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

• ?Go -nFEcell
• n moles of electrons transferred
• F Faradays constant, 96,500 C/mol e-
• This gives an answer in J, since 1 J 1 CV
• Convert to kJ since thats what ?Go is usually
  expressed in
• Ecell (0.0592 / n) log K (Nernst equation)

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Practice with ?Go and K

• Calculate ?Go and K for the following cell
  Zn(s) Zn2(aq) Cu2(aq) Cu(s)
• ?Go -nFEcell
• Ecell (0.0592 / n) log K

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