Their
conductance is low as they dissociate only to a small extent even at very high dilutions.
Eg: CH3COOH, NH4OH, sparingly soluble salts like AgCl, AgBr, AgI, BaSO4, PbSO4 etc

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ELECTROCHEMISTRY
Cells or Devices in Electrochemistry
Devices that converts electrical energy chemical energy
Reactions: Redox reaction
Oxidation (OIL) Loss of electrons
Reduction (RIG) Gain of electrons
Conversion of energy occurs through cells
Electrochemical cells
Or Galvanic cells
Or voltaic cells
Electrolytic cells
Chemical cells (change in chemical state – no change of matter)
Concentration cells ( physical change transfer in matter)
Applications:
Biological reactions
Batteries/cells

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ELECTROCHEMISTRY
Galvanic or Voltaic Cells: These are the electrochemical cells, which converts chemical energy
into electrical energy . Electricity is produced from a spontaneous chemical reaction.
Ex: Daniel cell, Dry cell etc.
Electrolytic cells: devices which convert electrical energy into chemical energy.
It is a non-spontaneous chemical reaction.
Ex. Electrolysis of molten NaCl, Recharge process of lead acid battery
Galvanic or Voltaic cells are again classified into 3 types
 Primary cells
 Secondary cells
 Concentration cells

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ELECTROCHEMISTRY
Primary cells: These are the cells which serve as a source of energy only as long as the active
chemical species is present in the cell. The cell reactions are irreversible. These are designated
cannot be charged again
Eg. Dry cell, Zn-HgO cell, Zn-Ag2O cell
Secondary cells: These cells are chargeable and can be used again and again. The cell reactions
are reversible.
Discharging – behaves voltaic cell- chemical to electrical
Charging – electrolytic cell - electrical to chemical
Eg. Lead-acid battery, Ni-Cd cells, lithium ion cells
Concentration cells: same metal dipped in same electrolyte- 2 half cells
Eg. Cu/Cu2+ (M1) // Cu2+ (M2)/Cu (different concentration of metal ions)

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ELECTROCHEMISTRY
Electrochemical Cell Electrolytic Cell
In this cell, chemical energy is
electrical energy
In this cell electrical energy is
chemical energy
Here, anode is –ve electrode and
+ve electrode
Here, anode is +ve electrode and
–ve electrode
Salt bridge is required Salt bridge is not required
This process is reversible and This process is irreversible and
spontaneous
EMF of the cell is +ve EMF of the cell is -ve

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ELECTROCHEMISTRY
Electrode potential
When a piece of metal is placed into a solution containing its own ions, a potential difference
exists between the solution and metal. This potential difference is known as electrode
potential of the metal or of this single electrode system.
It depends on:
 Concentration of ions present in
solution
 Nature of metal electro positivity of
metal
 Temperature of the solution

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ELECTROCHEMISTRY
Single Electrode potential
 Each electrochemical cell is made up of 2 electrodes, at one electrode electrons are evolved
and at other electrode electrons are consumed.
 Each electrode dipped in its salt solution is called Half Cell
 The potential of half-cell is called single electrode potential
Zn/ZnSO4//CuSO4/Cu
Half cell Half cell
Ecell = Eright- Eleft (or) E cathode – E anode

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ELECTROCHEMISTRY
Electrochemical cell (Galvanic Cell)
 It converts the chemical energy into electrical energy
 –ΔG = nFE (spontaneous reaction)
 Electrochemical cell consists of 2 electrodes immersed in
one or more electrolytes and the are connected externally
(Voltameter)
 Chemical reaction occurs in the cell.
 Oxidation at one electrode and electrons given out are
consumed for reduction at other electrode
 Flow of electrons is from anode to cathode, therefore,
electricity from cathode to anode

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ELECTROCHEMISTRY
Example of electrochemical cell is Daniel cell
 It consists of Zn rod dipped in ZnSO4 and a Cu rod dipped in CuSO4
solution
 Zn metal has a greater tendency for oxidation than copper.
Therefore oxidation takes place at Zn compartment. Zinc act as
anode. Zn2+ goes to the solution leaving behind the 2e- on the Zn
rod. Hence Zn rod achieves –ve charge (anode)
 Cu2+ has a greater tendency for reduction than Zn2+, therefore Cu2+
accepts electrons and gets reduced and deposited on copper metal.
Cu acts as cathode
 Both the cells are connected internally through a bent glass tube
having saturated solution of KCl or KNO3 or Na2SO4 or NH4OH in
agar-agar is called salt bridge

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ELECTROCHEMISTRY
 At anode: Zn Zn2+ + 2e- (oxidation)
 At cathode: Cu2+ + 2e- Cu (reduction)
Zn + Cu2+ Zn2+ + Cu
 The flow of electrons will be
externally from anode to
cathode and internally from
cathode to anode through salt
bridge
Working of Daniel cell
 Flow of current is due to difference in
electrode potentials of both the
electrodes
 Potential difference is called EMF and
measured in volts.

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ELECTROCHEMISTRY
 The transfer of electrons from Zn anode to Cu cathode leads to
accumulation of +ve charge around the anode due to the formation of Zn2+
ions and the accumulation of –ve charge around the cathode due to
deposition of Cu2+ ions as copper on the cathode
 The positive charge so accumulated around the anode will prevent the flow
of electrons from it.
 Similarly accumulation of negative charge around the cathode will prevent
the acceptance of electrons from anode.
 As the transfer of electrons stops, the current flow stops

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ELECTROCHEMISTRY
Function of salt bridge
 The flow of current becomes slow after
using the electrodes for a long time due to
polarization of electrodes
 Then salt bridge restores the electrical
neutrality of solution in 2 half cells. When
the concentration of Zn2+ ions increases
sufficient number of Cl- ions migrate from
salt bridge to anode half cell
 Similarly sufficient no. of K+ ions migrate from salt bridge to cathode half
cell from neutralizing excess negative charge due to additional SO4
2- ions in
cathodic cell
 This maintains the electrical neutrality of 2 solutions in half cells
 Also prevents intermixing of solutions between 2 half cells

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ELECTROCHEMISTRY
Representation of Galvanic cell
 The electrode showing oxidation reaction is
anode and the other electrode where reduction
occurs is cathode
 As per IUPAC convention, the anode is always
represented on the left and cathode always
represented on the right side of the cell
 Anode Half-cell // Cathode Half-cell
 The electrode on left is written by writing the metal first and then the
electrolyte
 The two are separated by a vertical line or a semicolon
 Electrode/Anode solution // cathode solution /
electrode
Zn (s) / Zn2+ (1M) // Cu2+ (1M) / Cu (s)

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ELECTROCHEMISTRY
Electrolytic cell (molten NaCl)
 Conversion of electrical energy to chemical
energy
 A cell that requires electrical energy to cause
nonspontaneous redox reactions to occur
 An electrolytic cell is a system of 2 inert
electrodes (graphite or Pt) and an electrolyte
connected to a power supply like battery
 The electrode which is connected to negative
terminal of the battery is cathode and the
electrode connected to positive terminal of
battery is anode
Construction of
electrolytic cell

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ELECTROCHEMISTRY
Working: when electricity is passed to the cell, Na+ ions
migrate towards the cathode and Cl- ions towards the
Chemical reactions
At cathode: 2Na+ + 2e- 2Na (reduction)
At anode: 2Cl- Cl2 + 2e- (oxidation)
overall reaction 2 Na+
(l) + 2Cl-
(l) 2Na(l) + Cl2(g)
Net reaction of molten NaCl
2NaCl(l) 2Na(l) + Cl2(g)

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ELECTROCHEMISTRY
Standard Electrode Potential
Standard electrode potential – electropositive character
The absolute value of electrode potential cannot be measured/calculated
For this purpose, Normal Hydrogen Electrode (NHE) or Standard
Hydrogen Electrode (SHE) is used
Electromotive force (EMF): potential difference between 2 electrodes of a
galvanic cell that causes flow of current from an electrode with higher
reduction potential to electrode with lower reduction potential. It is
denoted as Ecell Ecell = Eright- Eleft or E cathode – E anode
Under standard conditions, the standard electrode potential occurs in an electrochemical
cell say the temperature = 298K, pressure = 1atm, concentration = 1M. The symbol ‘Eo
cell’
represents the standard electrode potential of a cell

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ELECTROCHEMISTRY
Significance of Standard Electrode Potential
 The electric potential that arises between the anode and cathode is due to the difference in
the individual potentials of each electrode
 It is measured with the help of a reference electrode known as the standard hydrogen
electrode (SHE).
 The potential of SHE is considered as 0 Volts
 Good oxidizing agents have high standard reduction potentials
whereas good reducing agents have low standard reduction potentials.
 For example, the standard electrode potential of Ca2+ is -2.87 V. and
that of F2 is +2.87 V. This implies that F2 is a good oxidizing agent
whereas Ca is a reducing agent

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ELECTROCHEMISTRY
Determination of Standard Electrode Potential
The calculation of the standard electrode potential of a zinc electrode with the help of the standard hydrogen electrode is
illustrated below. It can be noted that this potential is measured under standard
conditions where the temperature is 298K, the pressure is 1 atm, and
the concentration of the electrolytes is 1M

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ELECTROCHEMISTRY
Spontaneity of Redox Reactions
 If a redox reaction is spontaneous, the ΔGo (Gibbs free energy) must have a negative value.
 It is described by the following equation:
ΔGocell = -nFE0cell
 Where n refers to the total number of moles of electrons for every mole of product formed, F
is Faraday’s constant (approximately 96485 C.mol-1)
 The E0cell can be obtained with the help of the following equation:
E0cell = E0cathode – E0anode
 Therefore, the E0cell can be obtained by subtracting the standard
electrode potential of the anode from that of the cathode.

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ELECTROCHEMISTRY
Spontaneity of Redox Reactions
 For a redox reaction to be spontaneous, the E0cell must have a positive value (because both
n and F have positive positive values, and the ΔGo value must be negative).
 This implies that in a spontaneous process,
 E0cell > 0; which in turn implies that E0cathode > E0anode
 Thus, the standard electrode potential of the cathode and the anode
help in predicting the spontaneity of the cell reaction.
 It can be noted that the ΔGo of the cell is negative in galvanic cells
and positive in electrolytic cells.