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NIOS Chemistry 313 - Part 3 आपका स्वागत है हमारे Study Platform में आज हम आपके NIOS Chemistry 313 segment का 3rd part करने वाले हैं, जो आपके exam में अकेले 12 marks का weightage रखता है। यानि अगर आप ये सभी Question-answers cover कर लेते हैं तो आपकी पूरे 12 marks की तैयारी complete हो जाएगी। इसी तरह से हमने आपके पूरे syllabus को 7 parts में divide किया है। जो अगर आप पूरी तरह से read कर लेते हैं तो आप पूरे 100 marks के लिए eligible हो जाएंगे। बाकी के 6 parts, notes, mock test और प्रैक्टिकल files के लिए आप हमारे Guaranteed Pass Course को join कर सकते हैं। Course Join करें Redox and Electrochemical Concepts Redox and Electrochemical Concepts: Easy Explanation Q.1: Understand oxidation and reduction in terms of electron transfer concept Oxidation: Oxidation is the process where an atom or ion loses electrons. Example: In the reaction where magnesium reacts with oxygen (Mg → Mg²⁺ + 2e⁻), magnesium is oxidized because it loses two electrons. Reduction: Reduction is the process where an atom or ion gains electrons. Example: In the reaction where copper ions react with electrons (Cu²⁺ + 2e⁻ → Cu), copper is reduced because it gains two electrons. Q.2: Calculate oxidation number (ON) of an atom in a molecule or ion Oxidation Number (ON): It is the charge an atom would have if the compound was purely ionic. Some basic rules: The oxidation number of an element in its pure form (e.g., O₂, N₂) is zero. The oxidation number of oxygen is usually -2, except in peroxides where it's -1. The oxidation number of hydrogen is +1 when bonded with non-metals and -1 when bonded with metals. Example: In H₂O, oxygen is -2, and hydrogen is +1. The sum of oxidation numbers should be zero. (2 × +1) + (-2) = 0. Q.3: Balance the chemical equation for redox reactions To balance a redox reaction, we use the half-reaction method: Separate the oxidation and reduction half-reactions. Balance all elements except oxygen and hydrogen. Balance oxygen by adding H₂O molecules. Balance hydrogen by adding H⁺ ions. Balance the charges by adding electrons. Q.4: Explain electrolytic conduction, conductance, and molar conductivity Electrolytic conduction: The process by which electricity is conducted through a solution of an electrolyte. This happens because ions are free to move and carry charge. Conductance (G): The ability of a solution to conduct electricity. It depends on the concentration of ions and their mobility. Molar conductivity (Λ): It is the conductivity of a solution divided by its molar concentration. It shows how well a single ion can conduct electricity in a solution. Q.5: Describe the effect of dilution on conductivity and molar conductivity of an electrolyte Conductivity decreases with dilution: As we dilute the solution, the number of ions per unit volume decreases, which reduces the ability to conduct electricity. Molar conductivity increases with dilution: When diluted, the ions become more spread out, reducing ion-ion interactions, so each ion can move more freely, increasing its molar conductivity. Q.6: Differentiate between electrolytic and Galvanic cell Electrolytic Cell: A cell that uses electrical energy to drive a non-spontaneous reaction. It requires an external power source. Galvanic Cell: A cell that generates electrical energy from a spontaneous chemical reaction (e.g., batteries). Q.7: State Faraday's laws of electrolysis First Law: The amount of substance deposited or dissolved at an electrode is directly proportional to the quantity of electricity passed. Formula: m = Z × I × t (where m = mass of substance, Z = electrochemical equivalent, I = current, t = time) Second Law: The amounts of different substances deposited or dissolved by the same quantity of electricity are proportional to their equivalent masses. Q.8: Predict and justify the products of electrolysis of some common electrolytes Example 1: Electrolysis of water (H₂O): At the cathode: reduction of H₂O → H₂ (hydrogen gas). At the anode: oxidation of H₂O → O₂ (oxygen gas). Example 2: Electrolysis of NaCl: At the cathode: reduction of Na⁺ → Na (sodium metal). At the anode: oxidation of Cl⁻ → Cl₂ (chlorine gas). Q.9: State standard electrode potential and use it for calculation of standard electrode potential of a cell Standard Electrode Potential (E°): is the voltage associated with a half-reaction at standard conditions (25°C, 1 M concentration, 1 atm pressure). The cell potential of a galvanic cell is calculated by the difference between the reduction potentials of the two electrodes: Formula: Ecell = Ecathode - Eanode Q.10: Explain Standard Hydrogen Electrode (SHE) SHE is a reference electrode used to measure electrode potentials. It is based on the half-reaction: 2H⁺ + 2e⁻ → H₂ Its potential is defined as 0 volts under standard conditions. Q.11: Describe electrochemical series and its application The electrochemical series is a list of standard electrode potentials arranged in order of increasing reduction potential. Application: It helps in predicting the direction of electron flow in a redox reaction and determining which substance will be oxidized or reduced. Q.12: State the effect of concentration on electrode potential (Nernst equation) Nernst Equation: It relates the concentration of ions to the electrode potential. E = E° - (0.0591 / n) log ([Red] / [Ox]) As the concentration of reactants increases, the potential becomes more positive, favoring reduction. As the concentration of products increases, the potential becomes more negative, favoring oxidation. Q.13: Solve numericals based on Nernst equation You can use the Nernst equation to calculate the potential for any redox half-reaction by substituting the values for concentrations of the reactants and products. Example: For the reaction Cu²⁺ + 2e⁻ → Cu, if the concentration of Cu²⁺ is 0.1 M and Cu is solid, then: Formula: E = E° - (0.0591 / n) log ([Cu²⁺] / [Cu]) By substituting values, you can calculate the potential. Q.14: Find relationship between emf and Gibbs energy change The relationship between emf (electromotive force) and Gibbs free energy change (ΔG) is given by: ΔG = -nFEcell Where n is the number of electrons, F is Faraday’s constant, and Ecell is the cell potential.

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