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Electrochemistry

Electrochemistry. Prof. Dr. Sabine Prys. http://www.iccb.org/student/mod/science/mod_chem1/mod1/p1.html. @designed by ps. 3.3 Normal & Standard Conditions. normal conditions: normal pressure p = 1 atm = 101,325 kPa = 1013,25 mbar normal temperature T = 0°C = 273.15 K

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Electrochemistry

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  1. Electrochemistry Prof. Dr. Sabine Prys http://www.iccb.org/student/mod/science/mod_chem1/mod1/p1.html @designed by ps

  2. 3.3 Normal & Standard Conditions normal conditions:normal pressure p = 1 atm = 101,325 kPa = 1013,25 mbarnormal temperature T = 0°C = 273.15 K standard conditions: standard pressure p = 1 atm = 101,325 kPa = 1013,25 mbarstandard temperature T = 25°C = 298.15 K Such definitions can vary according to different sources: IUPAC, NIST, …

  3. 3.4 Enthalpy • Enthalpy is a measure of the total energy of a thermodynamic system including • the internal energy (energy required to create a system), • the amount of energy required to make room for it by displacing its environment and establishing its volume and pressure. • Enthalpy is a thermodynamic potential, a state function and an extensive quantity (i.e. depending on amount material). H enthalpy U internal energy P pressure V volume http://goldbook.iupac.org

  4. 3.8 GIBBS’ Free Energy G Useful energy, or energy available to do work G G = free energy H = GIBBS’ energy (enthalpy) U = internal energy T = Kelvin temperature S = entropy p = pressure V = volume TDS is the energy not available for doing work

  5. 3.9 Spontaneity of Redox Reactions DH DS Spontaneity Exothermic DH < 0 Increase DS > 0 + DG < 0 Exothermic DH < 0 Decrease DS < 0 + if |TDS| < |DH| Endothermic DH > 0 Increase DS > 0 + if TDS > DH Endothermic DH > 0 Decrease DS < 0 - DG > 0

  6. 3.10 Thermodynamical Equilibrium • Reversible processes ultimately reach a point where the rates in both directions are identical, so that the system gives the appearance of having a static composition at which the Gibbs energy G is a minimum DG = 0 • At equilibrium the sum of the chemical potentials of the reactants equals that of the products, so that: DG = DG298 + RT . lnK = 0  DG298 = - RT . lnK • The equilibrium constant K is given by the mass-law effect. http://goldbook.iupac.org

  7. 3.12 Maximum Work Wmax = maximum work G = Gibbs free energy R = gas constant T = Kelvin temperature K = equilibrium constant z = ion charge n = moles F = Faraday‘s constant E = galvanic cell potential U = voltage I = currant t = time

  8. 4.0 Chemical Solutions • Suspension (particle diameters 10-4 - 10-5 cm ) solid particles in homogeneous fluid • Colloid (particle diameters 10-5 - 10-7 cm ) microscopically dispersed particles in another substance • Solution (particle diameters 10-7 - 10-8 cm ) Homogeneous phase with at least 2 components: solvent and solute • gas in liquid e.g. O2 in H2O • gas in solid e.g. H2 in palladium • liquid in liquid e.g. petroleum • solid in liquid e.g. NaCl in H2O •  electrolytes in water

  9. 4.6 Ion activity High ion concentrations in aqueous solutions ð ion – ion interactions: pH measured < pH calculated (1m, 0.1 m solution of acids) ion activity: a = activity, f = activity coefficient, c = concentration f (HCl, 25°C): 0.001m/0.965 0.01m/0.905 0.1m/0.794 1m/0.809

  10. 4.7 Colloidal Solutions • Larger particles in solvent, e.g. macromolecules / polymers • Properties depend on solute size and not on solute concentration ! • Coagulation: growth of larger particles by smaller particles consumption • Hydrophobe colloids: large surface, large adsorption properties • Hydrophile colloids

  11. 4.9 Electrolytes Electrolyte: solution which conducts electrical current Hydrated H3O+ Hydrated OH-

  12. _ + - - - - - - + + + + + + cathode cat ions anode anions 4.9.1 Electrical Conductivity in Solutions Electrolytes • solutions which support ion transport • salts in aqueous solutions, e.g. KCl, ZnSO4, CuCl2, etc. • molten salts Conductivity L (resistance R) bad electrolyte: distilled water: 0.0548 µS/cm at 25 °C. H2O

  13. 4.9.2 Specific Conductivity absolute electrolyte conductivity R = solution resistance specific electrolyte conductivity A = electrode surface, l = electrode distance

  14. 4.9.3 Example: Proton Migration Grotthuss Diffusion structural defect migration mesomeric structures between H9O4+ and H5O2+,

  15. - + - + H2 Cl2 H2 Cl2 5.0 Electrochemical Cells electrolytic cell galvanic cell 2 HCl(aq) ð H2(g) + Cl2(g) H2(g) + Cl2(g) ð2 HCl(aq) electrical energy  chemical energychemical energy  electrical energy

  16. 5.1 Electrolysis electrolysis: decomposing materials by electric current H2SO4 + 2H2O ð 2H3O+ + SO42- water electrolysis cathodic reduction 4H3O+ + 4e-ð 2 H2ñ + 2 H2O anodic oxidation 4 OH-ð 2 H2O + O2ñ + 4 e- total 2 H2O (l) ð 2 H2(g) + O2(g) electrods battery ca. 15 V H20 + H2SO4 1:10

  17. Q = electric charge in C n = yield in mol F = Faraday‘s constant = 96485,309 As / mol Ec = electrochemical equivalent M = ion weight z = ion charge NL = Lohschmidt‘s number e = elementary charge 5.1.1 Electrochemical Equivalent

  18. m = Ec . Q = Ec .I. t m = mass yield in g Ec = electrochemical equivalent Q = electric charges in Coulomb I = current strength t = electrolysis time 5.1.2 Faraday‘s Laws ma ,mb = mass yield in g for material a / b Ma,Mb = molecular weight for material a / b za, zb = chemical valency for material a / b

  19. voltmeter ca 1,1 V Zn Cu 1 m ZnSO4 1 m CuSO4 diaphragma (pottery) bridge containing KCl solution 5.2 Galvanic Elements Daniell Element: 2 galvanic half cells + bridge Zn / ZnSO4 // CuSO4 / Cu electrode reactions Zn (cathode)ð Zn2+ + 2e- Cu2+ + 2e-ð Cu (anode) Zn metal in ionic solution Cu ions in Cu metal electrical current results from different oxidation affinities

  20. H2 gas Pt electrode 5.4 Standard Hydrogen Electrode standard hydrogen electrode (SHE) = reference potential = E0 = 0 V H2ñ ð 2H+ + 2e- p = 1,01325 bar T = 25°c a(H+) = 1 mol / l c(H+) = 1,235 mol / l (HCl)

  21. H2 gas Pt electrode 5.5 Metal Standard Potentials standard hydrogen electrode = reference potential E0 = 0 V metal electrode / metal salt solution at standard conditions = standard metal potential M ðMz+ + ze- p = 1,01325 bar T = 25°c c(Mz+ ) = 1 mol / l pH < 6 precipitation prevention

  22. 5.5.1 Metal Standard Potential Tables pH-dependant

  23. 5.5.2 Calvanic Corrosion Potential Chart Galvanic Corrosion Potential Chart K, Na, Mg, Al, Zn, Fe, Pb, Cu, Ag, Au passivation of Al, Mg, Mn, Cr alternative corrosion potential charts for industrial materials cathode least noble corroded metals strong oxidation affinity negative oxidation potential anode most noble protected metals weak oxidation affinity positive oxidation potential

  24. 5.6.3 Exercise: Gibbs Free Energy WhathappensifΔG = 0

  25. electrode potential dependency on temperature and concentration E = measured cell potential E0 = standard reaction potential R = gas constant ( 8,3145 J . mol-1 . K-1) T = Kelvin temperature z = charges F = Faraday’s constant [ ] = concentration of oxidant / reductant in mol / l 5.7 NERNST‘s Equation 1

  26. type electrode ( = metal electrode in metal salt solution) [red] = const E = measured cell potential E0 = standard reaction potential R = gas constant ( 8,3145 J . mol-1 . K-1) T = Kelvin temperature z = charges F = Faraday’s constant [ox] = concentrationen of oxidant in mol / l 5.7.1 NERNST’s Equation 2

  27. 5.7.2 Exercise: Maximum Electrical Voltage • Calculate the maximum electrical voltage for the Daniell element when standard conditions ! Daniell Element: Cu/Cu++//Zn++/Zn Cu/Cu++/ +0,34 V Zn++/Zn/ +0,76 V S = + 1,1 V • Calculate the maximum electrical voltage for a galvanic cell with Ni/Ni++//Zn++/Zn when standard conditions ! Ni/Ni++// -0,23 V Zn++/Zn/ +0,76 V S = + 0,53 V

  28. 5.7.3 Exercise: Nernst Equation What is the electrode potential for a silver electrode at 0°C when the Ag+ concentration is 1 mol ?

  29. 2. type electrode = metal electrode in saturated metal salt solution = electrode with constant potential (no concentration changes) T = 25 °C: 1 m KCl E0 = + 0,220 V sat. KCl E0 = + 0,1958 V Ag K+ Ag+ Cl- AgClsat 5.8 Ag / AgCl Electrode

  30. 5.8.1 Concentration Cells • Cu(s) | Cu2+ (0.05 M) || Cu2+ (2.0 M) | Cu(s) • half cellreactions : oxidation: Cu(s) → Cu2+ (0.05 M) + 2 e– reduction: Cu2+ (2.0 M) + 2 e– → Cu(s) overallreaction: Cu2+ (2.0 M) → Cu2+ (0.05 M) • cell'semf : • E = E°- (0.05916\2) log [0,05/2] = 0.0474 V • E° = 0 , (electrodesandionsarethe same in both half-cells)

  31. 5.15 Dry Cells moist electrolyte paste Leclanché's cell • anode is a zinc container surrounded by a thin layer of MnO2 • Cathode a carbon bar inserted on the cell's electrolyte • moist electrolyte paste NH4Cl + ZnCl2 mixed with starch Anode: Zn(s) → Zn2+(aq) + 2 e– Cathode: 2 NH4+(aq) + 2 MnO2(s) + 2 e– → Mn2O3(s) + 2 NH3(aq) + H2O(l) Overall reaction: Zn(s) + 2 NH4+(aq) + 2 MnO2(s) → Zn2+(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l) E = ~ 1.5 V

  32. 5.16 ZnBattery Graphics: http://en.wikipedia.org/wiki/File:Zincbattery.png

  33. 5.17 Mercury Battery amalgamated anode of mercury and zinc surrounded by a stronger alkaline electrolyte and a paste of ZnO and HgO Mercury battery half reactions are shown below: Anode: Zn(Hg) + 2 OH–(aq) → ZnO(s) + H2O(l) + 2 e– Cathode: HgO(s) + H2O(l) + 2 e– → Hg(l) + 2 OH–(aq) Overall reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l) no changes in the electrolyte's composition when working 1.35 V of direct current Not rechargeable Graphics: http://en.wikipedia.org/wiki/File:Mercurybattery.png

  34. 5.18 Lead-Acidbattery six identical cells assembled in series (6 x 2V ) = 12 V lead anode lead dioxide cathode Electrolyte sulfuric acid Anode: Pb(s) + SO42–(aq) → PbSO4(s) + 2 e– Cathode: PbO2(s) + 4 H+(aq) + SO42–(aq) + 2 e– → PbSO4(s) + 2 H2O(l) Overall reaction: Pb(s) + PbO2(s) + 4 H+(aq) + 2 SO42–(aq) → 2 PbSO4(s) + 2 H2O(l) Rechargeable (external voltage  electrolysis of the products) http://en.wikipedia.org/wiki/Lithium-ion_battery

  35. 5.19 Lithium rechargeablebattery (1) Positive electrodes Electrode material Average potential differenceSpecificcapacitySpecificenergy LiCoO2 3.7 V 140 mA·h/g 0.518 kW·h/kg LiMn2O4 4.0 V 100 mA·h/g 0.400 kW·h/kg LiNiO2 3.5 V 180 mA·h/g 0.630 kW·h/kg LiFePO4 3.3 V 150 mA·h/g 0.495 kW·h/kg Li2FePO4F 3.6 V 115 mA·h/g 0.414 kW·h/kg LiCo1/3Ni1/3Mn1/3O2 3.6 V 160 mA·h/g 0.576 kW·h/kg Li(LiaNixMnyCoz)O2 4.2 V 220 mA·h/g 0.920 kW·h/kg Negative electrodes Graphite (LiC6) 0.1-0.2 V 372 mA·h/g 0.0372-0.0744 kW·h/kg HardCarbon (LiC6) Titanate (Li4Ti5O12) 1-2 V 160 mA·h/g 0.16-0.32 kW·h/kg Si (Li4.4Si)[27] 0.5-1 V 4212 mA·h/g 2.106-4.212 kW·h/kg Ge (Li4.4Ge)[28] 0.7-1.2 V 1624 mA·h/g 1.137-1.949 kW·h/kg

  36. Lithium rechargeablebattery (2) The following equations are in units of moles, making it possible to use the coefficient x. Overdischarge supersaturates lithium cobalt oxide, leading to the production of lithium oxide Overcharge up to 5.2 Volts leads to the synthesis of cobalt(IV) oxide In a lithium-ion battery the lithium ions are transported to and from the cathode or anode, with the transition metal, cobalt (Co), in LixCoO2 being oxidized from Co3+ to Co4+ during charging, and reduced from Co4+ to Co3+ during discharge. http://en.wikipedia.org/wiki/Lithium-ion_battery

  37. Exercises 1 • What is the internal energy of 1 mole Ar at 0°C ? • What is the volume of 1 mole hydrogen gas at 25 °C ? • What is the entropy change in 1 mole hydrogen gas at standard conditions when increasing the volume to DV = 1 m3 ? • The equilibrium constant for acetic acid in water at 25°C is 4,76. What is Gibbs Free Energy at that temperature ? • Calculate the maximum electrical voltage for the DANIELL element when normal pressure and 10 °C ! • Calculate the maximum electrical voltage for a galvanic cell with Ni/Ni++//Zn++/Zn when normal pressure and 10 °C ! • Explain the difference between a galvanic and an electrolytic cell ! • What is the standard hydrogen potential ?

  38. Exercises 2 • What is the standard metal potential ? • How can you decide whether an ion will precipitated at a given electrode ? • What is the electrode potential for a silver electrode at 10°C when the Ag+ concentration is 1 mol ? • How can you calculate the amount of elementary metal to be formed on an electrode ? • How can you calculate the maximum energy which can be obtained from a battery • Explain the chemical potential ! • Explain the lead/acid battery ! • Explain the mercury battery !

  39. Web Links • http://en.wikipedia.org/wiki/Electrochemistry#Principles • http://www.jesuitnola.org/upload/clark/TeachResource.htm • http://goldbook.iupac.org/

  40. References • A. Burrows, A. Parsons , G. Price, J. Holman , G. Pilling; Chemistry: Introducing inorganic, organic and physical chemistry ; Oxford University Press 2009 • J. Hoinkins; E. Lindner; Chemie für Ingenieure; Verlag: Wiley-VCH Verlag GmbH & Co. KGaA, 2007 • P.W. Attkins; L. Jobnes; Chemie – einfach alles; Verlag: Wiley-VCH Verlag GmbH & Co. KGaA, 2006 • Römpp‘s Chemie Lexikon • DTV-Atlas zur Chemie

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