Lakshya JEE AIR (2025)
Electrochemistry
Electrochemistry
1. Introduction:
Chemical reactions can be used to produce electrical energy, conversely, electrical energy can be used to
carry out chemical reactions that do not proceed spontaneously.
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. The subject is of importance both for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine, fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells convert chemical energy into electrical energy and are
used on a large scale in various instruments and devices. The reactions carried out electrochemically can
be energy efficient and less polluting. Therefore, study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of sensory signals through cells to brain and vice
versa and communication between the cells are known to have electrochemical origin. Electrochemistry,
is therefore, a very vast and interdisciplinary subject. In this Unit, we will cover only some of its
important elementary aspects.
Electrochemical Cell:
If a metal electrode is immersed in an aqueous solution containing cations of that metal, an equilibrium
that leads to negative charge formation on the electrode is established. This configuration of electrode
and solution is called a half-cell. Two half-cells can be combined to form an electrochemical cell. The
equilibrium condition in an electrochemical cell is that the electrochemical potential, rather than the
chemical potential, of a species is the same in all parts of the cell. The electrochemical potential can be
changed through the application of an electrical potential external to the cell. This allows the direction of
spontaneous change in the cell reaction to be reversed. Electrochemical cells can be used to determine the
equilibrium constant for the cell reaction. Electrochemical cells can also be used to provide power, in
which case they are called batteries. Electrochemical cells in which the reactants can be supplied
continuously are called fuel cells.
Electrochemical cell are of two types
Galvanic cells or Voltaic cell Electrolytic cell
A spontaneous chemical reaction An electric current drives a
generates an electric current. nonspontaneous reaction.
Thus, the two types of cells are reverse of each other.
PHYSICS WALLAH 1
, ELECTROCHEMISTRY
2. Galvanic Cell or Voltaic Cell:
A galvanic cell is an electrochemical cell that converts the chemical energy of a spontaneous redox
reaction into electrical energy. In this device the Gibbs energy of the spontaneous redox reaction is
converted into electrical work which may be used for running a motor or other electrical gadget like
heater, fan, geyser, etc.
Constriction of galvanic cell:
Anode: Some metals (which are reactive) are found to have tendency to go into the solution phase when
these are placed in contact with their ions or their salt solutions.
For example: Zn rod is placed in ZnSO4 solution.
If a Zn electrode is partially immersed in an aqueous solution of ZnSO4, an equilibrium is established
between Zn(s) and Zn2+(aq) as a small amount of the Zn goes into solution as Zn2+(aq) ions as depicted in
above figure. Zn(s) Zn2+ +2e–
However, the electrons remain on the Zn electrode. Therefore, a negative charge builds up on the Zn
electrode, and a corresponding positive charge builds up in the surrounding solution. This charging leads
to a difference in the electrical potential between the electrode and the solution, which we call the half-
cell potential or electrode potential. This particular electrode is known as anode.
* On anode oxidation will take place. (release of electrons).
* To act as source of electrons.
* It is of negative polarity.
* The electrode potential is represented by EZn(s)/Zn2+ (aq.)
Cathode: Some metals (Cu, Ag, Au etc.,) are found to have the opposite tendency i.e., when placed in
contact with their aqueous ions, the ions from the solution will get deposited on the metal rod.
The following equilibrium will be established: Cu2+ +2e– Cu(s).
So, rod will have deficiency of electron (positive charge). Extra negative charge will surround this
positively charged rod and form double layer. An electrical double layer is developed in the system and
hence a potential difference is created between the rod and the solution which is known as half cell
potential or electrode potential. This will be known as cathode.
* At cathode reduction will take place. (gain of electrons will take place)
* To act as sink of electron.
* Positive polarity will be developed.
* Their electrode potential can be represented by : ECu2+ (aq.)/Cu(s)
PHYSICS WALLAH 2
, ELECTROCHEMISTRY
Is whereoxidation occurs Is where reduction occurs
Anode: Is whereelectronsareproduced Cathode: Is whereelectronsareconsumed
Hasa negativesign Hasa positivesign
* Half-cell potentials cannot be measured directly. They are measured relative to one another rather
than absolutely. To understand how this is done, it is useful to consider an electrochemical cell,
which consists of two half-cells, such as the one shown figure. This particular cell is known as the
Daniell cell, after its inventor. On the left, a Zn electrode is immersed in a solution of ZnSO4(aq).
2−
The solute is completely dissociated to form Zn2+(aq) and SO4 (aq). On the right, a Cu electrode is
2−
immersed in a solution of CuSO4, which is completely dissociated to form Cu 2+(aq) and SO4 (aq).
The two half-cells are connected by an ionic conductor known as a salt bridge. The salt bridge
consists of an electrolyte such as KCl suspended in a gel. A salt bridge allows current to flow
between the half-cells while preventing the mixing of the solutions. A metal wire fastened to each
electrode allows the electron current to flow through the external part of the circuit.
Salt bridge:
A salt bridge is a U–shaped inverted tube that contains a gel permeated with an inert electrolyte.
Generally, tube is filled with a paste of agar-agar powder with a natural electrolyte/generally not common
to anodic/cathodic compartment with porous plugs at each mouth of tube. The ions of the inert electrolyte
do not react with other ion in the solution and the ions are not oxidised or reduced at the electrodes.
The electrolyte in salt bridge should be such that speed of it's cation equals speed of it's anion in electrical
field. For that charge and sign of the ions should be almost equal
i.e. Mobility of cation = Mobility of anion
KCl is generally preferred but KNO3 or NH4NO3 can also be used.
* If Ag+, Hg22+, Pb2+, Tl+ ions are present in a cell then in salt bridge KCl is not used because there can
be formation of precipitate of AgCl, Hg2Cl2, PbCl2 or TlCl at mouth of tube which will prevent the
migration of ions and its functioning will stop.
PHYSICS WALLAH 3
, ELECTROCHEMISTRY
Functions of Salt Bridge:
(i) It connects the solution of two half cell to complete the circuit.
(ii) It minimize the liquid junction potential. The potential difference between the junction of two
liquids. (Liquid-Liquid Junction Potential: The potential difference which arises between two
solutions (during the progress of reaction) when in contact with each other.)
(iii) It maintains the electrical neutrality of the solution in order to give continuous flow or generation of
current.
"The simultaneous electrical neutrality of the anodic oxidation chamber and cathodic reduction
chamber is due to same mobility or velocity of K+ and NO3– ions taken into salt bridge."
(iv) If the salt bridge is removed then voltage drops to zero.
(v) It prevents mechanical mixing of two electrolytic solution.
Shorthand Notation for Galvanic Cells:
We require two half cells to produce an electrochemical cell, which can be represented by following few
rules;
* The anode half-cell is always written on the left followed on the right by cathode half cell.
* The separation of two phases (state of matter) is shown by a vertical line.
* The various materials present in the same phase are shown together using commas.
* The salt bridge is represented by a double slash (||).
* The significant features of the substance viz. pressure of a gas, concentration of ions etc. are
indicated in brackets immediately after writing the substance.
* For a gas electrode, the gas is indicated after the electrode for anode and before the electrode in case
of cathode. (i.e Pt H2/H+ or H+/H2 Pt)
Ex.1 Write short hand notation for the following reaction,
Sn2+ (aq) + 2Ag+ (aq) ⎯⎯ → Sn4+ (aq) + 2Ag(s).
Sol. The cell consists of a platinum wire anode dipping into an Sn+2 solution and a silver cathode dipping into
an Ag+ solution therefore
Pt(s) | Sn2+(aq), Sn4+ (aq) || Ag+ (aq) | Ag(s).
Ex.2 Write the electrode reaction and the net cell reaction for the following cells. Which electrode would be
the positive terminal in each cell?
(a) Zn | Zn2+ || Br–, Br2 | Pt (b) Cr| Cr3+ || I–, I2 | Pt
(c) Pt | H2, H+ || Cu2+ | Cu (d) Cd | Cd2+ || Cl–, AgCl | Ag
Sol. (a) Oxidation half cell reaction, Zn ⎯⎯
→ Zn2+ + 2e–
reduction half cell reaction, Br2 + 2e– ⎯⎯
→ 2Br–
Net cell reaction Zn + Br2 ⎯⎯
→ Zn2+ + 2Br–
(Positive terminal : cathode Pt)
(b) Oxidation half reaction, [Cr ⎯⎯
→ Cr3+ + 3e–] x 2
reduction half reaction, [I2 + 2e– ⎯⎯ → 2I–] x 3
Net cell reaction 2Cr + 3I2 2Cr3+ + 6I–
(Positive terminal : cathode Pt)
PHYSICS WALLAH 4
Electrochemistry
Electrochemistry
1. Introduction:
Chemical reactions can be used to produce electrical energy, conversely, electrical energy can be used to
carry out chemical reactions that do not proceed spontaneously.
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. The subject is of importance both for theoretical and practical considerations. A large
number of metals, sodium hydroxide, chlorine, fluorine and many other chemicals are produced by
electrochemical methods. Batteries and fuel cells convert chemical energy into electrical energy and are
used on a large scale in various instruments and devices. The reactions carried out electrochemically can
be energy efficient and less polluting. Therefore, study of electrochemistry is important for creating new
technologies that are ecofriendly. The transmission of sensory signals through cells to brain and vice
versa and communication between the cells are known to have electrochemical origin. Electrochemistry,
is therefore, a very vast and interdisciplinary subject. In this Unit, we will cover only some of its
important elementary aspects.
Electrochemical Cell:
If a metal electrode is immersed in an aqueous solution containing cations of that metal, an equilibrium
that leads to negative charge formation on the electrode is established. This configuration of electrode
and solution is called a half-cell. Two half-cells can be combined to form an electrochemical cell. The
equilibrium condition in an electrochemical cell is that the electrochemical potential, rather than the
chemical potential, of a species is the same in all parts of the cell. The electrochemical potential can be
changed through the application of an electrical potential external to the cell. This allows the direction of
spontaneous change in the cell reaction to be reversed. Electrochemical cells can be used to determine the
equilibrium constant for the cell reaction. Electrochemical cells can also be used to provide power, in
which case they are called batteries. Electrochemical cells in which the reactants can be supplied
continuously are called fuel cells.
Electrochemical cell are of two types
Galvanic cells or Voltaic cell Electrolytic cell
A spontaneous chemical reaction An electric current drives a
generates an electric current. nonspontaneous reaction.
Thus, the two types of cells are reverse of each other.
PHYSICS WALLAH 1
, ELECTROCHEMISTRY
2. Galvanic Cell or Voltaic Cell:
A galvanic cell is an electrochemical cell that converts the chemical energy of a spontaneous redox
reaction into electrical energy. In this device the Gibbs energy of the spontaneous redox reaction is
converted into electrical work which may be used for running a motor or other electrical gadget like
heater, fan, geyser, etc.
Constriction of galvanic cell:
Anode: Some metals (which are reactive) are found to have tendency to go into the solution phase when
these are placed in contact with their ions or their salt solutions.
For example: Zn rod is placed in ZnSO4 solution.
If a Zn electrode is partially immersed in an aqueous solution of ZnSO4, an equilibrium is established
between Zn(s) and Zn2+(aq) as a small amount of the Zn goes into solution as Zn2+(aq) ions as depicted in
above figure. Zn(s) Zn2+ +2e–
However, the electrons remain on the Zn electrode. Therefore, a negative charge builds up on the Zn
electrode, and a corresponding positive charge builds up in the surrounding solution. This charging leads
to a difference in the electrical potential between the electrode and the solution, which we call the half-
cell potential or electrode potential. This particular electrode is known as anode.
* On anode oxidation will take place. (release of electrons).
* To act as source of electrons.
* It is of negative polarity.
* The electrode potential is represented by EZn(s)/Zn2+ (aq.)
Cathode: Some metals (Cu, Ag, Au etc.,) are found to have the opposite tendency i.e., when placed in
contact with their aqueous ions, the ions from the solution will get deposited on the metal rod.
The following equilibrium will be established: Cu2+ +2e– Cu(s).
So, rod will have deficiency of electron (positive charge). Extra negative charge will surround this
positively charged rod and form double layer. An electrical double layer is developed in the system and
hence a potential difference is created between the rod and the solution which is known as half cell
potential or electrode potential. This will be known as cathode.
* At cathode reduction will take place. (gain of electrons will take place)
* To act as sink of electron.
* Positive polarity will be developed.
* Their electrode potential can be represented by : ECu2+ (aq.)/Cu(s)
PHYSICS WALLAH 2
, ELECTROCHEMISTRY
Is whereoxidation occurs Is where reduction occurs
Anode: Is whereelectronsareproduced Cathode: Is whereelectronsareconsumed
Hasa negativesign Hasa positivesign
* Half-cell potentials cannot be measured directly. They are measured relative to one another rather
than absolutely. To understand how this is done, it is useful to consider an electrochemical cell,
which consists of two half-cells, such as the one shown figure. This particular cell is known as the
Daniell cell, after its inventor. On the left, a Zn electrode is immersed in a solution of ZnSO4(aq).
2−
The solute is completely dissociated to form Zn2+(aq) and SO4 (aq). On the right, a Cu electrode is
2−
immersed in a solution of CuSO4, which is completely dissociated to form Cu 2+(aq) and SO4 (aq).
The two half-cells are connected by an ionic conductor known as a salt bridge. The salt bridge
consists of an electrolyte such as KCl suspended in a gel. A salt bridge allows current to flow
between the half-cells while preventing the mixing of the solutions. A metal wire fastened to each
electrode allows the electron current to flow through the external part of the circuit.
Salt bridge:
A salt bridge is a U–shaped inverted tube that contains a gel permeated with an inert electrolyte.
Generally, tube is filled with a paste of agar-agar powder with a natural electrolyte/generally not common
to anodic/cathodic compartment with porous plugs at each mouth of tube. The ions of the inert electrolyte
do not react with other ion in the solution and the ions are not oxidised or reduced at the electrodes.
The electrolyte in salt bridge should be such that speed of it's cation equals speed of it's anion in electrical
field. For that charge and sign of the ions should be almost equal
i.e. Mobility of cation = Mobility of anion
KCl is generally preferred but KNO3 or NH4NO3 can also be used.
* If Ag+, Hg22+, Pb2+, Tl+ ions are present in a cell then in salt bridge KCl is not used because there can
be formation of precipitate of AgCl, Hg2Cl2, PbCl2 or TlCl at mouth of tube which will prevent the
migration of ions and its functioning will stop.
PHYSICS WALLAH 3
, ELECTROCHEMISTRY
Functions of Salt Bridge:
(i) It connects the solution of two half cell to complete the circuit.
(ii) It minimize the liquid junction potential. The potential difference between the junction of two
liquids. (Liquid-Liquid Junction Potential: The potential difference which arises between two
solutions (during the progress of reaction) when in contact with each other.)
(iii) It maintains the electrical neutrality of the solution in order to give continuous flow or generation of
current.
"The simultaneous electrical neutrality of the anodic oxidation chamber and cathodic reduction
chamber is due to same mobility or velocity of K+ and NO3– ions taken into salt bridge."
(iv) If the salt bridge is removed then voltage drops to zero.
(v) It prevents mechanical mixing of two electrolytic solution.
Shorthand Notation for Galvanic Cells:
We require two half cells to produce an electrochemical cell, which can be represented by following few
rules;
* The anode half-cell is always written on the left followed on the right by cathode half cell.
* The separation of two phases (state of matter) is shown by a vertical line.
* The various materials present in the same phase are shown together using commas.
* The salt bridge is represented by a double slash (||).
* The significant features of the substance viz. pressure of a gas, concentration of ions etc. are
indicated in brackets immediately after writing the substance.
* For a gas electrode, the gas is indicated after the electrode for anode and before the electrode in case
of cathode. (i.e Pt H2/H+ or H+/H2 Pt)
Ex.1 Write short hand notation for the following reaction,
Sn2+ (aq) + 2Ag+ (aq) ⎯⎯ → Sn4+ (aq) + 2Ag(s).
Sol. The cell consists of a platinum wire anode dipping into an Sn+2 solution and a silver cathode dipping into
an Ag+ solution therefore
Pt(s) | Sn2+(aq), Sn4+ (aq) || Ag+ (aq) | Ag(s).
Ex.2 Write the electrode reaction and the net cell reaction for the following cells. Which electrode would be
the positive terminal in each cell?
(a) Zn | Zn2+ || Br–, Br2 | Pt (b) Cr| Cr3+ || I–, I2 | Pt
(c) Pt | H2, H+ || Cu2+ | Cu (d) Cd | Cd2+ || Cl–, AgCl | Ag
Sol. (a) Oxidation half cell reaction, Zn ⎯⎯
→ Zn2+ + 2e–
reduction half cell reaction, Br2 + 2e– ⎯⎯
→ 2Br–
Net cell reaction Zn + Br2 ⎯⎯
→ Zn2+ + 2Br–
(Positive terminal : cathode Pt)
(b) Oxidation half reaction, [Cr ⎯⎯
→ Cr3+ + 3e–] x 2
reduction half reaction, [I2 + 2e– ⎯⎯ → 2I–] x 3
Net cell reaction 2Cr + 3I2 2Cr3+ + 6I–
(Positive terminal : cathode Pt)
PHYSICS WALLAH 4
