CLASS XII · PHYSICAL CHEMISTRY · PREMIUM
NOTES EDITION
ELECTRO
CHEMISTRY
Cells · Electrode Potential · Nernst Equation · Conductance
Electrolysis · Batteries · Corrosion
01 Electrochemical & Galvanic Cells
02 Electrode Potential & EMF
03 Nernst Equation
04 Gibbs Energy & Equilibrium
05 Conductance of Solutions
06 Molar Conductivity & Kohlrausch Law
07 Electrolytic Cells & Faraday's Laws
08 Batteries & Fuel Cells
09 Corrosion
HANDWRITTEN-STYLE STUDY NOTES 9 SECTIONS
,CELLS
ELECTROCHEMISTRY · SECTION 01 GALVANIC CELLS
01 Electrochemical Cells
DEFINITION
Electrochemistry is the branch of chemistry that deals with the study of production of
electricity from energy released during a spontaneous chemical reaction, and the use of
electrical energy to carry out non-spontaneous chemical transformations.
Redox reactions set up the basis of electrochemical cells — they involve spontaneous oxidation
and reduction reactions. An electrochemical cell is a device in which chemical energy of the
redox reaction is converted into electrical energy.
Chemical Energy ⇌ Electrical Energy
FORWARD: ELECTROCHEMICAL CELL | REVERSE: ELECTROLYTIC CELL
Galvanic (Voltaic) Cells
Electrochemical cells which convert chemical energy of a spontaneous redox reaction into
electrical energy are called Galvanic or Voltaic cells.
An electrode is dipped into the solution of its own salt — this combination is called a half cell
or redox couple. In a half cell, either a reduction or an oxidation reaction occurs.
Daniell cell is a classic example of a Galvanic cell.
AT ANODE AT CATHODE
Oxidation occurs (loss of electrons) Reduction occurs (gain of electrons)
The Daniell Cell
Component Detail
Electrodes Zinc & Copper
Electrolytes ZnSO₄(aq) & CuSO₄(aq)
Voltmeter
V
SALT BRIDGE
Zn electrode Cu electrode
ZnSO₄(aq) CuSO₄(aq)
Fig 1.1 Daniell Cell — Zn / ZnSO₄ || CuSO₄ / Cu
Electrochemistry 01
, CELLS
ELECTROCHEMISTRY · SECTION 01 WORKING & SALT BRIDGE
Working of the Daniell Cell
Zn(s) → Zn²⁺ + 2e⁻ (oxidation half-reaction)
Cu²⁺ + 2e⁻ → Cu(s) (reduction half-reaction)
Zn(s) + Cu²⁺ → Zn²⁺ + Cu(s)
The two half cells are connected by a metallic wire through a voltmeter and a switch, externally.
The electrolytes of the two cells are connected internally through a salt bridge.
Important Features
1. An electrical potential of 1.1 V is generated when the concentration of Zn²⁺ and Cu²⁺ ions is 1
mol dm⁻³.
2. If an external potential is applied and increased slowly, the reaction continues until the
opposing voltage reaches 1.1 V, at which point the reaction stops and no current flows.
3. If the external potential is increased further, the reaction reverses direction — the cell now
functions as an electrolytic cell.
Salt Bridge
DEFINITION
A U-shaped glass tube containing electrolytes (KCl, KNO₃, NH₄Cl, etc.) mixed in gelatin or
agar-agar.
It completes the circuit and maintains the electrical neutrality of solutions in each of the
respective half-cells.
For smooth flow of current, the mobility (transport number) of the cation of the electrolyte
should equal that of the anion. e.g. in KCl, mobility of K⁺ = mobility of Cl⁻.
It should not react with the contents of either electrode.
Electrode Potential
DEFINITION
When a metal is placed in a solution of its ions, it acquires a positive or negative charge
with respect to the solution. This potential difference is called Electrode Potential. It
depends on the nature of the electrode, concentration of ions, and temperature.
OXIDATION POTENTIAL REDUCTION POTENTIAL
M → Mⁿ⁺ + ne⁻ Mⁿ⁺ + ne⁻ → M
STANDARD ELECTRODE POTENTIAL
The potential difference developed between a metal electrode and the solution of its ions
of unit molarity (1M) at 25°C (298 K).
Electrochemistry 02
NOTES EDITION
ELECTRO
CHEMISTRY
Cells · Electrode Potential · Nernst Equation · Conductance
Electrolysis · Batteries · Corrosion
01 Electrochemical & Galvanic Cells
02 Electrode Potential & EMF
03 Nernst Equation
04 Gibbs Energy & Equilibrium
05 Conductance of Solutions
06 Molar Conductivity & Kohlrausch Law
07 Electrolytic Cells & Faraday's Laws
08 Batteries & Fuel Cells
09 Corrosion
HANDWRITTEN-STYLE STUDY NOTES 9 SECTIONS
,CELLS
ELECTROCHEMISTRY · SECTION 01 GALVANIC CELLS
01 Electrochemical Cells
DEFINITION
Electrochemistry is the branch of chemistry that deals with the study of production of
electricity from energy released during a spontaneous chemical reaction, and the use of
electrical energy to carry out non-spontaneous chemical transformations.
Redox reactions set up the basis of electrochemical cells — they involve spontaneous oxidation
and reduction reactions. An electrochemical cell is a device in which chemical energy of the
redox reaction is converted into electrical energy.
Chemical Energy ⇌ Electrical Energy
FORWARD: ELECTROCHEMICAL CELL | REVERSE: ELECTROLYTIC CELL
Galvanic (Voltaic) Cells
Electrochemical cells which convert chemical energy of a spontaneous redox reaction into
electrical energy are called Galvanic or Voltaic cells.
An electrode is dipped into the solution of its own salt — this combination is called a half cell
or redox couple. In a half cell, either a reduction or an oxidation reaction occurs.
Daniell cell is a classic example of a Galvanic cell.
AT ANODE AT CATHODE
Oxidation occurs (loss of electrons) Reduction occurs (gain of electrons)
The Daniell Cell
Component Detail
Electrodes Zinc & Copper
Electrolytes ZnSO₄(aq) & CuSO₄(aq)
Voltmeter
V
SALT BRIDGE
Zn electrode Cu electrode
ZnSO₄(aq) CuSO₄(aq)
Fig 1.1 Daniell Cell — Zn / ZnSO₄ || CuSO₄ / Cu
Electrochemistry 01
, CELLS
ELECTROCHEMISTRY · SECTION 01 WORKING & SALT BRIDGE
Working of the Daniell Cell
Zn(s) → Zn²⁺ + 2e⁻ (oxidation half-reaction)
Cu²⁺ + 2e⁻ → Cu(s) (reduction half-reaction)
Zn(s) + Cu²⁺ → Zn²⁺ + Cu(s)
The two half cells are connected by a metallic wire through a voltmeter and a switch, externally.
The electrolytes of the two cells are connected internally through a salt bridge.
Important Features
1. An electrical potential of 1.1 V is generated when the concentration of Zn²⁺ and Cu²⁺ ions is 1
mol dm⁻³.
2. If an external potential is applied and increased slowly, the reaction continues until the
opposing voltage reaches 1.1 V, at which point the reaction stops and no current flows.
3. If the external potential is increased further, the reaction reverses direction — the cell now
functions as an electrolytic cell.
Salt Bridge
DEFINITION
A U-shaped glass tube containing electrolytes (KCl, KNO₃, NH₄Cl, etc.) mixed in gelatin or
agar-agar.
It completes the circuit and maintains the electrical neutrality of solutions in each of the
respective half-cells.
For smooth flow of current, the mobility (transport number) of the cation of the electrolyte
should equal that of the anion. e.g. in KCl, mobility of K⁺ = mobility of Cl⁻.
It should not react with the contents of either electrode.
Electrode Potential
DEFINITION
When a metal is placed in a solution of its ions, it acquires a positive or negative charge
with respect to the solution. This potential difference is called Electrode Potential. It
depends on the nature of the electrode, concentration of ions, and temperature.
OXIDATION POTENTIAL REDUCTION POTENTIAL
M → Mⁿ⁺ + ne⁻ Mⁿ⁺ + ne⁻ → M
STANDARD ELECTRODE POTENTIAL
The potential difference developed between a metal electrode and the solution of its ions
of unit molarity (1M) at 25°C (298 K).
Electrochemistry 02
