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ELECTROCHEMISTRY, , , , , , , , , , 1. ELECTROCHEMISTRY, , Electrochemistry is the study of production of el, from the energy released during a spontaneous chemical, , , , tricity, , reaction and the use of electrical energy to bring about, non-spontaneous chemical transformations., , 2. ELECTROCHEMICAL CELLS, , A spontaneous chemical process is the one which can, take place on its own and in such a process the Gibb’s, energy of the system decreases. It is this energy that, gets converted to electrical energy.The reverse process, is also possible in which we can make non-spontaneous, , occur by supplying external energy in the form, , , , , ctrical ent ¢ inter conversions are carried, , , , gy. Thes, , , , out in equipments called Electrochemical Cells., , aids), , Electrochemical Cells are of two types, , , , 3.1 Galvanic Cells, , , , , , Converts chemical energy into electrical energy, , , , Cells, , , , , , , , 3.2 Electroly, , , , Converts electrical energy into chemical energy., , 4. GALVANIC CELL, , Cell energy is extracted from a spontaneous chemical, , , , process or reaction and it is converted to electric current., For example, Daniell Cell is a Galvanic Cell in which Zine, and Copper are used for the redox reaction to take place, Zn (s) + Cu’* (aq) —» Zn™* (aq) + Cu(s), , Oxidation Half: Zn (s) —» Zn** (aq) + 2e7, , , , Reduction Half: Cu”*(aq) + 2e” —» Cu(s), , Zn is the reducing agent and Cu* is the oxidising, agent. The half cells are also known as Electrodes. The, oxidation half is known as Anode and the reduction half is, called Cathode. Electrons flow from anode to cathode in, , , , the external circuit. Anode signed negative polari:, and cathode igned positive polarity, In Daniell Cell,, Zn acts as the anode and Cu acts as the cathode., , , , , , ELECTROLYTIC CELL, , These electrodes are dipped in and electrolytic solution, containing cations and anions. On supplying current the, , ions move towards electrodes of opposite polarity and, , , , , , simultaneous reduction and oxidation takes place, , , , 5.1, , , , Preferential Discharge of ions, , , , , , Where there are more than one cation or anion the process, of discharge becomes competitive in nature. Discharge, of any ion requires energy and in case of several ions, being present the discharge of that ion will take place, , , , , , which requires the energy, , ELECTRODE POTENTIAL, , It may be defined as the tendency of an element, when it is, placed in contact with its own ions to either lose or gain, electrons and in turn become positively or negatively charged., , The electrode potential will be named as oxidation or, reduction potential depending upon whether oxidation or, reduction has taken place., , , , , , , , , , , , , , M(s)=S== M™ (aq) + neo, M™ (aq) + ne" See M (5), 6.1 Characte, , , , , , , , , , (a) Both oxidation and reduction potentials are equal in, magnitude but opposite in sign., , (b) It is not a thermodynamic property, so values of E, not additive., , , , , , STANDARD ELECTRODE POTENTI, , , , (E, , It may be defined as the electrode potential ofan electrode, determined relative to standard hydrogen electrode under, standard conditions. The standard conditions taken are

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8., , (i) 1M concentration of each ion in the solution., (ii) A temperature of 298 K., (iii) 1 bar pressure for each gas., , . ELECTROCHEMICAL SERIES, , The half cell potential values are standard values and are, represented as the standard reduction potential values, as shown in the table at the end which is also called, Electrochemical Series., , . CELL POTENTIAL OR EMF OF A CELL, , |, , The difference between the electrode potentials of two, half cells is called cell potential. It is known as electromotive, force (EMF) of the cell ifno current is drawn from the cell, , , , Ey =E, , For this equation we take oxidation potential of anode and, reduction potential of cathode., , Since anode is put on left and cathode on right, it follows, therefore,, , =E,+E,, For a Daniel cell, therefore, , Beet = EC cu, , -E,, , Zoltan, , = 0.34+(0.76)=1.10V, , 10. CELL DIAGRAM OR REPRESENTATION, OF ACELL, , The following conventions or notations are applied for writing, the cell diagram in accordance with IUPAC recommendations., The Daniel cell is represented as follows, , Zn(s) | Zn** (C,) || Cu** (C,) | Cu(s), , (a) Anode half cell is written on the left hand side while, cathode half cell on right hand side., , (b) A single vertical line separates the metal from aqueous, solution of its own ions., , Zn(s)|Zn** (aq);, , Anodic chamber, , Cu** (aq)| Cu(s), , Cathodic chamber, (c) A double vertical line represents salt bridge, , (d) The molar concentration (C) is placed in brackets after, the formula of the corresponding ion,, , (e) The value of e.m.f. of the cell is written on the extreme, right of the cell. For example,, , Zn(s) | Zn** (1 M)\|Cu** (1 M)|Cu EMF =+1.1V, , (f) If an inert electrode like platinum is involved in the, construction of the cell, it may be written along with the, working electrode in bracket say for example, when a zine, anode is connected to a hydrogen electrode, , Zn(s)|Zn** (C,)\) H* (C,)|H, | (Pt)(s), , 11. SALT BRIDGE, , Salt bridge is used to maintain the charge balance and to, complete the circuit by facilitating the flow of ions through it, It contains a gel in which an inert electrolyte like Na,SO, or, KNO, etc are mixed. Negative ions flow to the anode and, positive ions flow to the cathode through the salt bridge and, charge balance is maintained and cell keeps on functioning., , , , , , Voltmeter/Ammeter, , , , , , , , , , =_.", Ft Coe x, , Y, , , , , , Copper, , Zine Rod, , Rod Salt bridge, , , , CuSO,, , , , , ZnSO, La, Sol Sol., Anode Movement of Cations, Cathode, Zn-Zn"+2e Movement of Anions Cy" +26 —> Cu, , Electrochemical cell based on the redox reaction,, Zn + CuSO, ~» ZnSO, + Cu, , , , 12. SPONTANEITY OF A REACTION, , AG == FE,, , For a spontaneous cell reaction AG should be negative, and cell potential should be positive, , If we take standard value of cell potential in the above, equation we will obtain standard value of AG as well., , 3° = — nF?, AG? = ~ nFE G3,, , 13. TYPES OF ELECTRODES, , , , , , 13.1 Metal-Metal lon electrodes, , , , , , , , A metal rod/plate is dipped in an electrolyte solution, containing metal ions. There is a potential difference, between these two phases and this electrode can act as a, cathode or anode both

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Anode: M —> M™ + neo, , Cathode: M™ + ne —» M, , , , , , 13.2 Gas Electrodes, , , , , , Electrode gases like H,, Cl, etc are used with their respective, ions. For example, H, gas is used witha dilute solution of HCI, (H’ ions). The metal should be inert so that it does not react, with the acid., , , , , , A solution of acid, containing an, effective hydrogen|, ion concentration, of 1 molar, , , , , , , , , hydrogen gas, , at 1 atoms connecting wire, , Platinum electrod, coated with finely, divided platinum, {platinum black}, , , , ‘Standard Hydrogen Electrode, , , , Anode: H, —» 2H” + 2e7, , Cathode: 2H” + 2e” —» H,, , The hydrogen electrode is also used as the standard to, measure other electrode potentials. Its own potential is, set to 0 V as a reference. When it is used as a reference, the concentration of dil HCI is taken as 1 M and the, electrode is called “Standard Hydrogen Electrode (SHE)”., , , , , , 13.3 Metal-Insoluble salt electrode, , , , , , We use salts of some metals which are sparingly soluble, with the metal itself'as electrodes. For example, if we use, AgCl with Ag there is a potential gap between these two, phases which can be identified in the following reaction, , AgCK(s) + e" —» Ag(s) + Cr, , This electrode is made by dipping a silver rod ina solution, containing AgCI(s) and CI” ions., , , , , , 13.4, , Calomel Electrode, , , , , , Mercury is used with two other phases, one is a calomel, paste (Hg,Cl,) and electrolyte containing CI” ions., , , , , , , , , , , , , , , , Pt. wire, , To salt A, , boxer Saturated, [= KCI solution, , , , , , , , , , , , , , , , , , , , , , Hg,Cl, and, Hg paste, cf Mercuy, Calomel electrode, Cathode :, Hg,Cl,(s) +2 —» 2Hg(/) +2Cr(aq), Anode :, , 2Hg(1) + 2Ct(aq) —» Hg,CL(s) +2, This electrode is also used as another standard to measure, , other potentials. Its standard form is also called Standard, Calomel Electrode (SCE)., , , , , , Redox Electrode, , , , , , In these electrodes two different oxidation states of the, same metal are used in the same half cell. For example,, Fe** and Fe* are dissolved in the same container and an, inert electrode of platinum is used for the electron transfer., Following reactions can take place:, , Anode: Fe? —» Fe*+e, , Cathode: Fe** + e" —» Fe*, , 14. NERNST EQUATION, , It relates electrode potential with the concentration of ions., , Thus, the reduction potential increases with the increase, in the concentration of ions. For a general electrochemical, reaction of the type, , aA + bB—*>cC + dD, , Nernst equation can be given as, , ore, RT [TIPE, a oF" [AP Bl

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E., =E. BB pr tog (Cl PT., ne TY, , Substituting the values of R and F we get, , 0.0591 (cT PY 298K, , fay [By, , 15. APPLICATIONS OF NERNST EQUATION, , 15.1 Equilibrium Constant from Nernst Equation, , , , E., =Ety- log, , , , , , , , , , Fora Daniel cell, at equilibrium, , , , [zat, Eq =0= Ete a, - ge, = 2303RT [Zn**], , , , et = Fp “Tea |, , , , [za],, Butat equilibrium, [eu* ] ©, , , , , , 2.303RT, E‘, = ———log K,, ceil oF 0g K., 2.303x8.314x 298 ‘, “<a _»_acenn ~~ K,, 2x 96500, 0.0591 ., =—7 bak., 0.0591, Ingeneral, Et = log K,, nE*., “ los Ke= 99591, , 16. CONCENTRATION CE, , If two electrodes of the same metal are dipped separately, into two solutions of the same electrolyte having different, , , , concentrations and the solutions are connected through, salt bridge, such cells are known as concentration cells., For example, , H,| H"(c,) || H’ (c,)|H,, Cu| Cu (¢,)|| Cu*(,)|Cu, , These are of two types:, , , , 16.1 Electrode concentration cells, , , , , , , , H,(P,)|H*(C)||H"(C)|HP), , , , where p, <p, for spontaneous reaction, , , , , , 16.2 Electrolyte concentration cell, , , , , , The EMF of concentration cell at 298 K is given by, Zn | Zn* (c,) | Zn* (c,)|Zn, , 0.0591, Ea = log,, ny C, , , , , , where c, > c, for spontaneous reaction, , 17. CASES OF ELECTROLYSIS, , 17.1 Electrolysis of molten sodium chloride, , , , , , , , , , 2NaC/(/) = 2Na*(/)+2Cr(/), , The reactions occurring at the two electrodes may be, shown as follows :, , At cathode :, , 2Na* + 2¢° ->2Na E°=-2.71V, Atanode:, , 2Cr > Cl, +2e7 E=-1.36V, Overall reaction :, , 2Na* (/)+2CI (J) Electrolysis, 2Na(I) +Cl, (g), , or 2NaClI(/) Electrolysis, 2Na(/) + Cl,(g), , Atcathode At anode

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17.2 Electrolysis ofan aqueous solution of sodium chloride, , , , , , , , , , 18.3 Mercury cells, , , , NaCl (aq) —> Na*(aq) + CI" (aq), (almost completely ionized), H,O(/) = H*(aq) + OH(aq), (only slightly ionized), Atcathode :, 2Na* +2e7 —» 2Na E*=-2.71V, 2H,O + 2e" —» H, + 20H” E*=- 0.83V, , Thus H, gas is evolved at cathode value Na” ions remain, in solution., , Atanode :, 2H,O —» 0,+4H*+4e" EE =-1.23V, 2Cr —» Cl, + 2e° E*=-1.36V, , Thus, Cl, gas is evolved at the anode by over voltage, concept while OH" ions remain in the solution., , eG TBE, , When Galvanic cells are connected in series to obtain a, higher voltage the arrangement is called Battery., , , , 18.1 Primary Batteries, , , , , , Primary cells are those which can be used so long the, active materials are present. Once they get consumed the, cell will stop functioning and cannot be re-used. Example, Dry Cell or Leclanche cell and Mercury cell, , , , , , 18.2 Drycell, , , , , , Anode : Zn container, , Cathode : Carbon (graphite) rod surrounded by powdered, MnO, and carbon., , Electrolyte : NH,Cl and ZnCl,, , Reaction:, , Anode : Zn —» Zn™* + 2e7, , Cathode : MnO, + NH} +e” —> MnO (OH)+ NH,, , The standard potential of this cell is 1.5 V and it falls as, the cell gets discharged continuously and once used it, cannot be recharged., , ‘These are used in small equipments like watches, hearing aids., Anode : Zn— Hg Amalgam, , Cathode : Paste of HgO and carbon, , Electrolyte : Paste of KOH and ZnO, , Anode : Zn (Hg) + 20H” —> ZnO (s) +H,O+ 2e7, Cathode : HgO (s)+ H,O+ 2e" —» Hg () + 20H™, Overall Reaction : Zn (Hg) + HgO (s) —> ZnO (s)+ Hg (), , The cell potential is approximately 1.35V and remains, constant during its life, , , , , , , , 18.4 Secondary Batteries, , , , Secondary cells are those which can be recharged again, and again for multiple uses. e.g. lead storage battery and, Ni-Cd battery., , , , 18.5 Lead Storage Battery, , , , , , , , , , Anode : Lead (Pb), Cathode : Grid of lead packed with lead oxide (PbO.), Electrolyte : 38% solution of H,SO,, Discharging Reactions, Anode: Pb(s) + SO,*(aq) —» PbSO (s) + 2eCathode: PbO,(s) + 4H"(aq)+ SO}(aq) +20», PbSO ,(s) + 2H,O(), Overall Reaction : Pb(s) + PbO,(s) + 2H,SO (aq) —>, 2PbSO (s) + 2H,O(), , To recharge the cell, it is connected with a cell of higher, potential and this cell behaves as an electrolytic cell, and the reactions are reversed. Pb(s) and PbO.(s) are, regenerated at the respective electrodes. These cells, deliver an almost consistent voltage, , Recharging Reaction : 2PbSO ,(s) + 2H,O(/) —> Pb(s) +, , 19. FUEL CELLS, , A fuel cell differs from an ordinary battery in the sense, that the reactants are not contained inside the cell but are, , PbO,(s)+2H,SO (aq)