Chapter: Electrochemistry – Detailed Notes for NEET/JEE Mains

1. Introduction to Electrochemistry

• Electrochemistry: The branch of chemistry that deals with 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.
• Electrochemical Cells (Voltaic or Galvanic Cells): Devices that convert chemical energy of spontaneous redox reactions into electrical energy.
  • Anode: Electrode where oxidation occurs (negative terminal).
  • Cathode: Electrode where reduction occurs (positive terminal).
  • Salt Bridge: An inverted U-tube containing an inert electrolyte (e.g., KCl, KNO3​, NH4​NO3​) in agar-agar gel. It connects the two half-cells, completes the electrical circuit, and maintains electrical neutrality by allowing the flow of ions.
  • Electrode Potential: The potential difference developed between the electrode and the electrolyte solution.
    • Standard Electrode Potential (E∘): Electrode potential measured under standard conditions (1 M concentration for ions, 1 bar pressure for gases, 298 K temperature).
  • Standard Hydrogen Electrode (SHE): A reference electrode with a standard electrode potential of 0.00 V. It consists of a platinum electrode dipped in 1 M H+ solution, over which hydrogen gas at 1 bar pressure is bubbled.
  • Cell Potential / EMF of the Cell (Ecell​): The potential difference between the two electrodes of a galvanic cell.
    • Ecell​=Ecathode​−Eanode​ (where Ecathode​ and Eanode​ are reduction potentials).
    • For standard conditions: Ecell∘​=Ecathode∘​−Eanode∘​.

2. Nernst Equation

• The Nernst equation relates the electrode potential (or cell potential) to the concentrations of species involved in the electrode reaction (or cell reaction).
• For an electrode reaction: Mn+(aq)+ne−→M(s)
  • E=E∘−nFRT​ln[Mn+(aq)][M(s)]​
  • Since concentration of solid is taken as unity: E=E∘−nFRT​ln[Mn+(aq)]1​
  • At 298 K, converting ln to log10​: E=E∘−n0.0592​log10​[Mn+(aq)]1​
• For a cell reaction: aA+bB→cC+dD
  • Ecell​=Ecell∘​−nFRT​lnQ
  • Where Q is the reaction quotient, Q=[A]a[B]b[C]c[D]d​
  • At 298 K: Ecell​=Ecell∘​−n0.0592​log10​Q
  • n = number of electrons exchanged in the balanced redox reaction.
  • F = Faraday constant (96485 C mol−1).

3. Gibbs Energy and Cell Potential

• For a spontaneous cell reaction, the Gibbs energy change (ΔG) is negative.
• The relationship between ΔG and Ecell​ is:
  • ΔG=−nFEcell​
  • Under standard conditions: ΔG∘=−nFEcell∘​
• Relationship between ΔG∘ and Equilibrium Constant (Kc​):
  • ΔG∘=−RTlnKc​
  • Combining with the standard cell potential: −nFEcell∘​=−RTlnKc​
  • Ecell∘​=nFRT​lnKc​
  • At 298 K: Ecell∘​=n0.0592​log10​Kc​

4. Electrolytic Cells

• Electrolytic Cell: A device that uses electrical energy to drive a non-spontaneous chemical reaction (electrolysis).
• Electrolysis: The process of chemical decomposition of an electrolyte by the passage of electric current.
  • Anode: Electrode where oxidation occurs (positive terminal).
  • Cathode: Electrode where reduction occurs (negative terminal).
• Faraday’s Laws of Electrolysis:
  • First Law: The mass of any substance deposited or liberated at any electrode is directly proportional to the quantity of electricity passed through the electrolyte.
    • W∝Q⇒W=ZQ=ZIt
    • W = mass deposited, Q = quantity of electricity (Coulombs), I = current (Amperes), t = time (seconds).
    • Z = Electrochemical equivalent (Z=Molar mass/(n×F)).
  • Second Law: When the same quantity of electricity is passed through solutions of different electrolytes connected in series, the masses of the substances deposited or liberated at the electrodes are directly proportional to their equivalent masses (or molar masses divided by number of electrons involved).
    • W2​W1​​=E2​E1​​ (where E is equivalent mass).

5. Conductance of Electrolytic Solutions

• Resistance (R): Obstruction to the flow of current. Unit: Ohm (Ω).
  • R=ρAl​ (where ρ is resistivity, l is length, A is area of cross-section).
• Resistivity (ρ / Specific Resistance): Resistance of a conductor of unit length and unit area of cross-section. Unit: Ohm-meter (Ω m) or Ohm-cm (Ω cm).
• Conductance (G): Reciprocal of resistance. Unit: Siemens (S) or Ω−1 or Mho.
  • G=1/R
• Conductivity (κ / Specific Conductance): Reciprocal of resistivity. Unit: Siemens per meter (S m−1) or S cm−1.
  • κ=1/ρ=G×(l/A)
  • The term (l/A) is called the Cell Constant (G∗). So, κ=G×G∗.
• Molar Conductivity (Λm​): Conductivity of a solution containing one mole of electrolyte placed between two electrodes of unit area of cross-section and unit distance apart.
  • Λm​=Molarity (M)κ×1000​ (Units: S cm2 mol−1)
  • Λm​ increases with dilution for both strong and weak electrolytes. For strong electrolytes, it’s due to decreased interionic attraction. For weak electrolytes, it’s primarily due to increased degree of dissociation.

6. Kohlrausch’s Law of Independent Migration of Ions

• Kohlrausch’s Law: At infinite dilution, when the dissociation is complete, each ion makes a definite contribution to the molar conductivity of the electrolyte, irrespective of the nature of the other ion with which it is associated.
  • Λm∘​=λc∘​+λa∘​ (where Λm∘​ is molar conductivity at infinite dilution, λc∘​ and λa∘​ are limiting molar conductivities of cation and anion respectively).
• Applications:
  • Calculation of molar conductivity of weak electrolytes at infinite dilution (which cannot be found by extrapolation).
  • Calculation of the degree of dissociation (α) of weak electrolytes: α=Λm∘​Λm​​
  • Calculation of dissociation constant (Ka​) for weak electrolytes.

7. Batteries (Electrochemical Cells in Practical Use)

• Primary Batteries: Non-rechargeable (cannot be reused). The reaction goes to completion.
  • Dry Cell (Leclanché Cell):
    • Anode: Zinc container
    • Cathode: Carbon (graphite) rod surrounded by powdered MnO2​ and carbon.
    • Electrolyte: Paste of NH4​Cl and ZnCl2​.
    • Cell Potential: Approximately 1.5 V.
  • Mercury Cell:
    • Anode: Zinc-mercury amalgam
    • Cathode: Paste of HgO and carbon.
    • Electrolyte: Paste of KOH and ZnO.
    • Cell Potential: Approximately 1.35 V (constant during its life due to no net ion change in solution).
• Secondary Batteries: Rechargeable (can be reused).
  • Lead-Acid Battery:
    • Anode: Lead (Pb)
    • Cathode: Lead dioxide (PbO2​)
    • Electrolyte: 38% H2​SO4​ solution.
    • Reactions (Discharging):
      • Anode: Pb(s)+SO42−​(aq)→PbSO4​(s)+2e−
      • Cathode: PbO2​(s)+SO42−​(aq)+4H+(aq)+2e−→PbSO4​(s)+2H2​O(l)
      • Overall: Pb(s)+PbO2​(s)+2H2​SO4​(aq)→2PbSO4​(s)+2H2​O(l)
    • Charging: Reactions are reversed by applying external current. Density of H2​SO4​ increases.
  • Nickel-Cadmium Cell (NiCad):
    • Anode: Cadmium (Cd)
    • Cathode: Nickel dioxide (NiO2​)
    • Electrolyte: KOH solution.
    • Longer life than lead-acid but more expensive.
• Fuel Cells: Galvanic cells that convert the chemical energy from the combustion of fuels (like H2​, CH4​, CH3​OH) directly into electrical energy.
  • Hydrogen-Oxygen Fuel Cell:
    • Fuel: H2​
    • Oxidant: O2​
    • Electrolyte: Concentrated KOH or NaOH solution.
    • Anode: 2H2​(g)+4OH−(aq)→4H2​O(l)+4e−
    • Cathode: O2​(g)+2H2​O(l)+4e−→4OH−(aq)
    • Overall: 2H2​(g)+O2​(g)→2H2​O(l)
    • Advantages: High efficiency, pollution-free, continuous energy supply as long as reactants are supplied.

8. Corrosion

• Corrosion: The process of slow deterioration of a metal by the action of air, moisture, or chemicals on its surface. It is an electrochemical process.
• Rusting of Iron:
  • At a specific spot on the surface of iron, oxidation occurs.
    • Anode: Fe(s)→Fe2+(aq)+2e−
  • Electrons released move through the metal to another spot on the metal surface (often where oxygen is available).
  • Cathode: O2​(g)+4H+(aq)+4e−→2H2​O(l) (The H+ ions come from H2​CO3​ formed by CO2​ in air dissolving in water).
  • Overall redox reaction: 2Fe(s)+O2​(g)+4H+(aq)→2Fe2+(aq)+2H2​O(l)
  • The Fe2+ ions are further oxidized by atmospheric oxygen to Fe3+ ions, which then form hydrated ferric oxide (Fe2​O3​⋅xH2​O), commonly known as rust.
    • 4Fe2+(aq)+O2​(g)+4H2​O(l)→2Fe2​O3​(s)+8H+(aq)
    • Fe2​O3​(s)+xH2​O(l)→Fe2​O3​⋅xH2​O(s) (Rust)
• Prevention of Corrosion:
  • Barrier Protection: Painting, oiling, greasing, plastic coating.
  • Sacrificial Protection (Cathodic Protection): Coating iron with a more electropositive metal (e.g., Zn in galvanization) which gets oxidized preferentially.
  • Electrochemical Protection: Connecting the metal to be protected to a more active metal.
  • Anti-rust solutions: Phosphates or chromates form a protective layer.