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ElectrochemistryElectrochemistry Prof. Dr. Sabine Prys by ps.

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1 ElectrochemistryElectrochemistry Prof. Dr. Sabine Prys by ps

2 Contents Chemical Reactions Chemical Reaction Kinetics Chemical Reaction Thermodynamics Chemical Solutions Electrochemical Reactions Galvanic Cells & Batteries Standard Potential Specific Electrodes NERNST’s equation Corrosion

3 1.0 Chemical Reactions redox reaction –Anodic oxidation Ag  Ag + + e - Fe 2+  Fe 3+ + e - –Cathodic reduction Ag + + e -  Ag Fe 3+ + e -  Fe 2+ stochiometrical yields Law of mass conservation !

4 1.0.1 Driving Forces Kinetics (chemical process rates) Thermodynamics (chemical process energy flow) reaction energy change: standard enthalpy*  H 298  H =  U + p.  V + v  p  U =  Q +  W chemical reactions in open vessels:  H ≈  U * enthalpy at standard conditions: T = 298 K, p= 1,013 bar HenthalpyUinternal energyQflow of heatWWork TtemperaturePpressureVvolume

5 2.0 Chemical Reaction Kinetics reversible chemical reaction e.g. in aqueous solution [A],[B] =reactant concentrationa,b = mols of corresponding reactants [C],[D]= product concentrationc,d = mols of corresponding products k  = velocity of reaction k  = velocity of back reaction K = equilibrium constant

6 (reversible) chemical reaction K x =equilibrium constant x B =equilibrium value either of pressure, fugacity, amount concentration, amount fraction, molality, relative activity or reciprocal absolute activity defining the pressure based, fugacity based, concentration based, amount fraction based, molality based, relative activity based or standard equilibrium constant B =stoichiometric number of a reactant (negative) or product (positive) 2.1 Equilibrium Constant

7 2.1.1 Reaction Velocity 1 st order kinetics velocity of reaction [t] [c]

8 2.1.2 Mass-Law Effect chemical potential of reaction (  ) = chemical potential of back reaction (  ) equilibrium constant K depends on temperature chemical equilibrium: concentrations are not independent

9 2.1.3 Example: Neutralisation Reaction Acid Base Reaction : acis + base = salt + water But also: + +

10 3.0 Chemical Reaction Thermodynamics U Internal energy – total energy contained by a one component system Uinternal energyE kin kinetic energyE pot potential energy Qflow of heatWWorkTtemperature SentropyppressureVvolume one component thermodynamic system: e.g. Ideal gas

11 3.1 „Ideal“ Gas one component thermodynamic system: e.g. Ideal gas e.g. N 2, H 2, O 2, He, Ne,.... V large, p small System contains only one particle type System contains statistically sufficient particles System contains point particles System contains non interacting particles System contains randomly moving particles All particle collisions are elastic

12 3.1 „Ideal“ Gas Law p= pressure [Pa] = [N/m 2 ] V= volume[m 3 ] n= number of mols[mol -1 ] R= gas constant[J/(mol+K)] T= Kelvin temperatur[K] gas constant:R = 8,3145 J.K -1.mol -1 one component thermodynamic system: e.g. gas

13 3.2 The BOLTZMANN Constant R = gas constantN A = Avogadro constant The BOLTZMANN constant is the physical constant relating energy at the individual particle level with temperature observed at the collective or bulk level

14 3.3 Normal & Standard Conditions normal conditions: T = 0°C = K normal conditions: normal pressure p = 1 atm = 101,325 kPa = 1013,25 mbar normal temperature T = 0°C = K standard conditions: standard conditions: T = 25°C = K standard pressure p = 1 atm = 101,325 kPa = 1013,25 mbar standard temperature T = 25°C = K Such definitions can vary according to different sources: IUPAC, NIST, …

15 3.4 Enthalpy EnthalpyEnthalpy 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. an extensive quantityEnthalpy is a thermodynamic potential, a state function and an extensive quantity (i.e. depending on amount material). HenthalpyUinternal energy PpressureVvolume

16 3.4.1 Exercise: gas volume Calculate the volume of 1 mol of an ideal gas ! –under normal conditions –under standard conditions

17 3.4.2 Exercise: internal energy 1.What is the internal energy for 1 Mol He at 20°C ?

18 3.4.3 Exercise: internal energy change 1.What is the internal energy change for 1 Mol He at 20 ° C, p = 1000 Pa and  Q = 0 when expanding the volume from 1 m 3 to 2 m 3 ?

19 3.4.4 Exercise: enthalpy 1.What is the enthalpy for 1 Mol He at 20 ° and standard pressure ?

20 3.5 Entropy reversible processQuantity the change in which is equal to the heat brought to the system in a reversible process at constant temperature divided by that temperature. Entropy is zero for an ideally ordered crystal at 0 K. In statistical thermodynamics where k is the Boltzmann constant and W the number of possible arrangements of the system the entropy is  S = k. lnW

21 3.5.1 Exercise: entropy change none expanding the volume 1.What is the entropy change for 1 Mol He at 20 ° C, p = 1000 Pa and  Q = 0 when expanding the volume from 1 m 3 to 2 m 3 ?

22 3.5.1 Exercise: entropy change 1.What is the entropy change for 1 Mol He at 20 ° C when the flow of heat  Q = 1000 J ?

23 3.6 Chemical Reaction Energy Internal energy of a one component thermodynamic system Internal energy of multi-component systems Uinternal energy E kin kinetic energy E pot potential energy Qflow of heat WWork Ttemperature Sentropy Ppressure Vvolume N i component i µ i chemical potential GGibbs’ energy multi component thermodynamic system: e.g. Chemical reaction

24 3.7 Enthalpies of Multi Component Systems Enthalpy of reaction Enthalpy of formation Enthalpy of combustion Enthalpy of hydrogenation Enthalpy of atomization Enthalpy of neutralization Enthalpy of solution Enthalpy of hydration Enthalpy of fusion Enthalpy of vaporization Enthalpy of sublimation Lattice enthalpy …..

25 3.7.1 Endothermic Reaction Endothermic A reaction for which the overall standard enthalpy change  H 298 >0 e.g. t E activation energy +  H

26 3.7.2 Exothermic Reaction Exothermic A reaction for which the overall standard enthalpy change  H 298 < 0 e.g. t E activation energy -  H

27 3.8 GIBBS’ Free Energy G Useful energy, or energy available to do work G G = free energyH= GIBBS’ energy (enthalpy) U = internal energyT= Kelvin temperature S= entropyp= pressure V= volume T  S is the energy not available for doing work

28 3.9 Spontaneity of Redox Reactions  H  SSpontaneity Exothermic  H 0+  G < 0 Exothermic  H < 0Decrease  S < 0+ if|T  S| < |  H| Endothermic  H > 0Increase  S > 0+ ifT  S >  H Endothermic  H > 0Decrease  S 0

29 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  G = 0 At equilibrium the sum of the chemical potentials of the reactants equals that of the products, so that:  G =  G RT. lnK = 0   G 298 = - RT. lnK The equilibrium constant K is given by the mass-law effect.

30 Examples 6 C (s) + 3 H 2 (g)  C 6 H 6 (l)  G 298 =+ 124 kJ C 6 H 6 (l)  6 C (s) + 3 H 2 (g)  G 298 =- 124 kJ 4 NH 3 (g) + 5 O 2 (g)  4 NO (g) + 6 H 2 O  G 298 =- 959,42 kJ

31 Reduction of SiO 2 reduction = electron acceptance production of metallurgical grade silicon Electric arc furnace with carbon electrodes 14 kWh / kg Si !!!

32 oxidation = electron donation SiHCl 3 Synthesis oxidation of silicon to SiHCl 3

33 3.11 Chemical Potential 1 G= Gibbs free energyn + = moles of cations n b = moles of type bn - = moles of anions n j = moles of type jµ b = chemical potential of type b T= Kelvin temperatureµ + = chemical potential of cations p= pressureµ - = chemical potential of anions

34 Standard Chemical Potential µ= chemical potential µ 0 = standard chemical potential R= gas constant T= Kelvin temperature a= activity of an ion type chemical potential µ i : change in the energy of the system when an additional constituent particle i is introduced, with the entropy and volume held fixed

35 Chemical Potential 2

36 3.12 Maximum Work W max = 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

37 Examples What is the vaporisation enthalpy of water at 298 K ? What is the heating energy of 1 m 3 methane ? What is the combustion energy of 1 m 3 methane ?

38 4.0 Chemical Solutions SuspensionSuspension (particle diameters cm ) solid particles in homogeneous fluid ColloidColloid (particle diameters cm ) microscopically dispersed particles in another substance SolutionSolution(particle diameters cm ) Homogeneous phase with at least 2 components: solvent and solute »gas in liquide.g. O 2 in H 2 O »gas in solide.g. H 2 in palladium »liquid in liquide.g. petroleum »solid in liquide.g. NaCl in H 2 O »  electrolytes in water

39 4.1 Solvation & Dissolution Solvation IUPAC: IUPAC: any stabilizing interaction of a solute … and the solvent … Such interactions generally involve electrostatic forces and van-der-Waals forces, as well as chemically more specific effects such as hydrogen bond formation. Dissolution IUPAC: IUPAC: mixing of two phases with the formation of one new homogeneous phase (i.e. the solution) kinetic process, quantified by its rate

40 4.2 Solubility IUPAC: The analytical composition of a saturated solution, expressed in terms of the proportion of a designated solute in a designated solvent, is the solubility of that solute. The solubility may be expressed as a concentration, molality, mole fraction, mole ratio, etc. quantifies the dynamic equilibrium state achieved when the rate of dissolution equals the rate of precipitation – solubility of salts: lattice energy & solvation enthalpy – ideal dilute solution

41 4.2.1 Lattice Energy CompoundLattice Energy [kJ/mol] LiF -1019,9 LiCl- 838,6 LiBr- 798,8 LiI- 742,3 NaCl- 766,2 KCl RbCl- 665,7 CsCl- 649,4 CompoundLattice Energy [kJ/mol] MgO3931,4 CaO3479,2 SrO3207,1 BaO3043,8 Al 2 O ,2 CaF ,6 CaCl ,8 CaBr ,4 CaI ,7 Elektrolytgleichgewichte; Ackermann et. Al

42 4.3 Aqueous Solutions Solvatation: Solvatation: lattice disintegration break of hydrogen bonds  „hole“ formation formation of ion dipole interactions Compound  H solution_∞ Compound  H solution _∞ Compound  H solution _∞ LiF + 4,6NaF+ 2,5KF- 17,2 LiCl- 36,0NaCl+ 5,4KCl+ 18,4 LiBr- 46,5NaBr+ 0,8KBr+ 21,4 LiI- 62,0NaI - 5,9KI+ 21,4 Elektrolytgleichgewichte; Ackermann et. Al

43 4.4.1 Hydratisation Enthalpies at 25°C Ion  H Hydr Ion  H Hydr Ion  H Hydr [J/mol] [J/mol] [J/mol] H+-1130,4NH ,7OH-- 318,2 Li+- 552,7Mg ,3F-- 410,3 Na+- 443,8Ca ,5Cl-- 330,8 K+- 401,9Br-- 297,3 Rb+- 272,1I-- 259,6 Cs+- 334,9 Increasing ion radius  less energy gain ! Elektrolytgleichgewichte; Ackermann et. Al

44 4.5 Saturated Solutions IUPAC: temperature pressure. A solution which has the same concentration of a solute as one that is in equilibrium with undissolved solute at specified values of the temperature and pressure. loading capacity (maximum loading, saturation capacity saturation loading) The maximum concentration of solute(s) that a solvent can contain under specified conditions. solute

45 4.5.1 Gas Solutions IUPAC: The partial pressure (fugacity) of a solute (B) in a solution is directly proportional to the rational chemical activity (a x ) of the solute: Henry's law: p B = partial pressure a x,B = chemical activity, a x,B ∞ = solubility coefficient

46 4.5.2 Solubility Coefficients of Gases in H 2 O Temp.[°C] O 2 N 2 CO 2 [g(Gas)/kg(H2O)/bar] 0 0,0676 0,0281 3, ,0526 0,0226 2, ,0428 0,0190 1, ,0364 0,0166 1, ,0291 0,0137 0, ,0258 0,0129 0, ,0246 0,

47 4.5.3 Solubility of Salts in H 2 O Compound 20 °C 80 °C NaCl 26,5 27,5 KCl 25,5 33,6 NH 4 Cl 27,0 40,0 K 2 SO 4 10,0 17,5 CaSO 4 0,199 0,10 Ca(OH) 2 0,17 0,087 CaCO 3 0,0015 0,002 (100 °C) ZnCl 2 78,7 84,5 In mass % Solubility depends on: ion radii, solvent dielectricity constant

48 4.5.4 Densities of NaCl Solutions at 20°C Concentration [mass%] Density [kg/m³] Concentration [mass%] Density [kg/m³]

49 4.5.6 Solubility Curves for Unhydrated Salts figure reference: Elektrolytgleichgewichte; Ackermann et. Al Often heat production when dissolving non hydrated salts e.g CaCl 2, Na 2 CO 3

50 4.5.7 Solubility Curves for Hydrated Salts figure reference: Elektrolytgleichgewichte; Ackermann et. Al Often heat consumation when dissolving hydrated salts e.g CaCl 2 * 6 H 2 0, Na 2 CO 3 * 10 H 2 O Change of solid structure & composition

51 4.5.8 Solubility Equilibrium IUPAC: IUPAC: solubility product The product of the ion activities raised to appropriate powers of an ionic solute in its saturated solution expressed with due reference to the dissociation equilibria involved and the ions present.

52 4.5.9 Low & High Solubility at 25°C Compound K sol AgCl1, mol 2 /l 2 AgBr5, mol 2 /l 2 AgI8, mol 2 /l 2 Ag 2 CrO 4 1, mol 2 /l 2 Ag 2 S5, mol 2 /l 2 Fe(OH) 2 1, mol 2 /l 2 Fe(OH) 3 6, mol 2 /l 2 FeS4, mol 2 /l 2  low solubility Compound K sol NaHCO 3 1, mol 2 /l 2 MgCO 3 1, mol 2 /l 2 PbCl 2 1, mol 2 /l 2 BaF 2 1, mol 2 /l 2 Ba(OH) 2 5, mol 2 /l 2 CaSO 4 2, mol 2 /l 2 Ag 2 SO 4 1, mol 2 /l 2  high solubility

53 4.6 Ion activity High ion concentrations in aqueous solutions  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/ m/ m/ m/0.809

54 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

55 4.8 Gels

56 4.9 Electrolytes Electrolyte: solution which conducts electrical current Hydrated H 3 O + Hydrated OH -

57 4.9.1 Electrical Conductivity in Solutions Electrolytes solutions which support ion transport salts in aqueous solutions, e.g. KCl, ZnSO 4, CuCl 2, etc. molten salts Conductivity L Conductivity L (resistance R) bad electrolyte: distilled water: µS/cm at 25 °C. cathode cat ions anode anions _ H2OH2O

58 4.9.2 Specific Conductivity absolute electrolyte conductivity R = solution resistance specific electrolyte conductivity A = electrode surface, l = electrode distance

59 4.9.3 Example: Proton Migration Grotthuss Diffusion structural defect migration mesomeric structures between H 9 O 4 + and H 5 O 2 +,

60 5.0 Electrochemical Cells Cl 2 H2H2 + - H2H2 + - electrolytic cell galvanic cell 2 HCl (aq)  H 2 (g) + Cl 2 (g) H 2 (g) + Cl 2 (g)  2 HCl (aq) electrical energy  chemical energychemical energy  electrical energy electrical energy  chemical energychemical energy  electrical energy

61 5.1 Electrolysis electrolysis: decomposing materials by electric current H 2 SO 4 + 2H 2 O  2H 3 O + + SO 4 2- water electrolysis cathodic reduction 4H 3 O + + 4e -  2 H 2  + 2 H 2 O anodic oxidation 4 OH -  2 H 2 O + O 2  + 4 e - total 2 H 2 O (l)  2 H 2 (g) + O 2 (g) H H 2 SO 4 1:10 electrods battery ca. 15 V

62 5.1.1 Electrochemical Equivalent Q = electric charge in C n =yield in mol F=Faraday‘s constant = 96485,309 As / mol E c =electrochemical equivalent M=ion weight z=ion charge N L =Lohschmidt‘s number e=elementary charge

63 5.1.2 Faraday‘s Laws m a,m b = mass yield in g for material a / b M a,M b =molecular weight for material a / b z a, z b =chemical valency for material a / b m = E c. Q = E c.I. t m = mass yield in g E c =electrochemical equivalent Q =electric charges in Coulomb I=current strength t=electrolysis time

64 5.2 Galvanic Elements Daniell Element: 2 galvanic half cells + bridge Zn / ZnSO 4 // CuSO 4 / Cu electrode reactions Zn (cathode)  Zn e - Cu e -  Cu (anode) Zn metal in ionic solution Cu ions in Cu metal electrical current results from different oxidation affinities voltmeter ca 1,1 V ZnCu 1 m ZnSO 4 1 m CuSO 4 diaphragma (pottery) bridge containing KCl solution

65 5.2.1 Half Cells & Bridge Daniell Element‘s Half Cells: Zinc electrode (anode): Zn(s) → Zn 2+ (aq) + 2 e – Copper electrode (cathode): Cu+ (aq) + 2 e – → Cu(s) simple battery: flow of electrical current from the anode to the cathode Bridge completion of the electric circuit: salt bridge ionic conduction path between the anode and cathode electrolytes –porous plug that allows ion flow without electrolyte mixing –salt bridge from electrolyte saturated gel in an inverted U-tube reduction oxidation

66 5.3 Electrical Potential cell diagram (path of the electrons in the electrochemical cell) Daniell cell: Zn(s) | Zn 2+ (1M) || Cu 2+ (1M) | Cu(s) reduced form  oxidised form electrical potential between anode and cathode  electrochemical cell voltage  electromotive force  emf 1 m CuSO 4 voltmeter ca 1,1 V ZnCu 1 m ZnSO 4 diaphragm (pottery) bridge containing KCl solution

67 5.4 Standard Hydrogen Electrode standard hydrogen electrode (SHE) =reference potential =E 0 = 0 V H 2  2H + + 2e - p = 1,01325 bar T = 25°c a(H + ) = 1 mol / l c(H + ) = 1,235 mol / l (HCl) Pt electrode H 2 gas

68 Pt electrode H 2 gas Metal Standard Potentials standard hydrogen electrode = reference potential E 0 = 0 V metal electrode / metal salt solution at standard conditions standard metal potential = standard metal potential M  M z+ + ze - p = 1,01325 bar T = 25°c c(M z+ ) = 1 mol / l pH < 6 precipitation prevention

69 5.5.1 Metal Standard Potential Tables pH-dependant

70 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 Calvanic Corrosion Potential Chart cathode least noble corroded metals strong oxidation affinity negative oxidation potential anode most noble protected metals weak oxidation affinity positive oxidation potential

71 5.6 Standard Cell Potential E° cell = E° red (cathode) – E° red (anode) = E° red (cathode) + E° ox (anode) cell diagram: Pt(s) | H 2 (1 atm) | H + (1 M) || Cu 2+ (1 M) | Cu(s) Standard cell potential E° cell = E° red (cathode) – E° red (anode) E° cell = E°(Cu 2+ /Cu) – E°(H + /H 2 ) E° cell = 0,34 V - = V = 0,34

72 5.6.1 Free Energy in Electrochemical Cells operation of electrochemical cells: chemical energy  electrical energy E electr = Q·E cell = n· F·E cell W max = W electrical = – n·F·E° cell free energy: maximum amount of work that can be extracted ΔG = – n·F·E cell

73 5.6.2 Maximum Work in Electrochemical Cells Δ G < 0 spontaneous electrochemical reaction: –electric current electrochemical cells batteries fuel cells Δ G < 0 electrolysis Δ G = Δ G° + RT·lnK nF Δ E = nF Δ E° – RT lnK n = the number of electrons/mole product F = the Faraday constant (coulombs/mole) ΔE = cell potential

74 5.6.3 Exercise: Gibbs Free Energy What happens if Δ G = 0

75 5.7 NERNST‘s Equation 1 electrode potential dependency on temperature and concentration E = measured cell potential E 0 = 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

76 5.7.1 NERNST’s Equation 2 1.type electrode ( = metal electrode in metal salt solution) [red] = const E = measured cell potential E 0 = 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

77 5.7.2 Exercise: Maximum Electrical Voltage 1.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  =+ 1,1 V 2.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  =+ 0,53 V  =+ 0,53 V

78 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 ?

79 5.8 Ag / AgCl Electrode 2. type electrode = metal electrode in saturated metal salt solution = electrode with constant potential (no concentration changes) T = 25 °C: 1 m KCl E 0 = + 0,220 V sat. KCl E 0 = + 0,1958 V Ag AgCl sat K + Ag + Cl -

80 5.8.1 Concentration Cells Cu(s) | Cu 2+ (0.05 M) || Cu 2+ (2.0 M) | Cu(s) half cell reactions : oxidation: Cu(s) → Cu 2+ (0.05 M) + 2 e– reduction: Cu 2+ (2.0 M) + 2 e – → Cu(s) overall reaction: Cu 2+ (2.0 M) → Cu 2+ (0.05 M) cell's emf : E = E°- ( \2) log [0,05/2] = V E° = 0, (electrodes and ions are the same in both half-cells)

81 5.10 pH Electrode measurement electrode = glas electrode reference electrode = silver chlorid flask cork KCl - solution Ag-stick AgCl solid KCl- solution glas membran = solid electrolyt buffer solution

82 5.11 Battery

83 5.12 Electrochemical cell types

84 5.13 Primary electrochemical cells

85 5.14 Secondary electrochemical cells

86 5.15 Dry Cells 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 NH 4 Cl + ZnCl 2 mixed with starch Anode: Zn(s) → Zn 2+ (aq) + 2 e – Cathode: 2 NH 4+ (aq) + 2 MnO 2 (s) + 2 e – → Mn 2 O 3 (s) + 2 NH 3(aq) + H2O(l) Overall reaction: Zn(s) + 2 NH 4+ (aq) + 2 MnO 2 (s) → Zn 2+ (aq) + Mn 2 O 3 (s) + 2 NH 3(aq) + H 2 O(l) E = ~ 1.5 V moist electrolyte paste

87 5.16 Zn Battery Graphics:

88 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) + H 2 O(l) + 2 e – Cathode:HgO(s) + H 2 O(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:

89 5.18 Lead-Acid battery six identical cells assembled in series (6 x 2V ) = 12 V lead anode lead dioxide cathode Electrolyte sulfuric acid Anode: Pb(s) + SO 4 2– (aq) → PbSO 4 (s) + 2 e – Cathode: PbO 2 (s) + 4 H + (aq) + SO 4 2– (aq) + 2 e– → PbSO 4 (s) + 2 H 2 O(l) Overall reaction: Pb(s) + PbO 2 (s) + 4 H + (aq) + 2 SO 4 2– (aq) → 2 PbSO 4 (s) + 2 H 2 O(l) Rechargeable (external voltage  electrolysis of the products)

90 5.19 Lithium rechargeable battery (1) Positive electrodes Electrode material Average potential difference Specific capacity Specific energy LiCoO2 3.7 V 140 mA·h/g kW·h/kg LiMn2O4 4.0 V 100 mA·h/g kW·h/kg LiNiO2 3.5 V 180 mA·h/g kW·h/kg LiFePO4 3.3 V 150 mA·h/g kW·h/kg Li2FePO4F 3.6 V 115 mA·h/g kW·h/kg LiCo1/3Ni1/3Mn1/3O2 3.6 V 160 mA·h/g kW·h/kg Li(LiaNixMnyCoz)O2 4.2 V 220 mA·h/g kW·h/kg Negative electrodes Graphite (LiC6) V 372 mA·h/g kW·h/kg Hard Carbon (LiC6) Titanate (Li4Ti5O12) 1-2 V 160 mA·h/g kW·h/kg Si (Li4.4Si)[27] V 4212 mA·h/g kW·h/kg Ge (Li4.4Ge)[28] V 1624 mA·h/g kW·h/kg

91 Lithium rechargeable battery (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.

92 5.20 Flow battery

93 5.21 Fuel cells

94 6.0 Practical Potential Charts 94

95 6.1 Galvanic series (most noble at top)

96 6.1.1 Leiterplattenherstellung 96

97 7.0 Corrosion less noble metal =anodeM 1  M 1 z+ + z e - noble metal = cathodeM 2 z+ + z e -  M 2 acidic electrolyte2 H e -  H 2  less acidic electrolyte2 H 2 O + O e -  4 OH - different local electrolyte concentrations corrosion from different local oxygen content

98 Fe Corrosion in a Drop of Water drop of water corrosion zone Fe ++ few O 2 OH - corrosion zone more O 2 Fe OH -  Fe(OH) 2  Fe(OH) 3 OH -

99 7.2 Iron corrosion

100 7.3 Corrosion of common metals

101 7.4 Prevention of corrosion

102 7.5 Coating

103 7.6 Sacrificial anodes

104 7.7 Electrolysis

105 7.8 Electrolysis of molten sodium chloride

106 7.9 Electrolysis of aqueous solutions

107 7.10 Electrolysis of a solution of sodium chloride

108 7.11 Applications

109 Exercises 1 1.What is the internal energy of 1 mole Ar at 0°C ? 2.What is the volume of 1 mole hydrogen gas at 25 °C ? 3.What is the entropy change in 1 mole hydrogen gas at standard conditions when increasing the volume to  V = 1 m 3 ? 4.The equilibrium constant for acetic acid in water at 25°C is 4,76. What is Gibbs Free Energy at that temperature ? 5.Calculate the maximum electrical voltage for the DANIELL element when normal pressure and 10 °C ! 6.Calculate the maximum electrical voltage for a galvanic cell with Ni/Ni ++ //Zn ++ /Zn when normal pressure and 10 °C ! 7.Explain the difference between a galvanic and an electrolytic cell ! 8.What is the standard hydrogen potential ?

110 Exercises 2 9.What is the standard metal potential ? 10.How can you decide whether an ion will precipitated at a given electrode ? 11.What is the electrode potential for a silver electrode at 10°C when the Ag+ concentration is 1 mol ? 12.How can you calculate the amount of elementary metal to be formed on an electrode ? 13.How can you calculate the maximum energy which can be obtained from a battery 14.Explain the chemical potential ! 16.Explain the lead/acid battery ! 17.Explain the mercury battery !

111 Exercises 3 16.Explain the pH electrode ! 17.What is the calomel electrode? 18.What is the standard metal potential ? 19.Explain the fuel cell ! 20.Explain corrosion processes ia a water drop !

112 Web Links

113 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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