Electrode Potentials

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Chapter 14

• Electrode Potentials

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14-1 Redox Chemistry Electricity
Oxidation a loss of electrons to an oxidizing
agent Reduction a gain of electrons from a
reducing agent Reduction-Oxidation reaction
(redox reaction) ex Half-reactions r
e Fe3 e- ? Fe2 ox V2 ? V3 e-
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1. Chemistry Electricity

• Electrochemistry the study of the interchange of
  chemical electrical energy.
• Electric charge (q) is measured in coulombs(C).
• The magnitude of the charge of a single electron
  (or proton) is 1.60210-19 C. A mole of electrons
  therefore has a charge of (1.60210-19
  C)(6.0221023 /mol) 9.649104 C/mol, which is
  called the Faraday constant, F.
• Example at p.310

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2. Electric current is proportional to the rate
of a redox reaction

• I (ampere A) electric current
• a flow of 1 coulomb per second 1C/s
• Example at p. 310
• Sn4 2e- ? Sn2
• at a constant rate of 4.24 mmole/h.
• How much current flows into the solution?

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3. Voltage Electrical Work
The difference in electric potential between two
points measures the work that is needed (or can
be done) when electrons move from one point to
another.

wire q I E
hose H2O VH2O PH2O
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Ask yourself at p.312

• Consider the redox reaction

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14.2 Galvanic Cells

• Chemical reaction spontaneously occurs to produce
  electrical energy.
• Ex lead storage battery
• When the oxidizing agent reducing agent are
  physically separated, e transfer through an
  external wire.
• ? generates electricity.

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A cell in action
Electrodes the redox rxn occur anode
oxidation occur cathode reduction occur Salt
bridge connect two solns. External wire
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Cell representation Line Notation
Example Interpreting Line Diagrams of Cells
Figure 14-4 Another galvanic cell.
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14-3 Standard Potentials

• Cell potential ( Ecell)
• The voltage difference between the electrodes.
• ? electromotive force (emf)
• can be measured by voltmeter.
• emf of a cell depends on
• The nature of the electrodes ions
• Temp.

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14-3 Standard Potentials

• S.H.E. (standard hydrogen electrode )
• It is impossible to measure Ecell of a half-rxn
  directly, ?need a reference rxn.
• standard hydrogen electrode

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The standard reduction potential (E0) for each
half-cell is measured by an experiment shown in
idealized form in Fig.14-6.
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Table 14-1 Appendix C
(?1953, the 17th IUPAC meeting??????????????? )
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Standard Reduction Potentials for reaction
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Standard Reduction Potentials for reaction

• rxn

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Formal potential

• AgCl (s) e- ? Ag (s) Cl-
• 0.222 V
• 0.197 V in saturated KCl (formal potentional)
• E0 0.222V
• S.H.E. Cl- (aq, 1M) AgCl (s) Ag(s)
• E0 (formal potential) 0.197 V (in saturated
  KCl)
• S.H.E. KCl (aq, saturated) AgCl (s) Ag(s)

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Formal Potential

• Ex Ce4 e- ? Ce3 E1.6V
• with HA- E?1.61V
• Formal potential (E)
• The potential for a cell containing a reagent
  ?1M.
• Ex Ce4/Ce3 in 1M HCl E1.28V

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14-4 The Nernst Equation
The net driving force for a reaction is expressed
by the Nernst eqn.
Nernst Eqn for a Half-Reaction
where

• E is the reduction potential at the specified
  concentrations
• n the number of electrons involved in the
  half-reaction
• R gas constant (8.3143 V coul deg-1mol-1)
• T absolute temperature
• F Faraday constant (96,487 coul eq-1) at 25C
  ? 2.3026RT/F0.05916

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Nernst equation for a half-reaction at 25ºC

• E E0 when A B 1M
• Q (Reaction quotient ) 1 ? E E0
• Where, Q Bb / Aa

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C Ecell

• standard conditions C1M
• what if C?1M?
• (ex)
• Al32.0M, Mn21.0M Ecelllt0.48V
• Al31.0M, Mn23.0M Ecellgt0.48V

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Dependence of potential on pH
Many redox reactions involved protons, and their
potentials are influenced greatly by pH.
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• Nernst Equation for a Complete Reaction
• 1. Write reduction half-reactions for both
  half-cells and find E0 for each in Appendix C.
• 2. Write Nernst equation for the half-reaction in
  the right half-cell.
• 3. Write Nernst equation for the half-reaction in
  the left half-cell.
• 4. Fine the net cell voltage by subtraction
  EE- E-.
• 5. To write a balanced net cell reaction.

P.321
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Nernst Equation for a complete reaction
Example at p. 321 Rxn 2Ag (aq) Cd (s) ? Ag
(s) Cd 2 (aq) 2Ag 2e- ? Ag (s) E0
0.799 Cd 2 2e- ? Cd (s) E0- -0.402
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• Electrons Flow Toward More Positive Potential
• Electrons always flow from left to right in a
  diagram like Figure 14-7.

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14-5 E0 and the Equilibrium Constant
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14-5 E0 and the Equilibrium Constant
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At equilibrium

• E 0 and Q K
• E0 gt 0 K gt 1,
• E0 lt 0, K lt 1

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Ex

• One beaker contains a solution of 0.020 M KMnO4,
  0.005 M MnSO4, and 0.500 M H2SO4 and a second
  beaker contains 0.150 M FeSO4 and 0.0015 M Fe2
  (SO4)3. The 2 beakers are connected by a salt
  bridge and Pt electrodes are placed one in each.
  The electrodes are connected via a wire with a
  voltmeter in between.
• What would be the potential of each half-cell (a)
  before reaction and (b) after reaction?
• What would be the measured cell voltage (c) at
  the start of the reaction and (d) after the
  reaction reaches eq.?
• Assume H2SO4 to be completely ionized and equal
  volumes in each beaker.

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• Ans
• 5Fe2 MnO4- 8H ? 5Fe3 Mn2 4H2O
• Pt Fe2(0.15 M), Fe3(0.003 M)MnO4-(0.02 M),
  Mn2(0.005 M), H(1.00 M) Pt
• (a) EFe EoFe(III)/Fe(II) (0.059/1) log
  Fe2/Fe3
• 0.771 0.059 log (0.150)/(0.0015 2)
  0.671 V
• EMn EoMnO4-/Mn2 (0.059/5)log
  Mn2/MnO4-H8
•   1.51 0.059/5 log
  (0.005)/(0.02)(1.00) 8 1.52 V
• (b) At eq., EFe EMn, ??????????,
• ??????????????,?
• EFe 0.771 0.059 log (0.05)/(0.103) 0.790
  V
• (c) Ecell EMn - EFe 1.52 0.671 0.849 V
• (d) At eq., EFe EMn, ??Ecell 0 V

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Concentration Cells

• Determine
• a) e- flow direction?
• b) anode? cathode?
• c) E ? at 25?

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Ex Systems involving ppt

• (ex) Calculate Ksp for AgCl at 25?
• 
  e0.58V
• soln

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14-6 Reference Electrodes
Indicator electrode responds to analyte
concentration Reference electrode maintains a
fixed potential

• ?

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Reference Electrodes

• Silver-Sliver Chloride
• AgCl e- ? Ag(s) Cl-
• E0 0.222 V
• E (saturated KCl) 0.197 V

• Calomel
• Hg2Cl2 2e- ? 2Hg(l) 2Cl-
• E0 0.268 V
• E (saturated KCl) 0.241 V
• saturated calomel electrode (S.C.E.)

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Voltage conversion between different reference
scales

• The potential of A ?

?