Electrochemistry Notes

ELECTROCHEMISTRY ELECTROCHEMISTRY Electrochemistry: Study of relationship between chemical energy and electrical energy and how one can be converted into another. Electrolytic cell: Electrical energy  Chemical energy Electrochemical cell : chemical energy  Electrical energy (Galvanic cell/ voltaic cell / Daniel cell) Electrode Potential: Potential difference between the electrode and the electrolyte. Standard electrode potential : Potential difference developed in a cell when the gives electrode which is in contact with its ions having concentration 1 mol L–1 is coupled with standard hydrogen electrode. Cell potential : Potential difference between the 2 electrodes of a galvanic cell. EMF : Potential difference between the 2 electrodes of a galvanic cell when no current is drawn through the cell. Ecell = ER - RL = Ecathode – Eanode. Cell Reaction : Cu(s) + 2Ag+(aq) Cu+2(aq) + 2Ag(s) At Cathode(Reduction) : 2Ag+(aq) + 2e 2Ag(s) At anode (Reduction) : Cu(s) Cu+2(aq) + 2e– Cell Representation : Cu | Cu+2 || Ag+ | Ag Ecell = E Ag / Ag  E Cu 2 / Cu SHE : Standard hydrogen electrode H2 gas 1 M HCl H+(aq) + e–  Eo  H / H2 1 H (g) 2 2 =0 Electrochemical Series : The arrangement of metal (or species) in their increasing reducing power (or decreeing Eº) is known as electrochemical series. Nernst Equation: To calculate electrode potential at any concentration.  Mn(aq )  ne  M( s) E Mn  / M  E o n  M /M  RT [M( s ) ] ln nF [Mn  ] But concentration of solid M is taken as unity ' EM  Eo n  n /M M /M  RT 1 ln nF Mn   ELECTROCHEMISTRY E Or Mn / M  Eo n M /M    RT ln Mn nF F = 96487 C mol–1 R = 8.314 Jk-1 mol–1 [Mn+] = concentration of species. T = 298 k n = no of moles of e indicated in balanced eqn- E n M /M  Eo n M /M    0.059 ln Mn n In Daniel cell (1) Zn | Zn+2||Cu+2 | Cu (2) Ni | Ni+2 ECell  Eocell    0.059 Cu 2 log 2 Zn  2 +1 || Ag | Ag E Cell  E ocell    0.059 [ Ag1]2 log 2 [Ni 2 ] For general electrochemical reaction aA + bB Ecell  Eocell  – ne cC + dD RT C Dd ln nF Aa Bb c At equilibrium condition Cc Dd A a Bb ECell = 0,  0 = Eº – = KC RT 2.303 RT logK C lnKC = nF nF RT lnKC nF Electrochemical cell and Gibbs energy of the reaction G = –nFE Total Charge = nF, E = EMF G = –RTlnKC Eº = Conductance (G) – The inverse or reciprocal or R The case with which current flows through a conductor. 1 SI unit  Siemen = ohm–1 R Conductivity (k) – The reciprocal of resistivity G= k= 1 l = SI unit Sm–1  RA Cell constant (G*) The quantity  G* is called cell constant. k = . A R Electronic conductance – electrical conductance though metal. Electrolytic or ionic conductance – The conductance by ions present in solution. ELECTROCHEMISTRY Conductivity of ionic solution depends on 1. Nature of electrolyte. 2. Size of the ions produced and their solvation. 3. Nature of the solvent and its viscosity 4. Concentration of electrolyte. 5. Temperature (Increases with increase of T) Molar conductivity.: m  k k  1000  C Molarity Equivalent conductivity  eq  k  1000 Normality Limiting molar conductivity : The molar conductivity of a solution whose concentration approaches zero. Debye – Huckel Onsanger equation  = º = AC1/2 Kohlrausch law of independent migration of ions the law states that the limiting molar conductivety of an electrolyte at infinite dilution is opposed as the sum of individual contributions of the anion and cation of the electrolyte. oNaCl  o Na   o Cl  , 0CaCl  0 2 Ca  2  2 0 Cl  In General  0m    0    0 v+ & v– are the no. of cations and anions performed of electrolyte. Applications of Kohlrausch Law : (i) In the calculation of 0m for weak electrolyte (ii) In the calculation of degee of dissociation Degree of Dissociation   m m (iii) In the calculation of dessociation of constant of weak electrolyte. Ka  c 2 C2m  1    0m 0m  m   Faradays laws of electrolysis ELECTROCHEMISTRY First law – It states that the amount of any substance produced at any electrode during electrolysis is directly proportional to the quantity of charge passed. W  Q W = zit Second law – When the same quantity of electricity is passed through a solution of different electrolytes, thus the weight of the substance produced at the respective electrodes are directly proportional to their equivalent weights. WA Eq. Wt of A  WB Eq. wt of B FARADAY : Thecharge on one mole of electrons is called 1 faraday. Thus 1 faraday = 96487 C So, charge on n mole of electrons will be Q = nF Products of electrolysis – depends on nature of materiel electrolysed and the types of electrodes used. Electrolysis of (i) Molten NaCl solution  Na + Cl2 (ii) Aqueous NaCl solution  NaOH + H2 + Cl2 u sing Pt electrodes Cu + Cl (iii) CuCl2   2 using Cu electrodes (iv) CuCl2   Cu is deposited at cathode and dissolves at anode. u sing Pt electrodes Silver at Cathode and O at anode (v) AgNO3   2 using (vi) AgNO3 Ag electrode Silver at Cathode and silver anode will be dissolved slowly.. Commercial Cells: Two type Primary Cells/Battries : Reaction occurs only once * Cannot be reused. * Electrode reaction cannot be reversed by external electric source. eq. Daniel cell, dry cell, Mercury cell . Dry cell /Lechlanche cell : At anode Zn(s)  Zn+2 + 2e Craphite(C) Cathode MnO2 + C At Cathode MnO2 + NH4+ + e  MnO(OH)2 + NH3 NH3 gas produced reacts with Zn+2 to form complex 2NH3 + Zn+2  [Zn(NH3)4]2+ Cell potential = 1.5 V Paste of NH4Cl and ZnCl2 Zn anode Mercury cell - for low current devices eq. hearing aids, watches etc. Anode - Zn –Hg Amalgam Cathode – paste of HgO and C. Electrolyte – paste of KOH and ZnO Anode: Zn(Hg) + 2OH–  ZnO(s) + H2O + 2e Cathode : HgO + H2O + 2e  Hg(l) + 2OH– ELECTROCHEMISTRY Overall Reaction : Zn(Hg) + HgO(s)  ZnO(s) + Hg(l) Cell potential = 1.35 (Remains constant through out life as the overall reaction does not involve any ion in solution ) Secondary cells: - Can be recharged by passing electric current eq. Lead storage battery, nickel cadmium cell Lead storage battery - in Automobiles and invertor. Anode -Pb Cathode - Grid of Pb packed with lead dioxide (PbO2). Electrolyte : H2SO4 (38% by mass) When battery in use (Discharging) Anode : Pb(s) SO4–2 (aq) PbSO4(s) + 2eCathode : PbO2(s) + SO4–2(aq) + 4H+(aq) + 2e  PbSO4(s) + 2H2O(l) Overall reaction Pb(s) PbO2(s) + 2H2SO4  2PbSO4(s) + 2H2O(l) On Recharging -Reaction is reversed In Recharging process PbSO4 (s) on anode and cathode is converted into Pb and PbO2 respectively. Ni–Cd Cell : It has longer life, more expensive then Pb storage battery but it is light therefore used in calculators, portable power tools. Overall reaction : Cd(s) + 2Ni(OH)3(s)  CdO(s) + 2Ni(OH)2(s) + H2O Fuel Cells : Galvanic Cells that are designed to convert the energy on combustion of fuel like H2, CH4 CH3OH etc directly into electrical energy are called fuel cells. Catalyst like Pt or palladilum are incorr proated into electrode H2O * Efficiency 70% Anode Cathode _ * Pollution free + H2 O 2 Cathode: Anode : O2 + 2H2O + 4e  4OH– 2H2 + 4OH–  4H2O + 4e NaOH(eq) Overall Reaction : 2H2 + O2  2H2O(l) * H2–O2 fuel cells were used in Apollo space program for providing electrical power. The cell runs continuously as long as reactants are supplied. Corrosion : In corrosion metal is oxidised by loss of electron to oxygen and formation of oxide. eq. tarnishing of Ag. Green coating on Cu and Bronge rusting of iron. (In presence of O2 & H2O) Anode (Oxidation): 2Fe 2Fe+2 + 4e Electrons released at anodic sport move through metal and go to another spot on the metal and reduce oxygen in presence of H+(obtained from H2CO3 formed due to disolution of CO2 into water.) Cathode : O2(g) + 4H+(eq) + 4e  2H2O. The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide ((Fe2O3x H2O) Prevention of corrosion : - Covering surface of metal by paint, chemicals (Bisphenol) or other metals (Sn, Zn,) etc, Galvanisation.