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Nernst Equation

Nernst Equation. G =. G o. + RT ln Q. G = -nF cell. G o = -nF o cell. standard. non-standard. nF.  o cell. - RT. ln Q.  cell =. Nernst Equation.  cell =  o cell - RT ln Q nF. anode. cathode. Cu (s) .

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Nernst Equation

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  1. Nernst Equation G = Go + RT ln Q G = -nFcell Go = -nFocell standard non-standard nF ocell - RT ln Q cell =

  2. Nernst Equation cell = ocell - RT ln Q nF anode cathode Cu(s) Cu2+(aq) (1M ) 4M Ag+(aq) (1M) Ag(s) a) oxidation oxidation b) reduction Cu2+ + 2e- Cu Cu(s) Cu2+ + 2e- 2( ) Ag+ + e- Ag(s) o= 0.34 V o= 0.80 V ored - oox= ocell= - 0.34 = 0.80 0.46 V

  3. [products]minitial [reactants]ninitial Nernst Equation cell = ocell - RT ln Q nF anode cathode Cu(s)Cu2+(aq) (4M) Ag+(aq) (1M) Ag(s) Cu (s)  Cu2+ + 2e- 2Ag+ + 2e- 2Ag (s) Q = = [Cu2+] = 4 = 4 [Ag+] 12 2

  4. Nernst Equation cell = ocell - RT ln Q nF anode cathode Cu(s)Cu2+(aq) (4M) Ag+(aq) (1M) Ag(s) Cu(s)  Cu2+ + 2e- 2Ag+ + 2e- 2Ag (8.314) ln 4 = 0.44 V cell = 0.46V - 2 (298) (96,500)

  5. Non-standard conditions Cu(s)Cu2+(aq) (4M) Ag+(aq) (1M) Ag(s) Cu(s)  Cu2+ + 2e- 2Ag+ + 2e- 2Ag G = -nF (0.46V) Go = -(2 mol e-) (96,500 C/mol e-) = -89 kJ (0.44V) G= -(2 mol e-) (96,500 C/mol e-) = -85 kJ

  6. V 1.62 x 10-10 ξ = 0.54 V Ksp AgCl = Ag+ + e-→ Ag (s) ξo = 0.80 V Ni2++ 2e-→ Ni (s) ξo = -0.23 V 38.17 ln Q = 0.54 = 1.03 -RT/nF ln Q Q = 3.78 x 1016 Ag (s) Ni (s) = [Ni2+] [Ag+]2 [Ag+] = 1.62 x 10-10 [Ag+][Cl-] Ksp = [AgCl] 1 x 10-3 M NiCl2 AgCl (s) 1.0 M HCl 1.0 M HCl

  7. Concentration cell Cu(s) Cu2+(aq) (0.5M)  Cu2+(aq) (2M) Cu(s)  Cu2+ + 2e- + 2e-  Cu(s) Cu(s) Cu2+ oxidation reduction 0 V ln Q cell = ocell - RT nF ocell= ored-oox = 0.34 - 0.34 = 0

  8. Concentration Cells Cu(s)Cu2+(aq) (0.5M)Cu2+(aq) (2M) Cu(s) Cu(s)  Cu2+ + 2e- Cu2+ + 2e- Cu(s) oxidation reduction = [products]minitial Q = [Cu2+] = 0.5 = 0.25 anode cathode [reactants]ninitial [Cu2+] 2 cathode -RT 2F cell = ocell -RT ln Q = 0 ln .25 = .02 V nF

  9. + - + Concentration Cells h QA P680 Mn + O2 2H2O 4H+ + 4e- pH = 3.0 cell = ocell - RT ln Q nF pH = 7.0 cell = ocell - 0.059 log Q n cell= Q = 10-7 / 10-3 G < 0 0.24 V G > 0 ADP + Pi ATP

  10. 2H+ + 2e- H2 o = 0.00 V o = 1.51 V MnO4- + 8H+ Mn2+ + 4 H2O +5e- 2MnO4- + 6H+ + 5H2 2Mn2+ + 8 H2O ocell= ored - oox = 1.51 - 0 = 1.51 V Go = -nFo = -1457 kJ = - (10)(96,500)(1.51) Go= -RT ln K = -1457 kJ K = e588 -RT -RT

  11. Balancing Redox Reactions MnO4- Mn2+ manganese is balanced balance oxygen using H2O MnO4- Mn2+ + 4H2O balance hydrogen using H+ MnO4- +  Mn2+ + 4 H2O 8H+ balance charge using e- MnO4- + 8H+ Mn2+ + 4 H2O +5e-

  12. Oxidation States MnO4- + 8H+ Mn2+ + 4 H2O +5e- Oxidation state of oxygen = 2- in compounds 0 in O2 Oxidation state of hydrogen = 1+ in compounds 0 in H2 Mn = 7+ MnO4- 4 x (-2) + 1(Mn) = -1 Mn2+ Mn = 2+ Mn7+ + 5 e- Mn2+

  13. 6 MnO4- + 8H+ Mn2+ + 4 H2O +5e- o = 1.51 V 2H+ + 2e- H2 o = 0.00 V anode cathode Pt(s) Mn2+(aq)(1M) MnO4-(aq)(1M), ,H+(1M) Pt(s) H2(g)(1atm) H+(aq)(1M)  ( )5 H22H++2e- ( )2 MnO4-+8H++5e-Mn2++4 H2O oxidation reaction reduction reaction 2MnO4- + 16H+ +10e- + 5H2 2Mn2+ + 8 H2O + 10H+ + 10e- 2MnO4- + 6H+ + 5H2 2Mn2+ + 8 H2O

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