Summary
Class 12 Chemistry Chapter 2, Electrochemistry, covers the interconversion of chemical and electrical energy, including galvanic cells, the Nernst equation, electrolytic conductance, Faraday's laws of electrolysis, batteries, fuel cells, and corrosion as an electrochemical process.
Electrochemistry (Class 12 Chemistry Part I, Chapter 2) studies how spontaneous redox reactions produce electrical energy in galvanic cells and how electrical energy drives non-spontaneous reactions in electrolytic cells. The chapter derives the Nernst equation relating cell EMF to ion concentrations, connects standard cell potential to Gibbs energy (DrG° = –nFE°cell) and the equilibrium constant, and explains molar conductivity and Kohlrausch's law of independent migration of ions. Practical applications covered include primary and secondary batteries (dry cell, lead storage, nickel-cadmium), hydrogen-oxygen fuel cells, and corrosion of iron as an electrochemical phenomenon.
Key points & formulas
- 01A galvanic cell converts Gibbs energy of a spontaneous redox reaction into electrical work; the Daniell cell (Zn/Cu²⁺) has a standard EMF of 1.1 V.
- 02The Nernst equation E(cell) = E°(cell) – (RT/nF) ln Q relates cell potential to the reaction quotient; at 298 K it simplifies to E(cell) = E°(cell) – (0.059/n) log Q.
- 03Standard cell potential, standard Gibbs energy, and equilibrium constant are interrelated: DrG° = –nFE°(cell) and E°(cell) = (2.303RT/nF) log Kc.
- 04Molar conductivity (Λm = κ/c) increases on dilution for both strong and weak electrolytes; for strong electrolytes Λm = Λ°m – A√c (Debye–Hückel–Onsager equation).
- 05Kohlrausch's law states that limiting molar conductivity equals the sum of individual ionic contributions: Λ°m = ν₊λ°₊ + ν₋λ°₋; it is used to find Λ°m of weak electrolytes.
- 06Corrosion of iron is an electrochemical process in which the metal surface acts as a galvanic cell: iron is oxidised at the anode (E° = –0.44 V) and oxygen is reduced at the cathode (E° = +1.23 V), giving an overall E°(cell) of 1.67 V.
Frequently asked questions
01What is the Nernst equation and how is it used in Class 12 Chemistry Chapter 2?
The Nernst equation is E(cell) = E°(cell) – (RT/nF) ln Q, where R is the gas constant (8.314 J K⁻¹ mol⁻¹), T is temperature in kelvin, n is the number of electrons transferred, F is the Faraday constant (96487 C mol⁻¹), and Q is the reaction quotient. At 298 K it simplifies to E(cell) = E°(cell) – (0.059/n) log Q. It is used to calculate cell EMF at non-standard concentrations and, at equilibrium (E = 0), to derive the equilibrium constant Kc of the cell reaction.
02How are standard cell potential, Gibbs energy, and equilibrium constant related in electrochemistry?
They are related by two equations from the chapter: DrG° = –nFE°(cell) connects standard Gibbs energy to standard cell potential, and DrG° = –RT ln K connects Gibbs energy to the equilibrium constant. Combining these gives E°(cell) = (2.303RT/nF) log Kc, so a positive standard cell potential corresponds to a spontaneous reaction with Kc > 1.
03What is Kohlrausch's law of independent migration of ions?
Kohlrausch's law states that the limiting molar conductivity (Λ°m) of an electrolyte equals the sum of the limiting molar conductivities of its individual ions: Λ°m = ν₊λ°₊ + ν₋λ°₋, where ν₊ and ν₋ are the numbers of cations and anions per formula unit. The law is used to calculate Λ°m for weak electrolytes (like acetic acid) that cannot be determined by direct extrapolation, and to find the degree of dissociation and dissociation constant at a given concentration.
04Is the NCERT Class 12 Chemistry Chapter 2 PDF free to download?
Yes, the NCERT Class 12 Chemistry Part I Chapter 2 (Electrochemistry) PDF is completely free to download on cbseprepmaster.com.
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