Electrochemistry

1
Chapter 13 14

• Electrochemistry
• and
• Electrode Potentials

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• Oxidation a loss of electrons to an
  oxidizing agent
• Reduction a gain of electrons from a
  reducing agent

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16A Principles

• Reduction-oxidation reaction
  (redox reaction)
• Ox1 Red2 ? Red1 Ox2
• An oxidizing substance
• Ma ne- ? M(a-n)
• An reducing substance
• Ma ? M(a-n) ne-

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16B Electrochemical Cells

• (1) Galvanic (Voltaic) cell
• a chemical reaction spontaneously occurs to
  produce electrical energy.
• Ex lead storage battery
• (2) Electrolytic cell
• electrical energy is used to force a
  nonspontaneous chemical reaction to occur.
• Ex electrolysis of water

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• the anode oxidation occurs
• the cathode reduction occurs
• Salt bridge allows charge transfer through the
  solutions but prevents mixing of the solutions.
• Fig 16-2

Cu reducing agent Ag oxidizing agent
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• Electrode potential the tendency of the ions to
  give off or take on electrons.
• Normal Hydrogen Electrode (NHE)
• or Standard Hydrogen Electrode (SHE)
• 2H 2e- H2 or H e- 1/2H2
• Eo (????) 0.000 V
• Table 16.1
• (?1953, the 17th IUPAC meeting??????????????? )

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• Potential are dependent on con. temp.
• Standard reduction potential activity1
• The more positive the electrode potential, the
  greater the tendency of the oxidized form to be
  reduced.
• The more negative the electrode, the greater the
  tendency of the reduced form to be oxidized.

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Related to free energy
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The Nernst Equation
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Ex

• Predict whether 1M HNO3 will dissolve gold metal
  to form 1M Au3?

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Cell representation

• anode?solution?cathode

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The Nernst Equation C effect

• Activities should be used in the Nernst equation.
  
• We will use concentrations here because
  titrations deal with large potential changes, and
  the errors are small by doing so.
• Table 16-1, ???????(activity)??1,
• ?????????

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Dependence of the cell potential on C

• 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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ex

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

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After?????reached eq., the cell voltage
necessarily becomes zero and the reaction is
complete.
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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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• 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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16C-7 Limitation to use E0

• The sources of differences
• For Fe3 e- ? Fe2
• (1) Use conc vs ax (activities)
• gFe(II)/ gFe(III) 0.4/0.18 at m 0.1M
• (2) Other equilibria
• complexes Fe(III) with Cl-, SO4-2 are more
  stable than those of Fe(II).

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16C-7 Formal Potential

• Ex Ce4 e- ? Ce3 E1.6V
• with HA- E?1.61V
• ?????????????????,???????? 
• ????????Eo??????Eo
• Formal potential (E)
• The standard potential of a redox couple with the
  oxidized and reduced forms at 1M concentrations
  and with the solution conditions specified.
• Ex Ce4/Ce3 in 1M HCl E1.28V

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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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Dependence of potential on complexation
Complexing one ion reduces its effective
concentration, which changes the potential.
In effect, weve stabilized the Fe3 by
complexing it, make it more difficult to reduce.
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Ex Systems involving ppt

• Calculate Ksp for AgCl at 25? E 0.58V

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Sol
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Ex 17-4 Calculate the cell potential for
AgAgCl(satd), HCl(0.0200 M)H2(0.800 atm), Pt

• Sol 2H 2e- ? H2(g) E0H/H2 0.000 V
• AgCl(s) e- ? Ag(s) E0AgCl/Ag 0.222 V

-0.0977 V
Eleft 0.222 0.0592 logCl- 0.222 0.0592
log 0.0200 0.3226 V
Ecell Eright Eleft -0.0977 0.3226
-0.420 V 2H 2Ag(s) ? H2(g) 2AgCl(s)
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17B Calculating Redox Equilibrium Constants

• Ex 17-6 Calculate the equilibrium constant for
  the reaction
• 2Fe3 3I- ? 2Fe2
  I3-

Sol 2Fe3 2e- ? 2Fe2 E0 0.771 V
I3- 2e- ? 3I- E0
0.536 V
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Example 17-7 Calculate the equilibrium constant
for the reaction 2MnO4- 3Mn2
2H2O ? 5MnO2(s) 4H Sol 2MnO4- 8H 6e- ?
2MnO2(s) 4H2O E0 1.695 V
3MnO2(s) 12H 6e- ? 3Mn2 6H2O E0
1.23 V EMnO4-/MnO2 EMnO2/Mn2
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17C Constructing Redox Titration Curves
Example 17-8 Obtain an expression for the
equivalence-point potential in the titration of
0.0500 M U4 with 0.1000 M Ce4. Assume that both
solutions are 1.0 M in H2SO4. U4
2Ce4 2H2O ? UO22 2Ce3 4H Sol UO22
4H 2e- ? U4 2H2O E0 0.334 V
Ce4 e- ? Ce3
E0' 1.44 V
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Table 17-1
Electrode Potential versus SHE in Titrations with 0.100 M Ce4 Electrode Potential versus SHE in Titrations with 0.100 M Ce4 Electrode Potential versus SHE in Titrations with 0.100 M Ce4 Electrode Potential versus SHE in Titrations with 0.100 M Ce4
Potential, V vs. SHE Potential, V vs. SHE Potential, V vs. SHE
Reagent Volume, mL 50.00 mL of 0.0500 M Fe2 50.00 mL of 0.02500 M U4
5.00 0.64 0.316
15.00 0.69 0.339
20.00 0.72 0.352
24.00 0.76 0.375
24.90 0.82 0.405
25.00 1.06 ?Equivalence ? Point 0.703
25.10 1.30 1.30
26.00 1.36 1.36
30.00 1.40 1.40
Note H2SO4 concentration is such that H 1.0
M throughout.
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17D Oxidation/Reduction Indicators

• Self-indication
• If the titrant is highly colored, this color may
  be used to detect the end point.
• Ex MnO4- ? Mn2
• purple faint pink

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• Starch indicator
• This indicator is used for titrations involving
  iodine
• Starch I2 ? dark-blue color complex

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• Redox Indicators
• These are highly colored dyes that are weak
  reducing or oxidizing agents that can be oxidized
  or reduced
• Oxind ne- ? Redind

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• A potential equal to 2(0.059/n)V is required for
  a sharp color change
• n 1 ? 0.12V
• n 2 ? 0.060V
• The redox indicator reaction must be rapid and
  reversible.
• Table 17.2
• Ex(1) Ferroin tris(1,10-phenanthroline)io
  n(II) sulfate
• for titrations with cerium(IV)
• (2) Starch/Iodine soln.

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18B Reducing Agents

• Thiosulfate stable to air oxidation
• Iron(II) E0 0.771V
• for titration of cerium(IV), chromium(VI),
  vanadium(V)
• indicatorferroin or diphenylamine sulfonate.

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18C Oxidizing Agents

• Potassium permanganate (KMnO4)
• E01.51
• In neutral solution MnO4-?MnO2
• In acid solution MnO4-?Mn2
• Autocatalytic decomposition
• Standardization Na2C2O4
• 5H2C2O42MnO4-6H
  ?10CO22Mn28H2O

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• Cerium (IV) Ce4 / H2SO4 E0 1.44V
• Ce4 / HClO4 E0 1.70V
• the rate of oxidation of chloride ion is slow
• is stable in H2SO4
• (NH4)2Ce(NO3)6 can be obtained as a primary
  standard.
• indicator Ferroin

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• Potassium dichromate K2Cr2O7
• a slightly weaker oxidizing agent than KMnO4
  primary standard
• Cr2O72- ? Cr3 E0 1.331.00V in 1M
  HCl