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Electrochemistry Class 12 Revision Notes: The electrochemistry chapter in CBSE Class 12 is an important part of the class 12 chemistry syllabus not only because the chapter carries the highest weightage for the board exams but also because it covers the fundamentals of electrochemistry, which is a vast field with applications in many areas of science and technology. Therefore, students must have all the concepts clear and well-learned. The short notes presented in this article will prove to be very helpful to recall all important concepts from the chapter. These notes have been prepared according to the new CBSE syllabus for Class 12 Chemistry. Subject experts have created and reviewed the revision notes for Class 12 Electrochemistry. Therefore, these notes are best to revise the chapter in a few minutes to save your time for the practice of important questions and previous years’ question papers. You can download the notes in PDF as well.
Also Read: CBSE Class 12 Chemistry Syllabus 2023-24
Revision Notes for CBSE Class 12 Chemistry Chapter 2, Electrochemistry
Electrochemical cell
It is a cell that converts the chemical energy of a spontaneous redox reaction into electrical energy.
In an electrochemical cell:
- The half-cell in which oxidation takes place is known as anode (negatively charged).
- The half-cell in which reduction takes place is known as cathode (positively charged).
Salt bridge: It is a U shaped tube containing an inert electrolyte in agar-agar and gelatine. It allows the flow of electric current by completing the electrical circuit.
Representation of an Electrochemical Cell
- Anode represents the oxidation half-cell and is is written on the left
It is written as Metal/Metal ion (Concentration)
- Cathode represents the reduction half-cell and is written on the right side.
It is written as Metal ion (Concentration)/Metal
- Salt bridge is indicated by placing two vertical lines between the anode and the
cathode.
Cell Potential: It is the potential difference between the two electrodes of a galvanic cell. It is measured in volts.
Ecell = Ecathode – Eanode
Electromotive Force (EMF): Cell potential is called the EMF when no current is drawn through the cell.
Electrode Potential: It is the potential difference between the electrode and the electrolyte.
- Electrode potential increases with increase in the concentration of the electrolyte and decrease in temperature.
- At equilibrium, Ecell = 0
Standard Electrode Potential: When the concentration of all the species involved in a half
cell is unity, then the electrode potential is known as standard electrode potential. It is
denoted as E0.
Standard Hydrogen Electrode (SHE): It is taken as a reference electrode and it is assigned a zero potential at all temperatures.
It is represented by Pt(s)|H2(g)|H+(aq),
- The standard hydrogen electrode can act as anode as well as
As an anode: H2(g) → 2H(aq) + 2e−
As a cathode: H(aq) + 2e− → H2(g)
Electrochemical Series:
In the electrochemical series, various elements are arranged as per their standard reduction
potential values.
- A substance with higher reduction potential value means that it has a higher tendency to get So, it acts as a good oxidising agent.
- A substance with lower reduction potential value means that it has a higher tendency to get
- So, it acts as a good reducing agent.
- The electrode with higher reduction potential acts as a cathode while the electrode with a
lower reduction potential acts as an anode.
Nernst Equation
It represents the relationship between standard electrode potential E0 and electrode potential E as follows:
Gibbs Energy: Work done by an electrochemical cell is equal to the decrease in Gibbs energy
ΔrG = – nFEcell
If the concentration of all the reacting species is unity, then Ecell = E0cell
⇒ ΔrG0 =– nFE0cell
Resistance: The obstruction offered by a conducting material to the flow of electricity which is called resistance. It is denoted by R and is measured in ohm (Ω).
- The resistance of a conductor is directly proportional to its length l and inversely proportional
to its area of cross section A.
R = ρ (l/A)
Ρ is the constant of proportionality and is called resistivity (specific resistance). Its SI unit is ohm metre (Ωm).
Conductance: The inverse of resistance, R, is called conductance and is represented as G.
G = 1/R = A/ρl = κ(A/l)
The SI unit of conductance is siemens, represented by the symbol ‘S’. It is equal to ohm–1 (also known as mho) or Ω–1.
Conductivity: The inverse of resistivity is called conductivity or specific conductance. It is represented by κ.
The SI unit of conductivity is S m–1 but it is also expressed in S cm–1.
Conductivity always decreases with decrease in concentration both, for weak and strong electrolytes.
Metallic or Electronic Conductance:
Electrical conductance through metals is called metallic or electronic conductance and is due to the movement of electrons. The electronic conductance depends on
(i) the nature and structure of the metal
(ii) the number of valence electrons per atom
(iii) temperature (it decreases with increase of temperature)
Electrolytic or Ionic Conductance:
The conductance of electricity by ions present in the solutions is called electrolytic or ionic conductance. It depends on:
(i) the nature of the electrolyte added
(ii) size of the ions produced and their solvation
(iii) the nature of the solvent and its viscosity
(iv) concentration of the electrolyte
(v) temperature (it increases with the increase of temperature)
Conductivity Cell
A conductivity cell consists of two platinum electrodes coated with platinum black. These have area of cross section equal to ‘A’ and are separated by distance ‘l’. The resistance of a column of solution (confined between two electrodes) is then given by the equation:
R = ρ (l/A) = l/(κA)
Cell Constant (G*): The quantity l/A is called cell constant.
Molar Conductivity (Λm)
Molar conductivity of a solution is defined as the conducting power of all the ions produced by dissolving 1 mole of an electrolyte in the solution.
Molar conductivity, Λm = κ/c
Or kΛm = (κ x 1000)/M
Where M is the molarity of the solution.
The unit for Λm is Scm2mol–1.
Molar conductivity increases with decrease in concentration.
Limiting Molar Conductivity (Λ0m): When concentration approaches zero, the molar conductivity is termed as limiting molar conductivity and is represented by the symbol Λ0m.
Kohlrausch Law of Independent Migration of Ions:
The law states that limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and cation of the electrolyte.
If an electrolyte on dissociation gives ν+ cations and ν– anions then its limiting molar conductivity is given by: Λ0m = ν+λ0+ + ν– λ0–
Degree of dissociation: It is the ratio of molar conductivity Λcm at a specific concentration ‘c’ to the limiting molar conductivity, Λ0m. It is represented as α.
α = Λcm/Λ0m
Electrolysis: It is a process of decomposition of an electrolyte by the passage of electricity through its aqueous solution or molten state.
Faraday’s Laws of Electrolysis:
(i) Faraday’s first law of electrolysis: The amount of chemical reaction which occurs at any electrode during electrolysis is proportional to the quantity of electricity passed through the electrolyte.
(ii) Faraday’s second law of electrolysis: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their chemical equivalent weights (Atomic Mass of Metal ÷ Number of electrons required to reduce the cation).
Faraday constant: It is equal to charge on 1 mole of electrons. It is taken as 96500 C/mol (approx.)
Products of electrolysis:
The products of electrolysis depend upon
(i)The nature of electrolyte being electrolyzed and the nature of electrodes. If the electrode is inert (e.g., platinum or gold), it does not participate in the chemical reaction and acts only as source or sink for electrons. If the electrode is reactive then it will take part in chemical reaction and products will be different as compared to inert electrodes.
(ii) The different oxidising and reducing species present in the electrolytic cell and their standard electrode potentials.
Primary Cells
A primary cell is a cell in which electrical energy is produced by the reaction occurring in the cell.
- Examples: Daniel cell, dry cell, mercury cell.
- It cannot be recharged.
- Uses: These are commonly used in transistors and watches.
(i) Dry Cell: The cell consists of a zinc container that also acts as anode and the cathode is a carbon (graphite) rod surrounded by powdered manganese dioxide and carbon. The space between the electrodes is filled by a moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2).
Anode: Zn(s) → Zn2+ + 2e–
Cathode: MnO2+ NH4+ + 2e– → 2MnO(OH) + NH3
(ii) Mercury Cell: It consists of zinc–mercury amalgam as anode and a paste of HgO and carbon as the cathode. The electrolyte is a paste of KOH and ZnO.
Anode: Zn (Hg) + 2OH– → ZnO(s) + H2O + 2e–
Cathode: HgO + H2O + 2e– → Hg (l) + 2OH–
Secondary Cell (lead storage battery)
It consists of a lead anode and a grid of lead packed with lead dioxide (PbO2) as cathode. A 38% solution of Sulphuric acid is used as an electrolyte.
- It can be recharged.
- Uses: These are commonly used in automobiles and invertors.
The cell reactions involved are:
Anode: Pb(s) + SO42–(aq) → PbSO4(s) + 2e–
Cathode: PbO2 (s) + SO42–(aq) + 4H+ (aq) + 2e– → PbSO4(s) + 2H2O (l)
Overall cell reaction is:
Pb(s) + PbO2 (s) + 2H2SO4 (aq) → 2PbSO4 (s) + 2H2O(l)
Fuel Cells: These are the galvanic cells that are designed to convert the energy of combustion of fuels like hydrogen, methane, methanol, etc. directly into electrical energy.
Fuel cells are pollution free.
The electrode reactions for fuel cell using H2 and O2 are given below:
Cathode: O2 (g) + 2H2O (l) + 4e– → 4OH–(aq)
Anode: 2H2 (g) + 4OH–(aq) 4H2O(l) + 4e–
Overall reaction is:
2H2(g) + O2(g) → 2H2O (l)
Corrosion: It is the process of gradual coating of the surfaces of metallic objects with oxides or other salts of the metal.
Examples: Rusting of iron, tarnishing of silver, development of green coating on copper and bronze.
Anode: Fe(s) → Fe2+ + 2e– E0(Fe2+/Fe) = – 0.44 V
Cathode: O2(g) + 4H+(aq) + 4e– → 2H2O (l) E0H+/O2/H2O = 1.23 V
The overall reaction is:
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l) E0cell =1.67 V
Prevention of corrosion:
- By covering the surface of the metal with paint or some chemicals (e.g. bisphenol).
- By covering the surface of the metal with of the o cover the surface by other metals like Sn, Zn, etc., that are inert or react to save the object.
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