Redox Reactions And Electrochemistry Revision Notes for Chemistry NEET

Redox Reactions and Electrochemistry covers fundamental concepts that are essential for understanding chemical changes and energy conversions involving electrons. The chapter begins by exploring how oxidation and reduction occur at the electronic level, emphasizing the gain and loss of electrons and how this is central to redox reactions in chemistry. These concepts are important for interpreting various chemical equations and identifying the direction of electron flow in reactions.

Electronic Concepts of Oxidation and Reduction Oxidation is defined as the loss of electrons, while reduction is the gain of electrons. In any redox reaction, both processes occur together, as one species loses electrons (is oxidized) and another gains electrons (is reduced). The substance that donates electrons is called the reducing agent, and the one that accepts electrons is the oxidizing agent. For example, in the reaction:

• Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Zinc gets oxidized by losing electrons, and copper ions get reduced by gaining electrons.

Oxidation Number and Rules for Assigning Oxidation Number The oxidation number (or state) is the apparent charge an atom has in a compound if all bonds are assumed to be ionic. To assign oxidation numbers, follow these basic rules:

• The oxidation number of an element in its free state is 0.
• For monoatomic ions, it equals the ionic charge.
• Oxygen is usually –2, except in peroxides where it is –1.
• Hydrogen is +1 with non-metals and –1 with metals.
• The sum of oxidation numbers in a neutral compound is zero, in polyatomic ions it equals the ion charge.

For example, in H₂SO₄, H is +1, O is –2; thus S is +6.

Redox Reactions Redox reactions involve the simultaneous oxidation and reduction of species. These reactions can be split into two half-reactions: one for oxidation and the other for reduction. Balancing redox reactions is crucial for correct stoichiometry. In acidic or basic medium, the ion-electron method (also called the half-reaction method) is commonly used for balancing:

1. Write separate half-reactions.
2. Balance atoms other than O and H.
3. Balance O using H₂O, then H using H⁺ (in acid) or OH^– (in base).
4. Balance charges by adding electrons.
5. Equalize electron transfer and add the half-reactions.

Electrolytic and Metallic Conduction Conduction in metals results from the flow of electrons, while in electrolytic solutions, it arises from the movement of ions. In electrolysis, an external voltage is used to drive a non-spontaneous chemical reaction. The electrolyte (solution or molten salt) decomposes, producing products at the electrodes. Metallic conduction is not affected by the nature of the electrolyte or solvent, while electrolytic conduction depends on the mobility and number of ions present in the solution.

Conductance in Electrolytic Solutions and Molar Conductivities The ability of a solution to conduct electricity is called conductance. It is measured using the cell constant and observed conductance. Two important types of conductance are:

• Specific conductance ($\kappa$): Conductance of one cm cube of solution.
• Molar conductivity ($\Lambda_m$): Conductance of all ions produced from one mole of electrolyte in solution.

Molar conductivity increases with dilution due to an increase in the degree of dissociation, especially for weak electrolytes.

Kohlrausch’s Law and its Applications Kohlrausch’s Law of Independent Migration of Ions states that at infinite dilution, each ion contributes to the molar conductivity independently. It is mathematically represented as: $\Lambda_m^\circ = \lambda^0_+ + \lambda^0_-$ This law is useful for:

• Calculating molar conductivity at infinite dilution for weak electrolytes.
• Determining the degree of ionization of weak electrolytes.
• Predicting solubility of sparingly soluble salts.

Electrochemical Cells: Electrolytic and Galvanic Cells Electrochemical cells convert chemical energy to electrical energy or vice versa. A Galvanic cell (like the Daniel cell) generates electricity from a spontaneous redox reaction, while an electrolytic cell consumes electricity to drive a non-spontaneous reaction. Table: Comparison between Galvanic and Electrolytic Cells

Feature	Galvanic Cell	Electrolytic Cell
Energy Conversion	Chemical to Electrical	Electrical to Chemical
Spontaneity	Spontaneous Reaction	Non-spontaneous Reaction
Anode Polarity	Negative	Positive
Electron Flow	Anode to Cathode	Cathode to Anode

Types of Electrodes and Electrode Potentials Electrodes can be classified as metal-metal ion, gas electrode, and redox electrodes. The potential developed at the interface of an electrode and its solution is the electrode potential. Standard Electrode Potential (SEP), denoted as $E^\circ$, is measured under standard conditions (1 M, 1 atm, 25°C). The SEP of the Standard Hydrogen Electrode (SHE) is defined as 0 V. Electrochemical series arranges elements in increasing order of their SEPs.

Half-cell and Cell Reactions, Emf of Galvanic Cell A Galvanic cell consists of two half-cells, each having an electrode and its ion solution. The cell potential (emf) is the difference between the electrode potentials of the two electrodes: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ Cell notation is written as: Anode | Anode Solution || Cathode Solution | Cathode The overall cell reaction is obtained by combining the oxidation half-reaction at the anode and the reduction half-reaction at the cathode.

Nernst Equation and Its Applications The Nernst equation relates the cell potential under non-standard conditions to the standard electrode potential: $E_{cell} = E^\circ_{cell} - \frac{0.0591}{n} \log Q$ where n is the number of electrons exchanged, and Q is the reaction quotient. This equation allows the calculation of cell potential at any concentration. It also helps in determining equilibrium constant and feasibility of cell reactions.

Relationship between Cell Potential and Gibbs Energy Change The free energy change (ΔG) for a cell reaction is directly related to the emf (E) of the cell by the formula: $\Delta G = -nFE_{cell}$ where F is the Faraday constant (96,500 C/mol e^–). A negative ΔG indicates a spontaneous reaction.

Types of Electrochemical Cells: Dry Cell, Lead Accumulator, Fuel Cells A dry cell (Leclanché cell) is a common non-rechargeable battery. The lead accumulator is a rechargeable battery, widely used in vehicles – it involves reversible chemical reactions. Fuel cells convert the chemical energy of a continuously supplied fuel (like H₂ and O₂) directly into electrical energy, and are used for clean energy applications.

NEET Chemistry Notes – Redox Reactions And Electrochemistry: Key Points for Quick Revision

These NEET Chemistry notes on Redox Reactions and Electrochemistry cover all crucial subtopics such as oxidation numbers, balancing redox equations, and cell potentials. Each concept is explained stepwise so students can revise quickly and recall facts for exams. Important formulas and rules are highlighted for faster memorization.

Using these revision notes, you can confidently solve NEET questions related to electrolytic conduction, Nernst equation applications, and real-world cells like dry cells and fuel cells. Practice with summarized tables, lists, and examples to strengthen your basics and apply concepts efficiently during the exam.