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https://docs.google.com/document/d/1ppOe2j5BtWgEawsgJsBqHr-RM2U8AXvQ/edit?usp=sharing&ouid=109474854956598892099&rtpof=true&sd=true Redox Reactions and Electrochemistry Electronic concepts of oxidation and reduction, redox reactions, oxidation number, rules for a ssigning oxidation number, balancing of redox reactions. Eectrolytic and metallic conduction, conductance in electrolytic solutions, specific and molar conductivities and their variation with concentration: Kohlrausch’s law and its applications. Electrochemical cells Electrolytic and Galvanic cells, different types of electrodes, electrode potentials including standard electrode potential, half - cell and cell reactions, emf of a Galvanic cell and its measurement; Nernst equation and its applications; Relationship between cell potential and Gibbs’ energy change; Dry cell and lead accumulator; Fuel cells; Corrosion and its prevention. CONCEPT OF OXIDATION & REDUCTION Oxidation Reduction Loss of electron Loss of hydrogen Gain of oxygen Increase in oxidation number Gain of electron Gain of hydrogen Loss of oxygen Decrease in oxidation number OXIDATION NUMBER Oxidation number is defined as the change (real or imaginary) which an atom appears to have undergone when it is present in combination. There are certain rules laid down in order to determine the oxidation number. O.N. of an atom in free elements is zero. O.N. of oxygen is -2 while in peroxides it is -1, in OF2 it is +2. O.N. of hydrogen is +1 while in metal hydrides it is -1. O.N. of an ion is equal to the electrical charge present on it. O.N. of group IA elements is +1 and that of group IIA elements is +2 For complex ions, the algebraic sum of oxidation numbers of all the atoms is equal to the net charge on the ion. In case of neutral molecules the algebraic sum of the oxidation number of all the atoms present in the molecules is zero. C H A P T E R CHAPTER INCLUDES Concept of Oxidation-Reduction Oxidation Number Oxidising and Reducing agent Balancing Redox Reactions Electrolytic cells & Galvanic cells Standard electrode potential Electrochemical series emf of Galvanic cell Gibb's energy change and cell potential Electrolysis Electrolytic conduction Molar conductivity Kohlrausch's law and its application TEACHING CARE Online Live Classes https://www.teachingcare.com/ +91-9811000616 Increase in oxidation number of an element in a reaction is known as oxidation while decrease in oxidation number of an element in a reaction is known as reduction. Besides +ve and -ve values, fractional values of oxidation number are also possible Ex. Na2S4O6 2×1 + 4x + 6(-2) = 0 x = 2.5 Ex. Fe3O4 3×x +4(-2) =0 x = 8 = 2 2 3 3 Ex. N3H 3x +1 =0 x = − 1 3 OXIDISING AND REDUCING AGENT Oxidising agent or oxidant The substance which oxidises others and itself get reduced. Reducing agent or reductant The substance which reduces others and itself get oxidised. Reduction Zn(s) + Cu+2(aq) e.g., Oxidation Zn+2(aq) + Cu In the above redox reaction Zn is acting as reducing agent and Cu+2 as oxdising agent. BALANCING REDOX REACTIONS It occurs by two methods, i.e. Oxidation number method Ion electron method Oxidation Number Method Ion Electron Method Write the skeleton equation representing chemical change. Assign oxidation numbers to find out which atoms are undergoing oxidation and reduction write separate equation for the atoms undergoing oxidation and reduction Find the change in oxidation number in each equation. Make the change equal in both the equation by multiplying with suitable integers. Add both the equations. Write down the redox reaction in ionic form. Split the redox reaction into two half reactions, one for oxidation and other for reduction Balance the number of atoms ions undergoing reduction and oxidation. 4. Now balance those elements which are not undergoing oxidation or reduction except H and O. Then balancing H and O with the help of H2O. Now add electrons on that side of reaction where they are deficient to equalise the charge on both sides. Multiply if required by suitable number to balance electron on both side of reaction. Now add both the half reaction and now balance atoms not undergoing reduction and oxidation (except H and O). Now balance H and O with help of water. Balancing of H2O can also be done as per the medium given Acidic Medium : Add H2O on that side of reaction where oxygen are deficient and double number of H+ on opposite side of reaction. Basic Medium : Add H2O on that side of reaction where oxygen are excess and double number of OH– on opposite side of reaction. ELECTROLYTIC CELLS & GALVANIC CELLS: The chemical changes which involve the flow of the electric current are called electrochemical changes. These are broadly of two types : Electrolytic Cells : The changes in which electrical energy causes chemical reaction to occur. This phenomenon is called electrolysis and the devices or cells used to carry out electrolysis are called electrolytic cells. These reactions are non-spontaneous and are forced to occur by the passage of electricity. Galvanic Cells : “A device employed to convert the chemical energy of a redox reaction into electrical energy” is called an electrochemical cell or simple chemical cell. The most common example is Daniel cell. (–) V e– flow current (+) Zinc electrode Salt bridge Copper electrode ZnSO4 solution Half-cell Inert electrolyte with agar-agar DANIEL CELL Ions Half-cell CuSO4 solution The net reaction is the sum of two half-cell reactions. At anode : Zn(s) ⎯⎯→ Zn+2(aq) + 2e− At cathode : Cu+2(aq) + 2e− ⎯⎯→ Cu(s) Net reaction : Zn(s) + Cu+2 ⎯⎯→ Zn+2(aq) + Cu(s) Schematic Representation : Zn | Zn+2 || Cu+2 | Cu STANDARD ELECTRODE POTENTIAL Standard Electrode Potential : The potential difference developed between metal electrode and the solution of its ions of unit molarity (1 M) at 298 K is called standard electrode potential. Standard Hydrogen Electrode – (SHE) SHE Half Reaction Electrode Potential H2 → 2H+ + 2e– 0.0 V (anode) 2H+ + 2e– → H2 0.0 V (cathode) With the help of SHE - the SRP values of all the electrodes has been determined and are placed in electrochemical series. ELECTROCHEMICAL SERIES Table in which the standard reduction potentials of various electrodes have been arranged in the increasing order is called electrochemical series or activity series or electromotive series. Electrode Electrode reaction E°(volts) Li+/Li Li+ + e– = Li –3.045 (Lowest) Zn2+/Zn Zn2+ + 2e— = Zn –0.763 H+/H2, Pt 2H+ + 2e— = H2 0.0 Standard Cu2+/Cu Cu2+ + 2e— = Cu +0.334 F2/F–, Pt F2 + 2e— = 2F— +2.887 (Highest) Applications of Electrochemical Series : The important applications of electrochemical series are: Relative strength of oxidising and reducing agents. In the series, the elements are arranged in the increasing order of reduction potentials or decreasing order of oxidation potential. Therefore, the elements at the top are good reducing agents while those at the bottom are good oxidising agents. Calculating e.m.f. of the cell : The e.m.f. of the cell can be determined by knowing the standard electrode potentials from the series as: º cell º (right) º (left) (If standard reduction potential are given) Predicting feasibility of a redox reaction. In general, a redox reaction is feasible only if the species which has higher reduction potential is reduced i.e., accepts the electrons and the species which has lower reduction potential is oxidised i.e. loses the electrons. Otherwise, a redox reaction is not feasible. In other words, the species to release electrons must have lesser reduction potential as compared to the species which is to accept electrons. To predict whether a metal can liberate hydrogen from acid or not only the metals which have negative reduction potentials, can liberate hydrogen from the acids. EMF OF GALVANIC CELL In galvanic cells, current is generated as a result of a spontaneous chemical reaction that occur in the cell. The main characteristics of galvanic cell are given below Cathode Anode (i) Sign Positive Negative (ii) Half reaction Reduction Oxidation (iii) Direction of electron movement Into the cell Out of the cell E.M.F is potential difference between the two electrodes of the cell when no current is flowing in the circuit (i.e., in an open circuit). It is the maximum voltage obtainable from the cell. Eºcell = EºR.P.(R.H.S) – EºR.P. (L.H.S.) = EºO.P. (L.H.S.) – EºO.P. (R.H.S.) = EºO.P. (L.H.S.) + EºR.P. (R.H.S.) Similarly Ecell = ER.P.(R.H.S.) – ER.P. (L.H.S.) ER.H.S. or EL.H.S. determined by Nernst equation Nernst equation It is a relation between electrode potential, emf of cell, temperature and concentration of ions in solution represented by the equation E = Eo RT ln [P] cell cell nF [R] where Ecell = EMF of cell at new concentration Eºcell = Standard emf of cell (1M, 25ºC and 1 atm) R = Gas constant = 8.314 JK–1 mol–1. T = Absolute temperature n = Number of electrons F = Faraday constant = 96500 C ln = 2.303 log10 [P] = Concentration of product in mol L–1 [R] = Concentration of reactants in mol L–1 Putting the values of R, F, ln, at 25ºC. E = Eo − 0.059 log [P] cell cell n 10 [R] Nernst equation is valid for cell as well as half cell. GIBB'S ENERGY CHANGE AND CELL POTENTIAL The electrical work done from the cell = nFE ∴ – ΔG = nFE ΔG = –nFE or ΔGº = – nFEº where ΔG = Gibb's energy change ΔGº = standard Gibb's energy change E = emf of cell/electrode potential Eº = standard emf of cell/standard electrode potential. ELECTROLYSIS The decomposition of the electrolyte due to the passage of electricity is known as electrolysis During electrolysis electrical energy changes into chemical energy. To have electrolysis we must use D/C. Electrolysis is a redox phenomenon. Faraday’s laws of electrolysis First Law : The weight of substance deposited at the electrode is directly proportional to the quantity of electric charge passed through the electrolytic solution W ∝ Q W = wt in g Q = quantity of charge ⇒ W = ZQ Z = constant known as electrochemical equivalent ⇒ W = ZIt I = Current in ampere t = time in sec Second law : If the same quantity of current is passed through different electrolytic solutions then the weights of different substances deposited at the respective electrodes is directly proportional to their chemical equivalents. wt. of electrolyte (A) = wt. of electrolyte (B) Eq. wt. of (A) Eq. wt. of (B) ELECTROLYTIC CONDUCTION Molten electrolyte and the aquous solution of electrolytes contain free ions and conduct electricity due to the movement of ions According to Ohm's law R ∝ l a or R = ρ × l a ∴ ρ = R × a l 1 = 1 × l ρ R a where R = Resistance of solution l = Length a = Area of cross section of the solution ρ = Resistivity of solution If l = 1 unit of length and a = 1 unit of area then Conductance (G) = 1 R Unit of conductance = ohm–1 or siemens (s) Specific conductance (K) = 1 = 1 × l = cell constant ( l is called cell constant) ρ R a Resistance a K = Conductance if l = 1 unit (length) and a = 1 unit (area) S.I. Unit of K = Sm–1 MOLAR CONDUCTIVITY (∧m ) The conducting power of all the ions furnished by one mole of an electrolyte in any solution is termed as its molar conductivity ∧m = Conductivi ty (K) Concentration of solution in moles per unit volume (cm3 ) Unit of ∧m = ohm–1 cm2 mol–1 Molar conductivity increases with increase in dilution or decrease in concentration due to increase in ionisation but specific conductivity decreases with increase in dilution because number of ions per cm3 of solution decreases. KOHLRAUSCH'S LAW It states that at infinite dilution the molar conductivity of an electrolyte can be expressed as the sum of the contribution from its individual ions. ∧∞ = V λ∞ + V λ∞ (where V and V are the number of cations and anions per formula unit of electrolyte m + + − − + – respectively and λ∞ and λ∞ are molar conductivities of the cation and anion at infinite dilution respectively). Application Some typical applications of Kohlrausch's law are Determination of ∞ for weak electrolyte Example : ∧∞ of CH COOH (weak acid) can be determined by knowing ∧∞ of three strong electrolytes m 3 m NaCl, HCl and CH3COONa ∧∞CH COOH = λ∞CH COO− + λ∞ + 3 3 H ∧∞ = ∧∞ + +λ∞ − NaCl Na Cl ∧∞ = λ∞ + + λ∞ − HCl ∧∞ H = λ∞ Cl − + λ∞ + CH3COONa CH3COO Na ∞ CH3COOH = HCl = λ∞ ∞ CH3COONa − + λ∞ + ∞ NaCl CH3COO H Determination of degree of dissociation (α) of a weak electrolyte ∧c α = m m Determination of ionisation constant of weak electrolyte of (AB) type K = ∞ C(∧c )2 Cα2 = ∧m (∧∞ − ∧c ) C(1− α) Determination of solubility of a sparingly soluble salt Concentration of sparingly soluble salt (C ) = 1000.Ksalt m (V λ∞ + V λ∞ ) + + − − ❑ ❑ ❑