Download - Chapter 20: Electrochemistry
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Chapter 20: Electrochemistry
Chemistry 1062: Principles of Chemistry II
Andy Aspaas, Instructor
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Oxidation-Reduction reactions
• Oxidation-reduction (redox) reaction: transfer of electrons from one species to another
• H3O+ becomes simply H+ when dealing with redox reactions to simplify balancing
– (still the same species, just different notation)• Skeleton oxidation-reduction equation: involves only the
species being oxidized and reduced. – Write oxidation numbers above each species. – No spectator ions, no balancing
• Half reaction: shows only one oxidation OR one reduction– Most redox reactions are split into an oxidation half-
reaction and a reduction half-reaction– LEO, GER
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Balancing redox equations in acidic solutions
• For each half reaction…
– Balance everything except H or O
– Balance O by adding H2O to one side
– Balance H by adding H+ to one side– Balance charge by adding e- to one side
• Multiply each half reaction by a factor so that the electrons cancel when the two half reactions are added together (e- cannot appear in the final equation)
• Add the reactions, cancel anything that appears on the left and right, and simplify the coefficients to the smallest integers
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Practice balancing acidic redox reactions
• Balance I2(s) + NO3-(aq) IO3
-(aq) + NO2(g) in acidic solution
• Half reactions
• Cancel electrons
• Add half-reactions
• Simplify
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Voltaic cells
• A voltaic cell consists of two half-cells
– Each half-cell contains a metal rod dipped in a solution containing that metal ion
– Anode: a species is being oxidized– Cathode: a species is being reduced
• Cell reaction: redox reaction for entire voltaic cell
• Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Zn(s) Zn2+(aq): oxidation half-reaction, anode
Cu2+(aq) Cu(s): reduction half-reaction, cathode
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Cell notation
• Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
• Cell notation: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
– Anode || Cathode
• Write half reactions and cell reactions for the following cell:
Tl(s) | Tl+(s) || Sn2+(aq) | Sn(s)
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emf, Standard Cell emf, Standard electrode potential
• Electromotive force, emf, Ecell = electrical pressure across the conductors of an electrochemical cell
– Unit: Volt, V– Measure of the driving force of a cell reaction
• Standard cell emf, Eocell = solutes are 1 M, gases are 1 atm,
temperature is 25 oC• Standard electrode potenital
– By convention, the standard hydrogen electrode has an emf of 0 V
– All reactions shown as reductions• Ecell = Ecathode – Eanode
• Ecell is positive for spontaneous reactions as written
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Practice calculating Ecell
• Using standard potentials, calculate Ecell for Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
• Standard cell potentials are an intensive property
– Do not depend on quantity!– If you have to multiply a half-reaction to cancel
electrons, do not multiply the Eo for that half-reaction
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Free energy and K from Ecell
• ΔGo = -nFEcell
n = moles of electrons transferred
F = Faraday’s constant, 96,500 C/mol e-
• This gives an answer in J, since 1 J = 1 C·V
• Convert to kJ since that’s what ΔGo is usually expressed in
• Ecell = (0.0592 / n) log K (Nernst equation)
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Practice with ΔGo and K
• Calculate ΔGo and K for the following cell: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
• ΔGo = -nFEcell
• Ecell = (0.0592 / n) log K
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