Transcript Chapter 17 Electrochemistry

John E. McMurry • Robert C. Fay
C H E M I S T R Y
Sixth Edition
Chapter 17
Electrochemistry
© 2012 Pearson Education, 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
Galvanic Cells
Zn(s) + Cu2+(aq)
Oxidation half-reaction:
Reduction half-reaction:
Zn2+(aq) + Cu(s)
Zn(s)
Zn2+(aq) + 2e
Cu2+(aq) + 2e
Cu(s)
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.
Galvanic Cells
•
Cathode:
• The electrode where
reduction occurs.
• The electrode where
electrons are
consumed.
• Is what cations
migrate toward.
• Has a positive sign.
Galvanic Cells
•
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
17/8
Galvanic Cells
Why do negative ions (anions) move toward the negative electrode (anode)?
Galvanic Cells
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
Zn(s)
Cu2+(aq) + 2e
Zn(s) + Cu2+(aq)
No electrons should be appeared in the overall cell reaction
Zn2+(aq) + 2e
Cu(s)
Zn2+(aq) + Cu(s)
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
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)
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.
Cell Potentials and Free-Energy
Changes for Cell Reactions
1J=1Cx1V
joule (J)
SI unit of energy
volt (V)
SI unit of electric potential
coulomb (C)
Electric charge
1 coulomb is the amount of charge transferred when a current of 1 ampere
(A) flows for 1 second.
Cell Potentials and Free-Energy
Changes for Cell Reactions
faraday or Faraday constant
The electric charge on 1 mol of electrons and is equal to 96,500 C/mol e
DG = nFE
Free-energy change
or
DG° = nFE°
Cell potential
Number of moles of electrons transferred in the
reaction
Cell Potentials and Free-Energy
Changes for Cell Reactions
The standard cell potential at 25 °C is 1.10 V for the reaction:
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Calculate the standard free-energy change for this reaction at 25 °C. Is the
reaction spontanous at this condition?
Examples
 Calculate the cell potential at standard state (Eocell) for
the following reaction. Then write the half reactions
I2(s) + 2 Br-(aq)  2I-(aq) + Br2(l)
Go = 1.1 x
105J
Standard Reduction Potentials
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
H2(g)
Cu2+(aq) + 2e
H2(g) + Cu2+(aq)
2H+(aq) + 2e
Cu(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
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
The standard hydrogen electrode (S.H.E.) has been chosen to be the
reference electrode.
Standard Reduction Potentials
The standard hydrogen electrode (S.H.E.) has been chosen to be the
reference electrode.
2H+(aq, 1 M) + 2e
H2(g, 1 atm)
H2(g, 1 atm)
2H+(aq, 1 M) + 2e
E°ox = 0 V
E°red = 0 V
Standard Reduction Potentials
Anode half-reaction:
H2(g)
2H+(aq) + 2e
Cu2+(aq) + 2e
Cathode half-reaction:
H2(g) + Cu2+(aq)
Overall cell reaction:
Cu(s)
2H+(aq) + Cu(s)
E°cell = E°ox + E°red
0.34 V = 0 V + E°red
A standard reduction potential can be defined:
Cu2+(aq) + 2e
Cu(s)
E° = 0.34 V
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)
b. Fe2+(aq) + 2e-  Fe(s)
Al3+(aq) + 3e-  Al(s)
Eored = -0.40 V
Eored = 0.80 V
Eored = -0.44 V
Eored = -1.66 V
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
n
log K
in volts, at 25°C
The 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 oC that has the following ion concentrations?
[Fe3+] = 1.0 × 104 M
[Cu2+] = 0.25 M
[Fe2+] = 0.20 M
Example
 Use the tabulated half-cell potentials to calculate K for
the following oxidation of copper by H+
2Cu(s) + 2H+(aq)  Cu2+(aq) + H2(g)