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20_Electrochem.pdf

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Department
Chemistry
Course
CHEM 1E03
Professor
Weisner
Semester
Winter

Description
ELECTROCHEMISTRY Balancing redox reactions in acid or base solutions Electrochemical cells oxidation-reduction ⇒ transfer of e s Half cells - cell reaction, cell potential Standard reduction potentials direction of redox reactions calculation of cell potentials ∆G and cell potential criterion for reaction Nernst eqn - connection to equilibrium constant concentratcolls Applications batteries → lead acid corrosion 141 Redox rxns involve a transfer of electrons 2+ 2+ − eg. Zn(s) + Cu (aq) → Zn (aq) + Cu(s) ←2 e s transfered ♣ − Can we harness the transferred e s? Yes, but we must physically separate the oxidation and reduction half-reactions - An electrochemical cell uses the natural direction of a chemical reaction to generate useful electricity 142 Why does the current flow? There is an electric potential difference → a voltage - measured in volts Potential difference = work done on unit charge upon completing one passage around circuit 1 J = 1 V × 1 C 1 Joule work done 1 Coulomb of charge 1 Volt potential passed through circuit difference What gives rise to the potential difference of an electrochemical cell & determines its value? A chemical driving force - Zn atoms get oxidized while Cu 2+ gets reduced → a favorable chemical reaction → the resulting cell potential, E , depends on T & the cell concentration of aqueous species • if there are gases involved in the cell reaction, the gas partial pressures affect the cell potential 143 Under standard conditions ... → T = 25°C, all concentrations = 1 M & all gas partial pressures = 1 atm We measure the standard cell potential, E°cell note - positive eg. E °cellZn  Cu) = 1.10 V blue solution another eg. Cu(s) + 2 Ag (aq) → Cu (aq) + 2 Ag(s) whereas Ag(s) + Cu(NO ) (aq) → no rxn 3 2 The natural direction of the reaction produces a positive cell potential - negative cell potential simply means the reaction proceeds in the opposite direction E° (Cu  Ag) = 0.46 V cell ↑ ↑ anode cathode LEO GER 144 other egs. 2 Ag (aq) + H (g) → 2Ag(s) + 2 H (aq) + 2 E° = 0.80 V + −cell + − H 2g) → 2 H (aq) + 2 e 2 Ag + 2 e → 2 Ag(s) LEO GER → at anode → at cathode + 2+ Zn(s) + 2 H (aq) → 2 Zn (aq) + H (g) 2 E° cell 0.76 V Zn(s) → Zn 2+ + 2 e H2 +(aq) + 2 e → H (g)2 LEO GER → at anode → at cathode ♣ This is the standard hydrogen electrode. It is used - with 1 atm pressure H (g) - as a 2 reference to establish a scale of standard cell potentials. 145 Cell diagrams: salt bridge Pt(s) │ H 2g) │ H SO2(a4) ║ Ag NO (aq) 3 Ag(s) ↑ ↑ ↑ | denotes a phase change Zn(s) │ ZnSO (a4) ║ H S2 (a4) │ H (g) 2 Pt(s) anode reduction oxidization cathode Cu(s) │ Cu(NO ) 3 2) ║ Ag NO (aq)3│ Ag(s) Reaction = Reduction Half-reaction + Oxidation Half-reaction ⇒ Ecell= E red’n + Eoxid’n reduction half-cell oxidation half-cell potential potential under standard conditions ... E° = E° + E° cell red’n oxid’n ↑ ↑ dirtctly accessible 146 We arbitrarily set + − E°[H 2 2H + 2e ] = 0 V Recall 2 Ag + H → 2 Ag + 2 H + ⇒ H → 2 H + 2 e+ − Ag + + e → Ag anode cathode R E G O E L ° E cell + 0.80 V ← measured E°cell= E° red+ E°ox + − + − E°[Ag + e → Ag] + E°[H → 2 H2+ 2e ] + − ⇒ 0.80 V = E°[Ag + e → Ag] + 0 ⇒ standard reduction potentials can be tabulated in this way NOTE: E° ox= − E° red + − red half-cell ⇒ E°[2 H + 2e → H ] = 02V ← potential for H-electrode 147 + 2+ eg. Cu + 2 Ag → 2 Ag + Cu E°cell= 0.46 V E°cell= E° red+ E° ox + −
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