CBSE Class 12 Chemistry Notes Chapter 3 Electrochemistry

Electrochemistry is the study of the relationship between electrical energy and chemical reactions. Electrochemical cells convert chemical energy into electrical work or vice versa. Galvanic cells harness the spontaneous redox reactions to generate electricity, while electrolytic cells utilise electrical energy to drive non-spontaneous redox reactions. 

• The standard electrode potential, referenced to the hydrogen electrode, helps determine the standard potential of the cell, which in turn relates to standard Gibbs energy and equilibrium constants of the reactions. 
• Nernst equation explains the concentration dependence of electrode potentials. 
• The conductivity and molar conductivity of electrolytic solutions vary with the concentration and nature of the solvent. 
• Kohlrausch's law of independent migration of ions elucidates molar conductivity at infinite dilution. 
• Electrochemical principles find application in batteries, fuel cells, and understanding corrosion phenomena. 

As per the CBSE Syllabus 2023-24, the Class 12 Electrochemistry chapter is now Chapter 2 with a weightage of 9 marks in the Class 12 Chemistry Exam.

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Class 12 Chemistry Chapter 3 Notes - Electrochemistry

Electrochemical Cells

• Electrochemical cells comprise metallic electrodes immersed in electrolytic solutions.
• They exist in two forms - 
  • Galvanic cells convert chemical energy to electricity.
  • Electrolytic cells drive non-spontaneous reactions using electricity.

Electrochemical cells

Components of Electrochemical Cells

• Two Half Cells
• Two Metal Electrodes
• Voltmeter
• Salt Bridge
• Two aqueous solutions

Galvanic Cells

• Galvanic cells convert chemical energy into electrical energy through spontaneous redox reactions.
• They consist of two half-cells with different electrode potentials, facilitating the flow of electrons through an external circuit.
• The anode undergoes oxidation, releasing electrons, while the cathode undergoes reduction, accepting electrons to complete the circuit and generate electrical current.

Galvanic Cells

Electrolytic Cells

• Electrolytic cells use electrical energy to drive non-spontaneous redox reactions.
• They consist of two electrodes submerged in an electrolyte solution.
• The anode is where oxidation occurs, while the cathode is where reduction takes place.

Measurement of Electrode Potential

• Electrode potential is measured using a reference electrode, typically the standard hydrogen electrode (SHE).
• The potential difference between the electrode under study and the reference electrode gives the electrode potential.
• The electrode potential is often expressed in volts (V) and is a measure of the tendency of the electrode to undergo reduction or oxidation reactions.
• Formula for EMF of a cell is given as: Eocell= Ecathode– Eanode

Nernst Equation

• The Nernst equation describes the relationship between the concentration of species involved in a redox reaction and the electrode potential.
• It states that the electrode potential varies logarithmically with the concentration of ions according to the formula Ecell = E0 – [RT/nF] ln Q, 
  • where E is the electrode potential, 
  • E° is the standard electrode potential, 
  • R is the gas constant, 
  • T is the temperature, 
  • n is the number of moles of electrons transferred, 
  • F is Faraday's constant, 
  • Q is the reaction quotient.
• The Nernst equation is fundamental in understanding how changes in concentration affect the redox equilibrium and electrode potential in electrochemical cells.

Nernst equation

\(E = E^0 - \frac{RT}{nF}InQ\)

E = Cell potential under non-standard conditions

E⁰ = Cell potential under conditions

R = universal gas constant

T = Temperature in Kelvin

n = Number of moles of electrons transferred in a balanced redox reaction

F = Faradey constant

Q = Reaction quotient

Equilibrium Constant from Nernst Equation

• The Nernst equation relates the equilibrium constant (K) of a reaction to the cell potential (E) and the standard cell potential (E∘).
• The Nernst equation provides a means to calculate the equilibrium constant from the cell potential, temperature, and standard cell potential.

Electrochemical Cell and Gibbs Energy of the Reaction

• The standard Gibbs energy (∆rG°) of a cell is related to its standard electrode potential (∆rE°) through the equation ∆rG° = -nF∆rE°, where 'n' is the number of moles of electrons transferred, and 'F' is the Faraday constant.
• The standard electrode potential is defined concerning the standard potential of the hydrogen electrode, which is arbitrarily set as zero. 
• It helps determine the spontaneity of the redox reactions occurring in the electrochemical cell.

Electrochemical Cell and Gibbs Energy of the Reaction

Conductance of Electrolytic Solutions

• The conductance of electrolytic solutions varies with the concentration of the electrolyte present in the solution.
• Molar conductivity (Λm) is defined as the conductivity of a solution divided by its concentration (Λm =κ/c). 
• It is a measure of the ions' ability to conduct electricity in the solution.
• While conductivity decreases with decreasing concentration, molar conductivity tends to increase. 
• This trend is more for strong electrolytes compared to weak electrolytes in highly dilute solutions.

Measurement of the Conductivity of Ionic Solutions

• Conductivity of ionic solutions is measured using a conductivity meter or conductivity probe.
• The conductivity of an ionic solution depends on the concentration of ions present in the solution.
• Conductivity also varies with temperature; generally, conductivity increases with increasing temperature due to enhanced ion mobility.

Variation of Conductivity and Molar Conductivity with Concentration

• Strong Electrolytes: Conductivity decreases as concentration decreases, while molar conductivity increases gradually.
• Weak Electrolytes: Conductivity decreases with decreasing concentration, but molar conductivity increases significantly at very low concentrations.
• Kohlrausch's Law: Molar conductivity at infinite dilution is the sum of contributions from ions, illustrating the law of independent migration of ions.

Variation of Molar Conductivity with \(\sqrt{C}\)

Variation of Conductivity and Molar Conductivity with Concentration

Electrolytic Cells and Electrolysis

• Electrolysis is the process occurring in electrolytic cells where electrical energy is used to decompose compounds into their constituent elements.
• Electrolytic cells are essential in various industrial processes like electroplating, extraction of metals, and production of chemicals.

Batteries

• Batteries store and provide electrical energy for various devices and applications, ranging from small gadgets to electric vehicles.
• Batteries typically consist of one or more electrochemical cells, each comprising electrodes, electrolytes, and separators.
• They facilitate the conversion of chemical energy into electrical energy and vice versa.
• Batteries come in various types, including lead-acid, lithium-ion, nickel-cadmium, and alkaline batteries.

Primary Batteries

• Primary batteries are designed for one-time use, meaning they cannot be recharged once depleted.
• They generate electrical energy through irreversible chemical reactions between the electrodes and electrolyte.
• Once the reactants are converted into products, primary batteries cease to produce electricity and become depleted.

Primary Battery

Secondary Batteries

• Secondary batteries, also known as rechargeable batteries, can be recharged and reused multiple times.
• Unlike primary batteries, the cell reaction in secondary batteries can be reversed by passing electricity through the battery, enabling repeated use.
• Common examples of secondary batteries include lead-acid batteries used in automobiles and lithium-ion batteries in electronic devices.

Secondary Battery

Fuel Cells

• Fuel cells are devices that convert chemical energy directly into electrical energy through a chemical reaction. 
• Fuel cells operate through the electrochemical reaction of hydrogen and oxygen to produce electricity, with water and heat as byproducts.
• Common types include proton exchange membrane fuel cells (PEMFCs), solid oxide fuel cells (SOFCs), and molten carbonate fuel cells (MCFCs), each with unique applications and operating temperatures.
• Fuel cells offer a promising pathway to clean and efficient energy production, particularly in transportation and stationary power generation applications.

Fuel Cells

Corrosion

• Corrosion is an electrochemical process involving the degradation of metals due to chemical reactions with the environment, typically oxygen and moisture.
• It leads to the formation of oxides, hydroxides, or salts on the metal surface, weakening its structural integrity and causing material loss.
• Factors such as humidity, temperature, pH, and the presence of impurities accelerate corrosion, emphasizing the electrochemical nature of this destructive process.

Fe²⁺ + 2H₂O → Fe(OH)₂ + 2H⁺

4Fe(OH)₂ + O₂ + x H₂O  → 2 Fe₂O₃ . (x+4) H₂O

Corrosion

Electromotive force (emf) serves as the driving potential for current flow, measured in volts, with the Standard Hydrogen Electrode (SHE) acting as the reference point at E° = 0 V. 

• The Electrochemical Series arranges elements based on their standard electrode potentials, while the standard emf of a cell reflects the difference between cathode and anode potentials. 
• Reduction potential indicates a substance's readiness to accept electrons, which can vary with ion concentration.
• Resistance and resistivity measure the obstruction to current flow, with conductance and specific conductance being their reciprocals. 
• Temperature influences conductance, and dilution affects equivalent and molar conductance. 
• Faraday's Laws of Electrolysis relate mass deposition to the quantity of electricity passed, with the Faraday Constant (F) representing the charge on one mole of electrons. 

There are Some important List Of Top Electrochemistry Questions Asked In CBSE CLASS XII