090643 Electrochemistry presentation.pdf

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Electrochemistry is the study of
production of electricity from energy
released during spontaneous chemical
reactions
and
the use of electrical energy to bring
about non-spontaneous chemical
transformations.

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Electrochemical Cells /Galvanic
cells /voltaic cell:
It is a cell in which electricity is
produced by a spontaneous reaction.
E.g. Daniell cell
Electrolytic cell:
It is a device used to carry out non-
spontaneous chemical reactions using
electricity.

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Galvanic Cells-Daniell Cell

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Galvanic Cells-Daniell Cell
spontaneous
redox reaction
anode
oxidation
cathode
reduction

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Zn (s) + Cu2+
(aq) → Zn 2+
(aq) + Cu (s)
cell reaction
Zn (s) → Zn 2+
+ 2e–
(oxidation half reaction)
Cu 2+
+ 2e–
→ Cu (s) (reduction half reaction)
Daniell Cell

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In a galvanic cell, the half-cell in which
oxidation takes place is called anode
and it has a negative potential with
respect to the solution.
The other half-cell in which reduction
takes place is called cathode
and it has a positive potential with
respect to the solution.

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SALT BRIDGE
Functions
1. It allows the flow of current by
completing the circuit
2. It maintains the electrical neutrality
Salt bridge is a U -shaped tube containing a
semi solid paste of some inert electrolyte
like KCl, KNO3
,NH4
Cl .... in agar agar or
gelatin

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The tendency of a metal to lose or gain
electron when it is in contact with its own
solution is called electrode potential
A potential difference develops
between the electrode and the
electrolyte which is called
electrode potential.
Electrode Potential

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Electrode Potential
When the concentrations of all the
species involved in a half-cell is
unity then the electrode potential is
known as standard electrode
potential.
According to IUPAC convention,
standard reduction potentials are
now called standard electrode
potentials.

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Cell Potential
The potential difference between the two
electrodes of a galvanic cell is called the
cell potential and is measured in volts.
The cell potential is the difference
between the electrode potentials
(reduction potentials) of the cathode and
anode.
It is called the cell electromotive force
(emf)of the cell when no current is
drawn through the cell

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The emf of the cell is positive and is
given by the potential of the half-cell
on the right hand side minus the
potential of the half-cell on the left
hand side
i.e.
Ecell = Eright – E left
Ecell = Ecathode
– Eanode
Cell Potential

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Representing a cell
the anode is kept on the left and the
cathode on the right
In anode metal is represented first
followed by metal ion
In cathode metal ion is represented
first followed by metal
put a vertical line between metal and
metal ion
a salt bridge is denoted by putting a
double vertical line

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Daniell Cell
Zn (s) + Cu 2+
(aq) → Zn 2+
(aq) + Cu (s)
Zn (s) → Zn
2+
+ 2e
–
(oxidation half reaction)
Cu
2+
+ 2e
–
→ Cu (s) (reduction half reaction)
The cell can be represented as:
Zn(s)|Zn
2+
(aq) || Cu
2+
(aq)| Cu(s)
Representing a cell

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Cell reaction:
Cu(s) + 2Ag+
(aq) → Cu2+
(aq) + 2 Ag(s)
Half-cell reactions:
Cathode (reduction): 2Ag+
(aq) + 2e–
→ 2Ag(s)
Anode (oxidation): Cu(s) → Cu2+
(aq) + 2e–
The cell can be represented as:
Cu(s)|Cu2+
(aq)||Ag+
(aq)|Ag(s)
Ecell = Eright – Eleft = EAg
+
⎥Ag – ECu
2+
⎥Cu

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Effect of opposing potential on the cell reaction

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Electrochemistry
PART 2

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Measurement of Electrode Potential
Measurement of Electrode Potential
The potential of individual half-cell cannot be
measured.
We can measure only the difference between the two
half-cell potentials that gives the emf of the cell.
For this purpose a half-cell called
Standard HydrogenElectrode (SHE) or
Normal Hydrogen Electrode (NHE) is used.

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Standard HydrogenElectrode (SHE)
It consists of a platinum
electrode coated with
platinum black.
The electrode is dipped in an
acidic solution of one molar
concentration
and
pure hydrogen gas at 1 bar
pressure and 298K is
bubbled through it.
By convention, the electrode potential of SHE is taken as zero.

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Standard HydrogenElectrode (SHE)
If SHE act as cathode
Reduction Takes place as
2e-
+ 2H+
(1 M) H2 (1 atm)
The half cell can be represented as
If SHE act as anode
Oxidation takes place as
H2 (1 atm) 2H+
(1 M) + 2e-
The half cell can be represented as
Pt (s) | H2 (1 atm) | H+
(1 M)
H+
(1 M) | H2 (1 atm) | Pt (s)

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Standard Electrode Potentials of Zn
19.3
Zn (s) | Zn2+
(1 M) || H+
(1 M) | H2 (1 atm) | Pt (s)
2e-
+ 2H+
(1 M) H2 (1 atm)
Zn (s) Zn2+
(1 M) + 2e-
Anode (oxidation):
Cathode (reduction):
Zn (s) + 2H+
(1 M) Zn2+
+ H2 (1 atm)

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19.3
E0
= 0.76 V
cell
Standard emf (E0
)
cell
0.76 V = 0 - EZn /Zn
0 2+
EZn /Zn = -0.76 V
0 2+
Zn2+
(1 M) + 2e-
Zn E0
= -0.76 V
E0
= EH /H - EZn /Zn
cell
0 0
+ 2+
2
Standard Electrode Potentials of Zn
E0
= Ecathode - Eanode
cell
0 0
Zn (s) | Zn2+
(1 M) || H+
(1 M) | H2 (1 atm) | Pt (s)

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Standard Electrode Potentials of Cu
19.3
Pt (s) | H2 (1 atm) | H+
(1 M) || Cu2+
(1 M) | Cu (s)
2e-
+ Cu2+
(1 M) Cu (s)
H2 (1 atm) 2H+
(1 M) + 2e-
Anode (oxidation):
Cathode (reduction):
H2 (1 atm) + Cu2+
(1 M) Cu (s) + 2H+
(1 M)
E0
= Ecathode - Eanode
cell
0 0
E0
= 0.34 V
cell
Ecell = ECu /Cu – EH /H
2+ +
2
0 0 0
0.34 = ECu /Cu - 0
0 2+
ECu /Cu = 0.34 V
2+
0

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Electrochemical series
It is a series in which various electrodes are
arranged in the increasing order of their
reduction potential

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• The more positive E0
the
greater the tendency for
the substance to be
reduced
The half-cell reactions are
reversible
• The sign of E0
changes
when the reaction is
reversed
• Changing the
stoichiometric
coefficients of a half-cell
reaction does not change
the value of E0

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Applications of Electrochemical series
Substances with higher reduction potentials are strong
oxidising agents
Fluorine has the highest electrode potential thus
Fluorine gas is the strongest oxidising agent
Substances with lower reduction potentials are strong
reducing agents
Lithium has the lowest electrode potential thus
lithium metal is the most powerful reducing agent
in an aqueous solution
1.Comparing the relative oxidising and
reducing powers

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2.Calculation of the standard EMF of the Cell
Eo
cell
= Eo
cathode
- Eo
anode
For Daniell cell
Zn (s)|Zn
2+
(aq) || Cu
2+
(aq)| Cu(s)
Eo
cell
= Eo
cathode
- Eo
anode
= 0.34 - 0.76 = 1.1v

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Fig 18-14 Pg 877
The copper/zinc electrochemical cell. The voltmeter measures the
difference in electrical potentials between the two electrodes.

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3. Predicting the Feasibility of Redox
Reaction
When Eo
is positive, the cell reaction is
spontaneous
Zn (s) + Cu2+
(aq) → Zn 2+
(aq) + Cu (s)
Eo
cell
=1.1 v Thus reaction is feasible
Zn 2+
(aq) + Cu (s)→ Zn (s) + Cu2+ (aq)
Eo
cell
= -1.1 v Thus reaction is not feasible

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4. Predicting the capability of a metal to
displace hydrogen from dilute acids
Metals having negative reduction potential
can displace hydrogen from dilute acids
Eg ;- Zn , Mg..

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NERNST EQUATION
It relate the electrode potential of an electrode (or,emf
of a cell ) with the concentration of electrolytic
Mn+
(aq) + ne-
M(s)
R is gas constant (8.314JK–1
mol–1
)
F is Faraday constant (96487 C mol–1
)
n is number of electrons
T is temperature in Kelvin and
[Mn+
] is the concentration of the species,Mn+
.

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NERNST EQUATION
E(M
n+
/M) = E°(M
n+
/M) - log (1 / [Mn+
])
concentration of solid M is taken as unity
2.303 RT
nF

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For Daniell cell
E(Cu
2+
/Cu)= E°(Cu
2+
/Cu) - (2.303RT/2F) log (1/[Cu2+
])
E(Zn
2+
/Zn) = E°(Zn
2+
/Zn) - (2.303RT/2F) log (1/[Zn2+
])
Ecell= E(Cu
2+
/Cu)- E(Zn
2+
/Zn)
= E°(Cu
2+
/Cu) - (2.303RT/2F) log (1/[Cu2+
]) -
{E°(Zn
2+
/Zn) - (2.303RT/2F) log (1/[Zn2+
])}

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For Daniell cell
= E°(Cu
2+
/Cu)- E°(Zn
2+
/Zn) - (2.303RT/2F) log
([Zn2+
]/[Cu2+
])
= E0
cell - (2.303RT/2F) log ([Zn2+
]/[Cu2+
])
Ecell = E0
cell – 0.059/2 log ([Zn2+
]/[Cu2+
])
R= 8.314JmolL-1
K-1
, T=298, F= 96487Cmol-1
Ecell increases with increase in concentration of
Cu 2+
and decreases with increase in
concentration of Zn2+

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Ni(s)|Ni2+
(aq)||Ag+
(aq)|Ag(s)
Ni(s) + 2Ag+
(aq) → Ni2+
(aq) + 2 Ag(s)
Nernst equation can be written as

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Equilibrium Constant from Nernst Equation
Daniell Cell
Zn (s) + Cu 2+
(aq) → Zn 2+
(aq) + Cu (s)
Ecell = E0
cell – 0.059/2 log ([Zn2+
]/[Cu2+
])
Ecell = 0 = E0
cell – 0.059/2 log ([Zn2+
]/[Cu2+
])

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Electrochemical Cell and Gibbs Energy of the
Reaction
The relationship between Gibbs energy and cell potential is
If the concentration of all the reacting species is unity

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Electrochemistry
PART 3

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Conductance of Electrolytic Solutions
Electrical resistance (
Electrical resistance (R
R)
) of any object is directly proportional
of any object is directly proportional
to its length (
to its length (l
l) and inversely proportional to its cross-
) and inversely proportional to its cross-
sectional area (
sectional area (A
A), i.e.,
), i.e.,
Or,
Or,
Where,
Where, ρ
ρ = resistivity or specific resistance
= resistivity or specific resistance
Its SI unit is ohm metre (Ω m) and is often expressed in
Its SI unit is ohm metre (Ω m) and is often expressed in
ohm centimetre.
ohm centimetre.
Resistivity − Resistance of a substance when it is 1
Resistivity − Resistance of a substance when it is 1
metre long and has 1 m
metre long and has 1 m2
2 cross-sectional area.
cross-sectional area.

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Conductance (G):
It is the inverse of resistance
ƙ is called conductivity. It is defined as the conductance of a
conductor having unit length and unit area of cross-section.
Its unit is ohm
Its unit is ohm -1
-1
m
m -1
-1
or mho m
or mho m -1
-1
or S m
or S m -1
-1
.
.
1 S cm
1 S cm -1
-1
= 100 S m
= 100 S m -1
-1
1 S m
1 S m -1
-1
= 10
= 10 -2
-2
S cm
S cm -1
-1

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Two types
Electronic or metallic conductance
Electrolytic or ionic conductance.
Types of Conductance
Types of Conductance

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Electrical conductance through metals is called
metallic or electronic conductance and it is due to
the movement of electrons.
Electronic or metallic conductance

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The electronic conductance depends on
(i) the nature and structure of the metal
(ii) the number of valence electrons per atom
(iii) temperature (it decreases with increase of
temperature).

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The conductance of electricity by ions present
in the solutions is called electrolytic or ionic
conductance
Ionic conductance

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The conductivity of electrolytic (ionic)
solutions depends on:
(i) the nature of the electrolyte added
(ii) size of the ions produced and their
solvation
(iii) the nature of the solvent and its viscosity
(iv) concentration of the electrolyte
(v) temperature (it increases with the increase
of temperature)

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Measurement of the Conductivity of Ionic Solutions
Conductivity cell
It consists of two platinum electrodes coated with
platinum black. The electrodes are separated by a
distance Ɩ and their area of cross-section is A.

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R = ρ.l/A ρ is resistivity
= l/ƙ A ƙ is conductivity
The quantity l/A is called
cell constant denoted by the symbol, G*.
R=G*/ ƙ
G* = R. ƙ
ƙ= G*/R

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Electrochemistry
PART 4

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It is the conductivity of 1 mole of an electrolytic
solution kept between two electrodes with unit
area of cross section and at a distance of unit
length
Molar conductivity (Λm)
1 S m2
mol –1
= 104
S cm2
mol–1
or
1 S cm 2
mol–1
= 10 –4
S m2
mol–1
The unit of Λm Ω -1
cm2
mol-1
or S cm 2
mol -1

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Conductivity always decreases with
decrease in concentration both, for
weak and strong electrolytes.
This is due to the fact that the number of
ions per unit volume that carry the current
in a solution decreases on dilution.
Variation of Conductivity and Molar
Conductivity with Concentration

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Molar conductivity of a solution at a
given concentration is the conductance of
the volume V of solution containing one
mole of electrolyte

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Molar conductivity increases with
decrease in concentration.
This is because the total volume, V, of
solution containing one mole of electrolyte
also increases. It has been found that
decrease in κ on dilution of a solution is
more than compensated by increase in its
volume.

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Limiting Molar Conductivity Λm
0
When concentration of solution
approaches zero, molar
conductivity is called limiting
molar conductivity

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For a strong
electrolyte, Λm
0
can be
determined by
extrapolating
the graph,as
Λmvaries
linearly.
Determination of Λm
0

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Strong Electrolytes
The value of the constant ‘A’ for a given solvent and
temperature depends on the type of electrolyte
i.e., the charges on thecation and anion produced on the
dissociation of the electrolyte in thesolution.
Thus, NaCl, CaCl 2
, MgSO 4
are known as 1-1, 2-1 and
2-2electrolytes respectively.
All electrolytes of a particular type have thesame value
for ‘A’.

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Kohlrausch law of independent
migration of ions
The law states that limiting molar
conductivity of an electrolyte can be
represented as the sum of the
individual contributions of the anion
and cation of the electrolyte.

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Λm
0
(NaCl) = λ0
Na
+
+ λ0
Cl
–
λ°Na+ and λ°Cl
–
are limiting molar conductivity
of the sodium and chloride ions respectively
In general, if an electrolyte on dissociation
gives v+ cations and v– anions then its limiting
molar conductivity is given by:
Λm
0
= ν+ λ0
+ + ν– λ0
–
Here,λ +
0
andλ−
0
are the limiting molar
conductivities of the cation and anion

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Applications of Kohlrausch law
1. Determination of Λm
0
of weak electrolytes
Λ o
m(CH3COOH)=
Λ o
m (CH3COONa)+Λ o
m(HCl)- Λ o
m(NaCl)

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Applications of Kohlrausch law
2. Determination of degree of dissociation of weak
electrolytes

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Electrochemistry
PART 5

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Electrolytic Cells and Electrolysis
In an electrolytic cell, the
electrical energy is converted to
chemical energy.
The dissociation of an
electrolyte by the passage of
electricity is called electrolysis.

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Faraday’s Laws of Electrolysis
First Law:
The amount of chemical reaction
which occurs at any electrode during
electrolysis by a current is
proportional to the quantity of
electricity passed through the
electrolyte (solution or melt).

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m α Q Or, m = zQ
Where z is a constant called
electrochemical equivalent (ECE).
But Q = It
I=current in ampere t= time in second
m= zIt
When I = 1 amp, t = 1 sec then Q = 1 coulomb, then
m = Z.
Thus, electrochemical equivalent is the
amount of the substance deposited or
liberated by passing 1A current for 1 sec
(i.e.. 1 coulomb)

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Ag +
(aq) + e–
→ Ag(s)
One mole of the electron is required for
the reduction of one mole of silver ions.
Charge on one electron = 1.6021×10–19
C.
Charge on one mole of electrons =
6.02 × 1023
mol–1
× 1.6021 × 10–19
C
= 96487 C mol–1
This quantity of electricity is called Faraday
and is represented by the symbol F.
Z = equivalent weight/96500

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Second Law:
The amounts of different substances
liberated by the same quantity of
electricity passing through the
electrolytic solution are proportional to
their chemical equivalent weights
(Atomic Mass of Metal ÷ Number of
electrons required to reduce the
cation).`

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Second Law:

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Products of Electrolysis
1. Molten NaCl
The products of electrolysis are
sodium metal and Cl2
gas.
Na+
is reduced at the cathode
Na+
+ e–
→ Na
Cl–
is oxidised at the anode
Cl–
→ ½Cl2+e–

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The products are NaOH, Cl2
and H2
.
In this case besides Na+
and Cl–
ions
we also have H+
and OH–
ions along
with the solvent molecules, H2
O.
At the cathode there is competition
between the following reduction
reactions:
Na+
(aq) + e–
→ Na(s) E0
cell = – 2.71 V
H+
(aq) + e–
→ ½ H2(g) E0
cell = 0.00 V
2. Aqueous NaCl Solution

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The reaction with higher value of E0
is
preferred and, therefore, the reaction at
the cathode during electrolysis is:
H+
(aq) + e–
→ ½ H2 (g)
but H+
(aq) is produced by the
dissociation of H2
O, i.e.,
H2
O (l ) → H+
(aq) + OH–
(aq)
Therefore, the net reaction at the cathode
may be written as the sum of the above
two and we have
H2
O (l ) + e–
→ ½H2(g) + OH–

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At the anode the following oxidation
reactions are possible:
Cl–
(aq) → ½ Cl2 (g) + e–
E0
cell =1.36 V
2H2
O (l )→ O2 (g) + 4H+
(aq)+ 4e– E0
cell =1.23 V
The reaction at anode with lower value of
E0
is preferred and therefore, water should
get oxidised in preference to Cl–
(aq).
However, on account of overpotential of
oxygen, the first reaction is preferred

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Net reaction:
NaCl(aq) + H2
O(l) →
Na+
(aq) + OH–
(aq) + ½H2(g) + ½Cl2(g)

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3. During the electrolysis of sulphuric acid, the following processes
are possible at the anode:
For dilute sulphuric acid, first reaction is preferred
but at higher concentrations of H2
SO4
, second
reaction is preferred

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Electrochemistry
PART 6

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Batteries
A battery is basically a galvanic cell in
which the chemical energy of a redox reaction is
converted to electrical energy.
Two types –
primary batteries and
secondary batteries.

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These are cells which cannot be recharged or
reused.
Here the reaction occurs only once and after use
over a period of time, they become dead
E.g. Dry cell, mercury button cell etc
Primary cells:

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1. Dry Cell
It is compact
form of
Laclanche cell

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Dry cell Consists of
* Zinc container as anode.
* Graphite rod surrounded by MnO2 and
carbon as cathode.
* A paste of NH4Cl and ZnCl2 as
electrolyte.
Anode- Zn(S) —› Zn2+
+ 2e-
Cathode - 2MnO2(s) +NH4
+
+ e-
—›
MnO(OH) + NH3
The emf of the cell is 1.5V

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Anode – Zinc amalgam.
Cathode – A paste of HgO and carbon.
Electrolyte - A paste of KOH and ZnO
Anode – Zn(Hg) + 2OH-
—› ZnO + H2O + 2e-
Cathode – HgO + H2O + 2e-
—› Hg + 2OH-
Overall reaction – Zn + HgO —› ZnO + Hg
The cell has a constant potential of 1.35 V,
since the overall reaction does not
involve any ion in solution

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Secondary Batteries
A secondary cell after use can be recharged by
passing current through it in the opposite
direction.
Thus it can be reused again and again
Eg;- Lead storage battery

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Lead Storage Batteries:

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Lead storage battery
Lead plates act as anode.
Grid of lead plates coated with lead dioxide
(PbO2
)act as cathode
A solution of H2SO4 (38% by mass)acts as
electrolyte.
Anode - Pb(s) + SO4
2-
(aq)
—› PbSO4 (s)+ 2e-
Cathode –PbO2(s) + 4H+
(aq)
+ SO4
2-
(aq)
+ 2e-
—›
PbSO4(s) +2H2O(l)
Cell reaction- Pb(s) + PbO2(s) + 2H2SO4(aq)
—›
2PbSO4 (s)+2 H2O (l)
On recharging cell reaction is reversed.

Nickel-Cadmium Battery
Nickel-Cadmium Battery

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Nickel – Cadmium cell
Cadmium metal act as anode.
Metal grid containing Nickel(IV)oxide act as
cathode.
KOH solution act as electrolyte.
Cell reaction –
Cd(S) + 2Ni(OH)3(s) —›CdO(S) + 2Ni(OH)2(s)+ H2O
Cell reaction is reversed on recharging.
emf is approximately 1.4 V .
Also called Nicad cell.

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FUEL CELLS
These are galvanic cells which convert
the energy of combustion of fuels like
hydrogen, methane,methanol, etc.
directly into electrical energy.
Eg:- H2 – O2 fuel cell.

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Hydrogen – Oxygen fuel cell
Here hydrogen and oxygen are bubbled through porous carbon
electrodes into concentrated aqueous sodium hydroxide solution.
To increase the rate of electrode reactions, catalysts like finely
divided platinum or palladium metal are filled into the
electrodes.
Iis used in the
Apollo space
programme.

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Anode – 2[H2(g) + 2OH-
(aq) —› 2 H2O(l) + 2e-
]
Cathode – O2 (g)+ 2 H2O(l) + 4e-
—› 4OH-
(aq)
Cell reaction -2 H2(g) + O2 —› 2 H2O(aq)
Cell works as long as H2 and O2 are supplied
Advantages
1) Efficient than any conventional source.
2) Pollution free working
3)The cell runs continuously as long as the
reactants are supplied

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It is the slow destruction of a metal due to attack of
atmospheric gases and moisture on their surface
resulting in the formation of compounds such as
oxide ,sulphide ,carbonate etc..
More reactive metals corrode more easily.
Rusting of iron.
Tarnishing of silver,
Formation of green coating on copper (verdigris)
Corrosion

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Mechnism of Rusting
(Electrochemical theory of rusting)
At a particular spot of the metal, oxidation takes place and that
spot behaves as anode. Here Fe is oxidized to Fe 2+
2Fe(s) —› 2Fe2+
+ 4e-
(Anode)
Electrons released at anodic spot move through the
metal and go to another spot on the metal and reduce
oxygen in presence of H + . This spot behaves as
cathode.
[H+ available from H2
CO3
or from H2
O]

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2Fe(s) —› 2Fe2+
(aq)+ 4e-
(Anode)
O2 (g)+ 4H+
(aq) + 4e-
—› 2 H2O(l) (Cathode)
2Fe(s) + O2 (g)+ 4H+
(aq) —› 2Fe2+
(aq)+ 2 H2O(l)
The ferrous ions(Fe2+
) are further oxidised by
atmospheric oxygen to ferric ions(Fe3+
) which come
out as rust in the form of hydrated ferric oxide,
Fe2O3 .x H2O and with further production of
hydrogen ions

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Prevention of corrosion
Barrier protection –
By coating with a suitable material –paint,grease
Sacrificial protection –
Coating with a more reactive metal.
The process of coating the surface of iron with
Zinc is called Galvanization.

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Cathodic protection - Here metal to be protected is set
as cathode by attaching a more reactive metal to it.
More reactive metal act as anode.
Now the more reactive metal undergo oxidation.
Antirust solutions – Alkaline phosphate or chromate
solutions are applied on iron surface to form a iron
phosphate or chromate coating which prevent corrosion.

YS Classes
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090643 Electrochemistry presentation.pdf

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