1

ELECTRO CHEMISTRY
ELECTROCHEMICAL CELLS:
A device for producing an electrical current from a chemical reaction (spontaneous redox
reaction) is called an electrochemical cell and also known as a galvanic cell or Voltaic cell. A simple
voltaic cell is shown in Fig.1. Here the spontaneous reaction of zinc metal with an aqueous solution of
copper sulphate is used.
Zn(s) + Cu2+ Zn2+ + Cu
A bar of zinc metal (anode) is placed in zinc sulphate solution (where oxidation takes place) in
the left container. A bar of copper metal
(cathode) is immersed in copper sulphate
solution (where reduction takes place) in
the right container. In other words, each
electrode may be regarded as a half-cell.
The zinc and copper electrodes are joined
by a copper wire. A salt bridge containing
potassium sulphate solution interconnects
the solutions in the anode compartment
and the cathode compartment.
The oxidation half-reaction occurs in the
anode compartment. Fig. 1 A simple voltaic (galvanic) cell.
Zn(s) Zn (aq) + 2e
2+ –

The reduction half-reaction takes place in the cathode compartment.

Cu2+(aq) + 2e– Cu(s)
The cell reaction: Zn(s) + Cu (aq)
+2
Zn+2(aq) + Cu(s)
When the cell is set up, electrons flow from zinc electrode through the wire to the copper
cathode. As a result, zinc dissolves in the anode solution to form Zn2+ ions. The Cu2+ ions in the cathode
half-cell pick up electrons and are converted to Cu atoms on the cathode. At the same time, SO42– ions
from the cathode half-cell migrate to the anode half-cell through the salt bridge. Likewise, Zn 2+ ions
from the anode half-cell move into the cathode half-cell. This flow of ions from one half-cell to the other
completes the electrical circuit which ensures continuous supply of current. The cell will operate till either
the zinc metal or copper ion is completely used up.
CELL TERMINOLOGY:
Before taking up the study of the electrochemical cells, we should be familiar with a few common
terms.
Current is the flow of electrons through a wire or any conductor.
Electrode: a metallic rod/bar/strip which conducts electrons into and out of a solution.
Anode is the electrode at which oxidation occurs. It sends electrons into the outer circuit. It has negative
charge and is shown as (–) in cell diagrams.
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Cathode is the electrode at which electrons are received (reduction occurs) from the outer circuit. It
has a positive charge and is shown as (+) in cell diagrams.
Electrolyte is the salt solutions in a cell.
Anode compartment is the compartment of the cell in which oxidation half-reaction occurs. It contains
the anode.
Cathode compartment is the compartment of the cell in which reduction half-reaction occurs. It
contains the cathode.
Half-cell. Each half of an electrochemical cell, where oxidation occurs and the half where reduction
occurs, is called the half cell.
Salt bridge: A salt bridge is a U-shaped device containing concentrated solution of an inert electrolyte
like KCl, KNO3, etc. or a solidified solution of those electrolytes in agar-agar solution and gelatin to
prevent intermixing of the solutions. It connects the oxidation and reduction half-cells of a galvanic cell.
The inert electrolytes present do not take part in redox reaction of the cell and don’t react with the
electrolyte that has been used.
Cell diagram or Representation of a Cell
A cell diagram is an abbreviated symbolic depiction of an electrochemical cell. For this purpose,
we will consider that a cell consists of two half-cells. Each half-cell is again made of a metal electrode
contact with metal ions in solution.
IUPAC Conventions: In 1953 IUPAC recommended the following conventions for writing cell diagrams.
We will illustrate these with reference to Zinc-Copper cell.
(1) A single vertical line (|) represents a phase boundary between metal electrode and ion
solution (electrolyte). Thus, the two half-cells in a voltaic cell are indicated as

Phase boundary

Zn | Zn2+ Cu2+ | Cu
Anode Half-Cell Cathode Half-Cell
It may be noted that the metal electrode in anode half-cell is on the left, while in cathode half- cell it
is on the right of the metal ion.
(2) A double vertical line (||) represents the salt bridge, porous partition or any other means of
permitting ion flow while preventing the electrolyte from mixing.
(3) Anode half-cell is written on the left and cathode half-cell on the right.
(4) In the complete cell diagram, the two half-cells are separated by a double vertical line
(salt bridge) in between. The zinc-copper cell can now be written as
Salt Bridge
Zn | Zn2+ || Cu2+ | Cu
Anode Cathode
Half-cell Half-cell
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(5) The symbol for an inert electrode, like the platinum electrode is often enclosed in a bracket. For
example,

The value of emf of a cell is written on the right of the cell diagram. Thus, a zinc-copper cell has emf
1.1V and is represented as

Electromotive force (emf) or Cell potential: In a Zn-Cu voltaic cell, electrons are released at the
anode and it becomes negatively charged. The negative electrode pushes electrons through the external
circuit by electrical repulsions. The copper electrode gets positive charge due to the discharge of Cu2+
ions on it. Thus, electrons from the outer circuit are attracted into this electrode. The flow of current
through the circuit is determined by the ‘push’, of electrons at the anode and ‘attraction’ of electrons at
the cathode. These two forces constitute the ‘driving force’ or ‘electrical pressure’ that sends electrons
through the circuit. This driving force is called the electromotive force (abbreviated emf) or cell
potential. The emf of cell potential is measured in units of volts (V) and is also referred to as cell
voltage.
The magnitude of the emf of a cell reflects the tendency of electrons to flow externally from one
electrode to another. The electrons are transported through the cell solution by ions present and pass
from the positive electrode (Cu in case of Daniel cell) to the negative electrode. This corresponds to a
clockwise flow of electrons through the external circuit. Thus, the emf of the cell is given the +ve sign.
If the emf acts in the opposite direction through the cell circuit, it is quoted as –ve value. For example,
Daniel cell has an emf of 1.1V and the copper electrode is positive. This can be expressed in two ways:
Zn | ZnSO4 || CuSO4 | Cu E = + 1.1 V
Cu | CuSO4 || ZnSO 4 | Zn E = – 1.1 V
The negative sign indicates that the cell is not feasible in the given direction. The reaction will
take place in the reverse direction.
Calculating the emf of a cell:
The electromotive force may be defined as the potential difference between two
electrodes of a galvanic cell or voltaic cell. Or
The difference of potential, which causes the current to flow from an electrode at higher
potential to the one of lower potential, is called the Electro motive force (emf) of the cell.
Mathematically, emf of an electrochemical cell is the algebraic sum of the single electrode
potential; provided proper signs are being given according to the actual reaction taking place on the
electrodes.
The emf of a cell can be calculated from the half-cell potentials of the two cells (anode and
cathode) by using the following formula
Ecell = Ecathode – Eanode
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= ER – EL (IUPAC convention 3)
Where ER and EL are the reduction potentials of the right-hand and left-hand electrodes respectively.
Standard emf of a cell:
The emf generated by an electrochemical cell is given by the symbol E. It can be measured with
the help of a potentiometer. The value of emf varies with the concentration of the reactants and products
in the cell solutions and the temperature of the cell. When the emf of a cell is determined under standard
conditions, it is called the standard emf. The standard conditions are (a) 1 M solutions of reactants
and products; and (b) temperature of 25°C. Thus, standard emf may be defined as: the emf of a cell
with 1 M solutions of reactants and products in solution measured at 25°C. Standard emf of a
cell is represented by the symbol E°. With gases 1 atm pressure is a standard condition instead of
concentration. For a simple Zn-Cu voltaic cell, the standard emf, E°, is 1.10 V. This means that the emf
of the cell operated with [Cu2+] and [Zn2+] both at 1 M and at 25°C is 1.10 V. That is,
Zn | Zn2+(aq, 1M) || Cu2+(aq,1M ) | Cu E = + 1.1 V
Single electrode potential:
An electrochemical cell consists of two half-cells. With an open-circuit, the metal electrode in
each half-cell transfers its ions into solution. Thus, an individual electrode develops a potential with
respect to the solution. The potential of a single electrode in a half-cell is called the Single electrode
potential. Thus, in a Daniel cell in which the electrodes are not connected externally, the anode Zn/Zn2+
develops a negative charge and the cathode Cu/Cu2+, a positive charge. The amount of the charge
produced on individual electrode determines its single electrode potential. The single electrode potential
of a half-cell depends on: (a) concentration of ions in solution; (b) tendency to form ions; and (c)
temperature.
The tendency of an electrode to lose or gain electrons, when it is in contact with the solution of
its own ions. The metal which has a greater tendency to lose electrons becomes the anode, while which
has a greater tendency to gain electrons will behave as cathode.
It may be noted that absolute values of these electrode potentials cannot be determined directly.
These are found by connecting the half-cell with a standard hydrogen electrode whose reduction
potential has been arbitrarily fixed as zero.
Determination of electrode potential or emf of a half-cell:
By a single electrode potential, we also mean the emf of an isolated half-cell or its half- reaction.
The emf of a cell that is made of two half-cells can be determined by connecting them to a voltmeter.
However, there is no way of measuring the emf of a single half-cell directly. A convenient procedure to
do so is to combine the given half-cell with another standard half-cell. The emf of the newly constructed
cell, E, is determined with a voltmeter. The emf of the unknown half-cell, E°, can then be calculated
from the expression
Emeasured = ER – EL
If the standard half-cell acts as anode, the equation becomes.
ER = Emeasured (∵ EL = 0)
On the other hand, if standard half-cell is cathode, the equation takes the form
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EL = – Emeasured (∵ ER = 0)
The standard hydrogen half-cell or Standard
Hydrogen Electrode (SHE) is selected for coupling
with the unknown half-cell. It consists of a platinum
electrode immersed in a 1 M solution of H+ ions
maintained at 25°C. Hydrogen gas at one atmosphere
enters the glass hood and bubbles over the platinum
electrode.
The hydrogen gas at the platinum electrode passes into
solution, forming H+ ions and electrons.
The standard hydrogen electrode:
The emf of the standard hydrogen electrode is
Fig. 3 Standard hydrogen electrode
arbitrarily assigned the value of zero volts. So, SHE
can be used as a standard for other
electrodes. The half-cell whose potential is
desired is combined with the hydrogen
electrode and the emf of the complete cell
determined with a voltmeter. The emf of the
cell is the emf of the half-cell.
For example, it is desired to
determine the electrode potential of the zinc
electrode, Zn | Zn2+. It is connected with the
SHE as shown in Fig. 4. The complete
electrochemical cell may be represented as:
Zn | Zn 2+ || H+ | H2 (1 atm), Pt
Ecell = ER – EL
Fig. 4. The Zinc electrode coupled with SHE
= 0 – 0.76 = – 0.76 V
The emf of the cell has been found to
be –0.76V which is the emf of the zinc half-cell. Similarly, the electrode potential of the copper
electrode, Cu |Cu can be determined by pairing it with the SHE when the electrochemical cell can be
2+

represented as:
Pt, H2 (1 atm) | H+ || Cu2+ | Cu
The emf of this cell has been determined to be 0.34 V which is the electrode potential of the copper
half-cell.
E0cell = E0Cu|Cu2+ - E0SHE
0.34 – 0 = 0.34 V
SHE can act both as cathode and anode when joined with another half-cell (Fig. 5).
When it is placed on the right-hand side of the Zinc electrode, the hydrogen electrode reaction is
2H+ + 2e– H2
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The electrons flow to the SHE

and it acts as the cathode. When
the SHE is placed on the left-
hand side, the electrode reaction
is
H2 2H+ + 2e–
The electrons flow to the copper
electrode and the hydrogen
electrode as the anode.
Evidently, the SHE can act
both as anode and cathode
and, therefore can be used to Fig. 5. The two situations are shown in
determine the emf of any other
half-cell electrode (or single electrode).
In the procedure for determining the emf of a given half-cell, the standard hydrogen electrode
can be placed on the left-hand or the right-hand. The electrons flow from left-to-right and the given
half-cell electrode gains electrons (reduction). The observed emf of the combined electrochemical cell
is then the emf of the half-cell on the right-hand. Such emf values of half-cells, or half reactions, are
known as the Standard reduction potentials or Standard potentials. However, if the SHE be placed
on the right-hand side of the given half-cell, the potential so obtained is called as the Standard
oxidation potential. The latter potentials are the standard potentials with the sign reversed, the only
difference being that cells have been turned around.
According to IUPAC convention, the standard reduction potentials alone are the
standard potentials.
Electrochemical series: The electrochemical series consists of a list of elements have been
arranged in the increasing order of their standard electrode potentials. Or
When elements are arranged in increasing order of their standard electrode potential,
a series called electrochemical series obtained.
Applications of electrochemical series:
1) Predicting the oxidizing or reducing ability 2) Predicting cell emf 3) Predicting feasibility of
reaction 4) Predicting whether a metal will displace another metal from its salt solution or not
5) Predicting whether a metal will displace hydrogen from a dilute acid solution. 5) The relative
corrosion tendencies of the metal.
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Table 1. Standard reduction potentials or Standard potentials

Element Electrode Standard Electrode
Reaction (Reduction) Reduction potential Eo, volt

Strong reductant Li Li+ + e- → Li -3.05

K K+ + e- → K -2.925

Ca Ca2+ + 2e- → Ca -2.87

Na Na+ + e- → Na -2.714
Increasing reducing strength

Mg Mg2+ + 2e- → Mg -2.37

Al Al3+ + 3e- → Al -1.66

Zn Zn2+ + 2e- → Zn -0.7628

Cr Cr3+ + 3e- → Cr -0.74

Fe Fe2+ + 2e- → Fe -0.44

Cd Cd2+ + 2e- → Cd -0.403

Ni Ni2+ + 2e- → Ni -0.25

Increasing oxidizing strength

Sn Sn2+ + 2e- → Sn -0.14

H2 2H+ + 2e- → H2 0.00

Cu Cu 2+
+ 2e → Cu
-
+0.337

I2 I2 + 2e- → 2I- +0.535

Ag Ag+ + e- → Ag +0.799

Hg Hg2+ + 2e- → Hg +0.885

Br2 Br2 + 2e- → 2Br- +1.08

Cl2 Cl2 + 2e- → 2Cl - +1.36

Au Au3+ + 3e- → Au +1.50

F2 F2 + 2e- → 2F- +2.87 (Strongest oxidant)

Predicting the Oxidizing or Reducing Ability:
Let us consider a series of elements Cu, H 2, Ni, Zn and their ions. These four elements could act
as reducing agents. On the other hand, their ions Cu2+, H+, Ni2+ and Zn2+ can act as electron acceptors
or oxidizing agents. If we list the respective half-reactions (or electrodes) in order of increasing E°
values, we will have placed the reducing agents in ascending order of their ability to attract electrons.
It is noteworthy that the value of E° becomes more positive down the series. This means that
Cu 2+
is the best oxidizing agent (most electron-attracting ion) of those in the list. That is, Cu 2+ shows
the greatest tendency to be reduced. Conversely, Zn 2+ is the worst oxidizing agent, being the least
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electron-attracting ion. Of the elements Cu, H2, Ni and Zn, Zn is the best reducing agent (best electron
donor), since E° for the half-reaction has the most positive value. By the reasoning, Cu is the worst
reducing agent.
Zn Zn2+ + 2e– E° = - 0.76 V
The table of half reaction potentials above tells us that at standard conditions the following reactions
occur spontaneously.
Some important points concerning the Table of Standard Reduction Potentials (Table 1) are:
(1) The more positive the value of E°, the better the oxidizing ability (the greater the tendency to be
reduced) of the ion or compound, on moving down ward in the Table.
(2) The more negative the value of E° the better the reducing ability of the ions, elements or compounds
on moving upward in the Table.
(3) Under standard conditions, any substance in this Table will spontaneously oxidize any other
substance higher than it in the Table.
Predicting cell emf: The standard emf, E°, of a cell is the standard reduction potential of right- hand
electrode (cathode) minus the standard reduction potential of the left-hand electrode anode). That is,
E°cell = E°right – E°left
= Cathode potential – Anode potential
Let us predict the emf of the cell
Zn(s) | Zn+2 (aq) || Ag+(aq) | Ag
By using the E° values from the Table
E°cell = E°right – E°left
= 0.80 – (– 0.763) = 0.80 + 0.763 = 1.563 V
The answer is so clear from Fig. 6.

Figure 6. Diagrammatic representation of Cell emf.

Predicting Feasibility of Reaction: The feasibility of a redox reaction can be predicted with the help
of the electrochemical series. The net emf of the reaction, E cell, can be calculated from the expression
E°cell = E°cathode – E°anode
In general, if E°cell = + ve, the reaction is feasible; E° cell = – ve, the reaction is not feasible
SOLVED PROBLEM 1. Predict whether the reaction
2 Ag(s) + Zn2+ (aq) ⎯⎯→ Ag+ (aq) + Zn(s) is feasible or not. Consult the table for the E° values.
SOLUTION: The cell half reactions are
Anode: 2Ag(s) 2Ag+(aq) + 2e– E° = 0.80 V ;
Cathode: Zn2+(aq) + 2e– Zn(s) E° = – 0.763 V
E°cell = E°cathode – E°anode
∴ E°cell = – 0.763 V – 0.80 = – 1.563
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Since E°cell is negative, the given reaction is not feasible.

SOLVED PROBLEM 2. Determine the feasibility of the reaction
2Al(s) + 2Sn4+(aq) 2Al3+ + 3Sn2+(aq) Consult the table for the E° values.
SOLUTION: The given reaction consists of the following half reactions
Anode: 2Al(s) 2Al3+ + 6e– E° = – 1.66 V
Cathode: 3Sn4+ + 6e– 3Sn2+ E° = + 0.15
E°cell = 0.15 – (– 1.66) = 1.81 V Since E° cell is positive, the reaction is feasible.
Predicting whether a metal will displace another metal from its salt solution or not: As already
shown, the metals near the top of the electrochemical series are strong reducing agents and are
themselves oxidized to metal ions. On the contrary, the metals lying bottom in the series are strong
oxidizing agents and their ions are readily reduced to the metal itself. For example, zinc lying down
above the series is oxidized to Zn2+ ion, while copper which is lower in the series is produced by reduction
of Cu2+ ion.
Zn Zn2+ + 2e–
Cu2+ + 2e– Cu↓
Thus, when zinc is placed in CuSO4 solution, Cu metal gets precipitated. In general, we can say
that a metal higher up in the electrochemical series can precipitate the one lower down in the
series. Silver cannot precipitate Cu from CuSO4, solution, since both metals have positions lower in the
series and are strong oxidizing agents.
Predicting whether a metal will displace hydrogen from a dilute acid solution: Any metal below
hydrogen in the electrochemical series is a weaker reducing agent than hydrogen itself and cannot
reduce H+ to H2. Any metal above hydrogen is a stronger reducing agent than hydrogen and will convert
H+ to H2. This explains why Zn lying above hydrogen reacts with dil. H2SO4 to liberate H2, while Cu lying
below hydrogen does not react.
Zn + H+ (dil. H2SO4) Zn2+ + H2 ↑
Cu + H+ (dil. H2SO4) Cu2+ + H2
The relative corrosion tendencies of the metal. The metals above the hydrogen in the series can
easily oxidised hence they undergo corrosion.
THE NERNST EQUATION:
The electrical energy generated by the galvanic cell can be quantitatively converted into work.
So, the emf of the cell is a measure of the maximum useful work that can be obtained under standard
condition.
The electrical energy or electrical work is equal to the product of emf of the cell and the electrical
charge that flow through the external circuit.
Wmax = nFEcell
Where, n is the no of moles of electrons transferred through wire
F is faraday constant i.e., 96500 coloumbs
Ecell is the emf of the cell.
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According to the thermodynamics, the free energy change (∆G) for a process is equal to the
maximum work that can be derived from a cell.
Wmax = ∆G
∆G = -nFEcell and ∆G0 = -nFE0cell
∆G0 = Standard free energy change
E0cell = Standard cell potential
The change in free energy is given by
∆ =∆ + Where Q= reaction quotient
On substituting the above expression

− =− +
On dividing the equation with –nF

= −

Where, R= gas constant= 8.314 J; T= Temp in Kelvin, 298K; F= Faraday 96500 C

.
= −
[ ]
Where, =
[ ]

. [ ]
= −
[ ]
For a given general reaction
aA + bB cC + dD

. [ ] [ ]
= −
[ ] [ ]
SOLVED PROBLEM. Calculate the emf of the cell.
Zn | Zn+ (0.001M) || Ag + (0.1M) | Ag
The standard potential of Ag+|Ag half-cell is + 0.80 V and Zn2+|Zn is – 0.76 V.
SOLUTION: Step 1. Write the half-cell reactions of the anode and the cathode. Then add the anode
and cathode half reactions to obtain the cell reaction and the value of E°cell.
Cathode : 2Ag+ + 2e– 2Ag E° = +0.80 V
Anode : Zn Zn 2+
+ 2e –
E° = – 0.76 V

Cell reaction: Zn + 2Ag+ Zn2+ + 2Ag E° = 1.56 V

Step2.

=
[ ]
Substitute the given values in the Nernst equation and solving for E cell, we have
 11

.
= −

.
= . −
[ ]
.
= . −
[ ]
= . + . = .
Nernst equation of electrode potential:
We know experimentally that the potential of a single electrode or half-cell varies with the
concentration of ions in the cell. In 1889 Walter Nernst derived a mathematical relationship which enable
us to calculate the half-cell potential, E, from the standard electrode potential, E°, and the temperature
of the cell.
This relation known as the Nernst equation can be stated as
.
= − ( )

Where E° = standard electrode potential; R = gas constant; T = Kelvin temperature

n = number of electrons transferred in the half-reaction F = Faraday of electricity
Q = reaction quotient
Calculation of Half-cell potential:
For a reduction half-cell reaction
Mn+ + ne– M
The Nernst equation takes the form
.
= − ( )

The concentration of solid metal [M] is considered as unity. Therefore, the Nernst equation can be
written as
.
= − ( )
[ ]
Substituting the values of R, F and T at 25°C, the quantity 2.303 RT/F comes to be 0.0592. Thus, the
Nernst equation (3) can be written in its simplified form as
.
= −
[ ]
This is the equation for a half-cell in which reduction occurs. In case it is an oxidation reaction, the sign
of E will have to be reversed.
SOLVED PROBLEM. What is the potential of a half-cell consisting of copper electrode in 0.015M CuSO4
solution at 25°C, E° =0.34 V.
SOLUTION: The half-cell reaction is
Cu2+ + 2e– Cu (reduction)
The Nernst equation for the oxidation half-cell reaction is
 12

. [ ]
= −
[ ]
The number of electrons transferred n = 2 and E° = 0.34 V. and solid copper [Cu] concentration is unity.
Substituting these values in the Nernst equation we have
.
= . −
.
= .
Calculation of Equilibrium constant for the cell reaction
The Nernst equation for a cell is
.
= −

At equilibrium ∆G=0 and E=0 and Q=K., Then

.
= −

=
.
SOLVED PROBLEM. Calculate the equilibrium constant for the reaction between silver nitrate and
metallic zinc.
SOLUTION: Step 1. Write the equation for the reaction

Zn + 2Ag+ Zn2+ + 2Ag E° = 1.56 V

Step 2. Substitute values in the Nernst equation at equilibrium

=
.
= . − .
− . =− .
− .
= =
. − .
K = 1×1052
Quinhydrone Electrode: It involves the redox reaction between quinone (Q) and hydroquinone (QH2).

Q + 2H+ + 2e– QH2

Figure.7 The Quinhydrone
electrode.
 13

The hydroquinone half-cell consists of

a platinum strip immersed in a saturated
solution of quinhydrone at a definite H+ ion
concentration. Quinhydrone is a molecular
compound which gives equimolar amounts of
quinone and hydroquinone in solution. The
electrode system may be represented as
Pt | QH2, Q, H+
The potential developed is measured against
a hydrogen electrode or calomel electrode. Fig.7 Quinhydrone electrode coupled with
Q + 2H + 2e
+ –
QH2 saturated calomel electrode (dipping calomel
electrode)
The electrode potential at 25 c is given by
0

Nernst equation

. [ ]
= −
[ ][ ]

. [ ]
= −
[ ][ ]

Since quinone and hydroquinone are taken in equimolar amounts

i.e., [Q] = [QH2] = 1

.
= −
[ ]

.
= + [ ]

= + . [ ]

= − .

[− [ ] = ]

= . − .

Since the electrode potential of the quinhydrone electrode depends upon the concentration of
hydrogen ions, it can be used for the determination of pH value just like a hydrogen electrode.
Construction: Quinhydrone electrode can very easily be set up by adding a pinch quinhydrone powder
(a sparingly soluble solid) to the experimental solution with stirring, until the solution is saturated and
a slight excess of it remains undissolved. Then, indicator electrode, usually of bright platinum, is inserted
in it.
 14

For determining the pH value, this half-cell is combined with any other reference electrode,
usually saturated calomel electrode and the emf of the cell so-formed (Fig. 7) is determined
potentiometrically. The complete cell may be represented as

(-) Hg, Hg2Cl2 | KCl (satd) || H2Q, Q, H+(unknown) | Pt (+)

= ER− EL

= Equinhydrone − Ecalomel

= (0.6994v – 0.0592 V pH) ‒ 0.2422

. − . ‒
=
.
Advantages:

1. The electrode is very easy to setup. 2) The pH value obtained is very accurate. 3) Very small
quantities of the solution are sufficient for the measurement.

Limitations: The electrode cannot be used in alkaline solutions (pH>8.5) 2) It cannot be used in
solutions containing redox system, which would react with either quinhydrone or quinone. (Ex: Fe+2,
Mno2, etc.,)

Potentiometric titrations
A titration in which the equivalent or the end point of the reaction is determined with the help
of measurement of potential of a reaction mixture is known as potentiometric titrations.
In a potentiometric titration, a suitable electrode immersed in the solution to be titrated acts as
the ‘indicator electrode’ (Fig. 7). The indicator electrode is paired with a reference electrode and the
two electrodes are connected to an electronic voltmeter. The emf of the indicator electrode changes
gradually with the change of concentration of ions caused by the addition of titrant from the burette.
The equivalence Point is indicated by a sharp change in cell potential.
Since the reference electrode potential has a constant value, any change in the indicator
electrode potential is reflected by a similar change in the cell potential. Therefore, the equivalence point
can be found by plotting a graph between the cell emf and the volume of titrant added from the burette.
A sharp rise of the curve shows the equivalence point and the corresponding volume on the graph is the
volume of the solution used for the titration.
The potentiometric titrations may be of three types:
1) Acid-base titrations 2) Oxidation-reduction titrations 3) Precipitation titrations
Acid-base titrations:
Potentiometric measurement of EMF of a cell constructed using the test solution, HCl is used for
locating the end point in acid-base titration. The cell consists of a Saturated Calomel Electrode (SCE) as
reference electrode and quinhydrone (Q, QH2) as an indicator electrode. The EMF of the cell depends on
the H+ concentration of test solution. The quinhydrone is a powdered organic substance having
equimolar mixture of quinone (Q) and hydroquinone (QH2). When a pinch of hydroquinone is added to
 15

the acid solution in contact with platinum electrode, quinone, hydroquinone and H + ions form a reversible
redox system.

Q + 2H+ + 2e- QH2

Initially, before titration, the EMF measured will be high due to high H + concentration and on
addition of NaOH solution, it decreases due to decrease in concentration of H+ as a consequence of
neutralization to form H2O.
HCl + NaOH NaCl + H2O
After neutralization, the EMF decreases even below zero as indicated by -ve value. The cell can
be represented as:
(-) Hg, Hg2Cl2(s) | KCl (saturated) || H+ ions ( c=?), quinhydrone | Pt (+)
An oxidation takes place on the Calomel electrode and reduction tales place at quinhydrone
electrode.
= ER− EL

= Equinhydrone − Ecalomel

= (0.6994v – 0.0592 V pH) ‒ 0.2422

. − . ‒
=
.

After each addition, the EMF of the cell is recorded. The EMF is then plotted against the volume
of alkali added. The shape of the curve for the titration of a strong acid against strong alkali (HCl versus
NaOH) is shown in Fig.8a. The steepest portion of the curve indicates the equivalent point of the titration.
However, the steepness of the curve is less marked and it is difficult to judge the end-point. So
now, we plot the curve, ΔE/ΔV against the volume of alkali used (Fig. 8b). The maximum of the curve
indicates the end-point.

Fig. 8 Curves of Acid-base titrations

Oxidation-reduction titrations:
Potentiometric measurement of EMF can be used for locating endpoint in the redox titration.
The cell consists of a saturated calomel electrode (SCE) as reference electrode and the platinum
electrode dipped in Fe+2 (test) solution as an indicator electrode. The EMF of the electrode depends upon
the ratio, [Fe +3/Fe +2].
 16

The cell can be represented as: (-) Hg, Hg2Cl2(s) |KCl (sat) || Fe+3, Fe+2 |Pt (+)
.
The EMF of the cell is given by, = + −
e.g.: Fe+2 titrated against KMNO4
Initially, the concentration of Fe +2 in the solution is very low and the EMF measured will be low
and on addition of KMnO4 solution to Fe+2 solution, the concentration of Fe+3 increases (due to oxidation
of Fe +2) and the EMF increases.
5Fe+2 + MnO4- + 8H+ 5Fe +3 + Mn+2 + 4H2O
Finally, at the end point the EMF increases sharply and an inflection in the titration curve (Fig.
9) can be observed due to sharp decrease in Fe+2 concentration.

Fig. 15 Potentiometric titration curve of redox titration.

Precipitation titrations:
In precipitation reactions or titrations also, an electrode reversible to one of the ions is involved.
e.g.: Titrations of AgNO3 with NaCl where AgCl precipitates out.
Here Silver electrode is used along with Calomel electrode. The NaCl is taken in the burette and
AgNO3 is taken in the beaker containing electrodes. The emf of the cell is measured and plotted against
volume of NaCl added. The steepest portion of the curve indicates the equivalent point of the titration.
Advantages of potentiometric titrations:
1) Coloured solutions can be titrated without the use of an indicator.
2) Even weak acids, weak base titrations can be carried out.
3) Results obtained are very accurate.
4) The apparatus required is inexpensive, reliable and readily available.
5) can be used in the titrations of mixture of acids, bases, halides
 17

Ion-selective electrode (ISE): Ion selective electrode possess the ability to respond only to certain
specific ions, there by developing a potential with respect to that species only in a mixture and ignoring
the other ions totally.

Example: glass electrode is only H+ ions selective.

Glass electrode: When two solutions of different PH values are separated by

a thin glass membrane, there develops a difference of potential between the
two surfaces of the membrane. The potential difference developed is
proportional to the difference in PH values. The Glass membrane functions as
an Ion-exchange Resin. There is an equilibrium set up between the Na+ ions
of glass and H+ ions in the solution. The potential difference varies with the
H+ ions concentration of the solution. The electrode potential is given by

At 250c = − 0.0592 [ ]

= + 0.0592

Construction: Glass electrode

A glass electrode is a type of ion-selective electrode and consists of thin-walled glass bulb
containing AgCl coated Ag electrode or simply a Pt electrode in 0.1M HCl (Fig. 11). The glass electrode
may be shown schematically

Ag│AgCl(s), HCl (0.1M)│Glass or Pt,0.1M HCl│Glass

Determination of pH of solution by using glass electrode: Glass electrode is dipped in a solution

whose PH is to be determined and it is
coupled with saturated calomel electrode
(Fig. 12). The EMF of the cell is measured.
The EMF of complete cell is given by

Ecell = Eright - Eleft

Ecell = ESCE - EGLASS

= 0.2422 − [ + 0.0592 ]

0.2422 V − E −E
P =
0.0592 V
The value of the potential of calomel
electrode is known while Ecell can be found
experimentally. Therefore, we can find PH
of a given solution if E° G is known. The
E°G value of a glass electrode can be
determined by using a solution of known Fig. 12 a) Glass electrode and saturated calomel electrode
PH in the cell and measuring Ecell. This immersed in a solution of unknown pH. b) Combined
value of E° G is constant for a particular probe or electrode of Glass electrode and reference
glass electrode and can be used for any electrode (saturated Calomel electrode)
subsequent determinations of pH of
unknown solutions
 18

Advantages of glass electrode: 1) It is simple and can easily be used. 2) Equilibrium is rapidly achieved
3) The results are accurate 4) It is not easily poisoned.

Limitations of glass electrode: 1) The glass electrode can be used in solutions with P H range of 0-10.
Electrodes composed of special glasses can be used for measurements up to a PH of 12. However, above
12 PH cations of solution affect the glass interface and render the electrode useless. 2) Although, glass
membranes of electrode is very thin, yet its resistance is very high, which cannot be measured by
ordinary potentiometers. It is therefore, necessary to use special electronic potentiometers.
 BATTERIES

Battery: battery is an electrochemical cell, or often several electrochemical cells connected in series
that can be used as a source of direct electric current at a constant voltage.
Classification: batteries are classified into three categories depending on their recharging
capabilities.
1) Primary battery (non rechargeable) 2) Secondary battery (rechargeable) 2) Flow battery
(Fuel cell)
Primary battery: Primary battery is a cell in which the cell reaction is not reversible. Thus, once
the chemical reaction takes place to release the electrical energy, the cell gets exhausted. They are
use and throw type.
Example: dry or leclanche cell, lithium cells

Secondary battery: Secondary battery is a cell in which the cell reaction is reversible. They are
rechargeable cells. Once the battery gets exhausted, it can be recharged.
Example: Lead-acid cell (storage cell), Nickel-Cadmium cell, lithium-ion cells etc.,

Differences between primary and secondary batteries

Primary batteries Secondary batteries

1.Cell reaction is irreversible 1.Cell reaction is reversible

2.Must be discarded after use 2.May be recharged

3.Have relatively short shelf life 3.Have long shelf life

4.Function only as galvanic cells 4. Functions both galvanic cell& as electrolytic cell.

5.They cannot be recharged 5.They can be recharged

6.EX:Dry cell, Li-MnO2 6.EX: Lead-acid cell (storage cell), Nickel-Cadmium

cell, lithium ion cells etc.,
Lithium cells: Lithium cells belong to primary cells. The cells having lithium anodes are called
lithium cells. Based on the cathode used, the lithium cells can be classified into two categories.

1) Lithium cells with solid cathodes, 2) Lithium cells with liquid cathodes

1) Lithium cells with solid cathodes: These batteries may have solid or liquid electrolyte.
The Most widely used cell is Li-MnO2 cell (3V).

Anode: Lithium metal

Cathode: MnO2

Electrolyte: Lithium perchlorate in propylene

carbonate or dimethoxyethane

voltage (EMF) 3V

MnO2 should be heated to over 300 0C to remove

water before keeping it in the cathode, thereby increasing
the efficiency of the cell.

Cell reactions:

At anode: Li Li+ + e-

At cathode: MnO2+ e- MnO2-

Net Cell reaction: Li+ MnO2 LiMnO2

Li-MnO2 primary battery

Applications:
1) The coin type cells are used in watches and
calculators.
2. Cylindrical cells are used in fully automatic cameras

Lithium-ion battery: A lithium-ion battery (sometimes Li-ion battery or LIB) is a member of

a family of rechargeable battery types in which lithium ions move from the negative electrode to the
positive electrode during discharge and back when charging.

Anode Graphite (Lithium intercalated graphite)

Cathode lithium cobalt oxide (LiCoO2)

Electrolyte lithium hexafluorophosphate (LiPF 6), lithium tetrafluoroborate (LiBF4), lithium

perchlorate (LiClO4), etc., which are dissolved in organic solvents such as
ethylene carbonate and dimethyl carbonate.

Voltage 3.7 V

The three primary functional components of a lithium-ion battery are the positive and
negative electrodes and electrolyte. The most commercially popular negative electrode is graphite.
The positive electrode is generally one of
three materials: a layered oxide (such
as lithium cobalt oxide), a polyanion (such
as lithium iron phosphate) or a spinel (such
as lithiummanganese oxide).
The electrolyte is typically a mixture
of organic carbonates such as ethylene
carbonate or diethyl carbonate containing
complexes of lithium ions. These non-
aqueous electrolytes generally use non-
coordinating anion salts such as lithium
hexafluorophosphate (LiPF6), lithium Lithium ion battery
hexafluoroarsenate monohydrate (LiAsF6),
lithium perchlorate (LiClO4), lithium tetrafluoroborate (LiBF4) and lithium triflate (LiCF3SO3).

Both electrodes allow lithium ions to move in and out of their interiors.
During insertion (or intercalation) ions move into the electrode. During the reverse process,
extraction (or deintercalation), ions move back out. When a lithium-ion based cell is discharging,
the positive Lithium ion moves from the negative electrode (usually graphite = "C 6" below) and
enters the positive electrode (lithium cobalt oxide). When the cell is charging, the reverse occurs.

The cell reaction:

At anode: Lix C xLi+ + C + xe-

At cathode: Li1-xCoO2 + xLi+ + xe- LiCoO2

Lix C + Li 1-x CoO2 C + LiCoO2

Applications:

1) Lithium-ion batteries are common in consumer electronics

2) They are one of the most popular types of rechargeable batteries for portable electronics, with a
High energy density, small memory effect, and only a slow loss of charge when not in use
3) LIBs are also growing in popularity for military, battery electric vehicle and aerospace applications
Advantages: Portable and rechargeable
Fuel cells: The device in which chemical energy of fuel-oxidant system converted into electrical
energy is known as fuel cell.
Principle: The basic principle of the fuel cell is same as that of an electrochemical cell. The fuel cell
operates like a galvanic cell. The only difference is that the fuel and oxidant stored outside of the
cell. Fuel and oxidant are supplied continuously and separately to the electrodes at which they
undergo redox reactions.
Fuel + oxidant Oxidation product + Electricity.
Examples: 1) Hydrogen – oxygen fuel cell 2) Methanol-oxygen fuel cell
1) Hydrogen – Oxygen fuel cell: A typical example of pollution free cell is H 2 –O2 fuel cell in
which the Fuel is hydrogen and the oxidizer is oxygen.

Construction and working:

Hydrogen – oxygen fuel cell consists of two porous electrodes made up of compressed carbon coated
with small amount of catalysts (Pt, Pd, Ag) and KOH or NaOH solution as the electrolyte.

During working, Hydrogen (the fuel) is bubbled through the anode compartment, where
it is oxidized. The oxygen (oxidizer) is bubbled through the cathode compartment, where it is
reduced. The following cell reactions occur.

At Anode: 2H2 + 4OH- 4H2O + 4e-

At cathode: O2 + 4e- + 2H2O 4OH -

Net reaction: 2H2 + O2 2H2O

The product discharged is water and the standard EMF of the cell is 1.23V

Applications: Hydrogen – oxygen fuel cells are used as auxiliary energy source in space vehicles
(e.g., Apollo space craft), submarines other military vehicles.
Limitations: 1) High energy cost of generating Hydrogen fuel
2) Problems in handling, storage and distribution of highly flammable hydrogen fuel.

2) Methanol-oxygen fuel cell:

Methanol – oxygen fuel cell, methanol used as the fuel and oxygen or air as the oxidant.

Construction and working: It consists of two electrodes separated by a proton exchange

membrane (PEM) and connected via an external circuit that allows the conversion of free energy
from the chemical reaction of methanol with air or oxygen to be directly converted into electrical
energy. Aqueous methanol is fed at the anode side. It diffuses through the diffusion layer to the
catalytic layer where it is electrochemically oxidized into mainly carbon dioxide, protons and
electrons. Protons formed during this reaction diffuse through the membrane to the cathode catalytic
layer. They participate in the reduction of oxygen to form water at cathode side. Oxygen may be
pure but can also come from air.
Cell reactions:

At anode: CH3OH+H2O CO2+6H ++6e-

At cathode: 3/2O2+6H++6e- 3H2O

Net reaction: CH3OH+3/2O2 CO2+2H2O

Discharging voltage (EMF) of methanol-oxygen fuel cell is 1.19V

Applications:
1. Space craft applications
2. Fuel cell vehicles.

Limitations: 1) During the methanol oxidation reaction, CO2 is formed which is strongly absorbed
on to platinum catalyst, reducing the surface area and lowering the performance.
2)

Methanol-oxygen fuel cell

Methanol is toxic and flammable.

3) Limited in the power supply.
Advantages of fuel cell:
1) It can’t give pollution
2) More efficiency
3) Maintenance cost is low
4) These are compact and transportable
5) They have quick start system
6) They save fossil fuels
7) The by products are environmentally acceptable
8) It has high reliability in electricity generation.