Transcript Slide 1

John E. McMurry • Robert C. Fay
C H E M I
S T R Y
Fifth Edition
Chapter 17
Electrochemistry
Lecture Notes
Alan D. Earhart
Southeast Community College • Lincoln, NE
Copyright © 2008 Pearson Prentice Hall, Inc.
Galvanic Cells
Electrochemistry: The area of chemistry concerned with the
interconversion of chemical and electrical energy.
Galvanic (Voltaic) Cell: A spontaneous chemical reaction which generates
an electric current.
Electrolytic Cell: An electric current which drives a nonspontaneous
reaction.
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/2
Galvanic Cells
Zn(s) + Cu2+(aq)
Oxidation half-reaction:
Reduction half-reaction:
Zn2+(aq) + Cu(s)
Zn(s)
Zn2+(aq) + 2e-
Cu2+(aq) + 2e-
Copyright © 2008 Pearson Prentice Hall, Inc.
Cu(s)
Chapter
17/3
Galvanic Cells
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Galvanic Cells
•
Anode:
• The electrode where oxidation occurs.
• The electrode where electrons are produced.
• Is what anions migrate toward.
• Has a negative sign.
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/5
Galvanic Cells
•
•
Cathode:
• The electrode where reduction occurs.
• The electrode where electrons are consumed.
• Is what cations migrate toward.
• Has a positive sign.
Salt Bridge: a U-shaped tube that contains a gel permeated with a
solution of an inert electrolytes
• Maintains electrical neutrality by a flow of ions
• Anions flow through the salt bridge from the cathode to anode
compartment
• Cations migrate through salt bridge from the anode to cathode
compartment
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/6
Galvanic Cells
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
Zn(s)
Zn2+(aq) + 2e-
Cu2+(aq) + 2eZn(s) + Cu2+(aq)
Cu(s)
Zn2+(aq) + Cu(s)
No electrons should be appeared in the overall cell reaction
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/7
17.2 Shorthand Notation for
Galvanic Cells
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/8
17.2 Shorthand Notation for
Galvanic Cells
 Cell involving gas
 Additional vertical line due to presence of addition phase
 List the gas immediately adjacent to the appropriate electrode
 Detailed notation includes ion concentrations and gas pressure
E.g
Cu(s) + Cl2(g)  Cu2+(aq) + 2 Cl-(aq)
Cu(s)|Cu2+(aq)||Cl2(g)|Cl-(aq)|C(s)
Example
 Consider the reactions below
 Write the two half reaction
 Identify the oxidation and reduction half
 Identify the anode and cathode
 Give short hand notation for a galvanic cell that employs the
overall reaction
Pb2+(aq) + Ni(s)  Pb(s) + Ni2+(aq)
Example
 Given the following shorthand notation, sketch out the
galvanic cell
Pt(s)|Sn2+,Sn4+(aq)||Ag+(aq)|Ag(s)
17.3 Cell Potentials and Free-Energy
Changes for Cell Reactions
Electromotive Force (emf): The force or electrical potential that pushes the
negatively charged electrons away from the anode (- electrode) and pulls them
toward the cathode (+ electrode).
It is also called the cell potential (E) or the cell voltage.
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/12
Cell Potentials and Free-Energy
Changes for Cell Reactions
1J=1Cx1V
joule
SI unit of energy
volt
SI unit of electric potential
coulomb
Electric charge
1 coulomb is the amount of charge transferred when a current of 1
ampere flows for 1 second.
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/13
Cell Potentials and Free-Energy
Changes for Cell Reactions
faraday or Faraday constant
the electric charge on 1 mol of electrons
96,5000 C/mol e-
DG = -nFE
free-energy change
or
DG° = -nFE°
cell potential
number of moles of electrons transferred in
the reaction
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/14
Cell Potentials and Free-Energy
Changes for Cell Reactions
The standard cell potential at 25 °C is 0.10 V for the reaction:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Calculate the standard free-energy change for this reaction at 25 °C.
DG° = -nFE°
= -(2 mol e-)
96,500 C
(1.10 V)
mol e-
1J
1CV
1 kJ
1000 J
DG° = -212 kJ
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/15
17.4 -17.5 Standard Reduction
Potentials
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
H2(g)
Cu2+(aq) + 2eH2(g) + Cu2+(aq)
2H1+(aq) + 2eCu(s)
2H1+(aq) + Cu(s)
The standard potential of a cell is the sum of the standard half-cell potentials for
oxidation at the anode and reduction at the cathode:
E°cell = E°ox + E°red
The measured potential for this cell: E°cell = 0.34 V
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/16
Standard Reduction
Potentials
 Eocell is the standard cell potential when both products and
reactants are at their standard states:
 Solutes at 1.0 M
 Gases at 1.0 atm
 Solids and liquids in pure form
 Temp = 25.0oC
Standard Reduction Potentials
 Spotaniety of the reaction can be determined by the positive
Eocell value
 The cell reaction is spontaneous when the half reaction with
the more positive Eo value is cathode
 Note: Eocell is an intensive property; the value is independent
of how much substance is used in the reaction
Ag+(aq) + e-  Ag(s)
Eored = 0.80 V
2 Ag+(aq) + 2e-  2 Ag(s)
Eored = 0.80V
Standard Reduction Potentials
H2(g) + Cu2+(aq)
2H1+(aq) + Cu(s)
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been chosen to be the
reference electrode.
2H1+(aq, 1 M) + 2eH2(g, 1 atm)
H2(g, 1 atm)
2H1+(aq, 1 M) + 2e-
Copyright © 2008 Pearson Prentice Hall, Inc.
E°ox = 0 V
E°red = 0 V
Chapter
17/20
Standard Reduction Potentials
Examples
 Of the two standard reduction half reactions below, write
the net equation and determine which would be the anode
and which would be the cathode of a galvanic cell.
Calculate Eocell
a.
Cd2+(aq) + 2e-  Cd(s)
Ag+(aq) + e-  Ag(s)
Eored = -0.40 V
Eored = 0.80 V
b.
Fe2+(aq) + 2e-  Fe(s)
Al3+(aq) + 3e-  Al(s)
Eored = -0.44 V
Eored = -1.66 V
17.6 The Nernst Equation
DG = DG° + RT ln Q
DG = -nFE and DG° = -nFE°
Using:
Nernst Equation:
RT
E = E° -
ln Q
nF
or
E = E° -
2.303RT
nF
log Q
or
E = E° -
0.0592 V
log Q
n
in volts, at 25°C
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/23
17.6The Nernst Equation
Consider a galvanic cell that uses the reaction:
Cu(s) + 2Fe3+(aq)
Cu2+(aq) + 2Fe2+(aq)
What is the potential of a cell at 25 °C that has the following ion concentrations?
[Fe3+] = 1.0 x 10-4 M
[Cu2+] = 0.25 M
Copyright © 2008 Pearson Prentice Hall, Inc.
[Fe2+] = 0.20 M
Chapter
17/24
Example
 Calculate the concentration of cadmium ion in the galvanic
cell below
Cd(s)|Cd2+(aq)(?M)||Ni2+(aq)(0.100M)|Ni(s)
Standard Cell Potentials and
Equilibrium Constants
Using
DG° = -nFE°
DG° = -RT ln K
and
-nFE° = -RT ln K
E° =
RT
ln K
=
nF
E° =
2.303 RT
log K
nF
0.0592 V
log K
in volts, at 25°C
n
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/26
Standard Cell Potentials and
Equilibrium Constants
Examples
 Calculate the equilibrium constant, Keq, for the reaction
below
Zn2+(aq) + 2e-  Zn(s)
Sn2+(aq) + 2e-  Sn(s)
Eored = -0.76 V
Eored = -0.14 V
Standard Cell Potentials and
Equilibrium Constants
Three methods to determine equilibrium constants:
1.
K from concentration data:
K=
[C]c[D]d
[A]a[B]b
2.
K from thermochemical data:
ln K =
-DG°
RT
3.
K from electrochemical data:
E° =
RT
ln K
nF
or
ln K =
nFE°
RT
Copyright © 2008 Pearson Prentice Hall, Inc.
Chapter
17/29