chapter 18.pptx

18 - 1
Chemistry: A Molecular Approach
Third Canadian Edition
Chapter 18
Electrochemistry
Copyright © 2020 Pearson Canada Inc.

Copyright © 2020 Pearson Canada Inc. 18 - 2
Chapter Outline
18.1 Pulling the Plug on the Power Grid
18.2 Voltaic (or Galvanic) Cells: Generating Electricity from
Spontaneous Chemical Reactions
18.3 Standard Electrode Potentials
18.4 Cell Potential, Gibbs Energy, and the Equilibrium
Constant
18.5 Cell Potential and Concentration
18.6 Batteries: Using Chemistry to Generate Electricity
18.7 Electrolysis: Driving Nonspontaneous Chemical
Reactions with Electricity
18.8 Corrosion: Undesirable Redox Reactions

18.1 Pulling the Plug on the Power Grid
2 2 2
2H ( ) O ( ) 2H O( )
g g g
 

BC Transit had 20 hydrogen fuel cell buses in service
in Whistler, BC, for the 2010 Winter Olympic games.
[Marcel Sloover]

18.2 Voltaic (or Galvanic) Cells: Generating
Electricity from Spontaneous Chemical Reactions
Figure 18.1 A Spontaneous Oxidation–Reduction Reaction
[© Richard Megna/Fundamental Photographs, NYC]

Voltaic Cell (1 of 4)
Figure 18.2 A Voltaic Cell

Rate of flow of electrons
–1
1 A = 1 C s
Driving force
–1
1 V = 1 J C
Cell Potential (Ecell) or cell electromotive force (emf)
The difference in potential of the two half cells

Figure 18.3 An Analogy for Electrical Current
[Alejandro Díaz Díez/AGE Fotostock]

Standard Cell Potential (E°cell) or standard emf
1 M concentration for reactants and 1 bar pressure for gases.
25°C unless otherwise noted
cell
2+ 2
1.1
Zn( )+Cu ( ) Zn ( ) Cu ) 0V
(
s aq aq E
s

  

 
cell
2+ 2
0.6
Ni( )+Cu ( ) Ni ( ) Cu ) 2V
(
s aq aq E
s

  

 
A measure of the overall tendency of the reaction to occur
spontaneously. A positive cell potential indicates that the
forward reaction is spontaneous.

Electrochemical Cell Notation (1 of 3)

– 2 2
4 2
( )
5 Fe(s) 2 MnO ( ) 16 H 5 ( )
Fe 2 Mn 8 H O
( ) ( )
aq aq aq aq l
  
  

 
2
(
Fe(s) Fe 2
)
aq e
 
 


Oxidation
+ 2
4 2
MnO ( ) 8H ( ) Mn ( )
5 8 H O( )
aq aq e aq l
  


  
Reduction

2 – 2+
4
| ( ) || ( )
Fe( ) Fe MnO , H ,
( ) ( ) |
M P )
t(
n
s aq aq aq aq s
 
Figure 18.4 Inert Platinum Electrode

18.3 Standard Electrode Potentials
Figure 18.5 An Analogy for Electrode Potential

Standard Hydrogen Electrode
Figure 18.6 The Standard Hydrogen Electrode
2
2 H ( ) 2 H ( )
e
aq g

 

cell
0.00 V
E 

Measuring Electrode Potential (1 of 3)
Figure 18.7 Measuring Electrode Potential

2
2
( ) | ( ) || (
Zn Zn H )
H
) | (
s aq aq g
 
cell cathode anode
–
E E E
   
2
Zn/Zn
0.76 V 0.00V – E 
 
2
Zn/Zn
– 0.76 V
E 
 
The Zn/Zn2+ electrode will be the anode when it is paired with the SHE.

Consider 2
2
H H C
( ) | ( u
) || ( ) u
| ( )
C
g aq aq s
 
cell cathode anode
–
E E E
   
2+
Cu/Cu
0.34 V – 0.00V
E
 
2
Cu/Cu
0.34 V
E 
  
The Cu/Cu2+ electrode will be the cathode when it is paired with the SHE or
an electrode with a more negative cell potential.

Electrochemical Cell with SHE as the Anode
Figure 18.8 Electrochemical Cell with SHE as the Anode

Summarizing Standard Electrode Potentials:
• The electrode potential of the standard hydrogen electrode (SHE) is
exactly zero.
• The electrode in any half-cell with a greater tendency to undergo
reduction is positively charged relative to the SHE and therefore has
a positive E°.
• The electrode in any half-cell with a lesser tendency to undergo
reduction (or greater tendency to undergo oxidation) is negatively
charged relative to the SHE and therefore has a negative E°.
• The cell potential of any electrochemical cell (E°cell) is the difference
between the electrode potentials of the cathode and the anode
(E°cell = E°cat− E°an).
• E°cell is positive for spontaneous reactions and negative for
nonspontaneous reactions.

Standard Electrode Potentials (1 of 3)

Predicting the Spontaneous Direction of an
Oxidation-Reduction Reaction (1 of 3)
cell cathode anode
–
–0.23 – ( –1.18 )
0.95
E E E
V V
V
   

 

Figure 18.9 Mn/Ni2+ Electrochemical Cell

Summarizing the Prediction of Spontaneous Direction
for Redox Reactions:
• The half-reaction with the more positive electrode potential
attracts electrons more strongly and will undergo reduction
(so substances listed at the top of Table 18.1 tend to
undergo reduction; they are good oxidizing agents).
• The half-reaction with the more negative electrode
potential repels electrons more strongly and will undergo
oxidation (so substances listed near the bottom of Table
18.1 tend to undergo oxidation; they are good reducing
agents).
• Any reduction reaction in Table 18.1 is spontaneous when
paired with the reverse of the reaction listed below it.

Predicting Whether a Metal Will Dissolve in Acid
When zinc is immersed in hydrochloric acid, the zinc is oxidized, forming
ions that become solvated in the solution. Hydrogen ions are reduced,
forming bubbles of hydrogen gas.
[© Richard Megna/Fundamental Photographs, NYC]
2 H ( ) 2
aq e
 
 2
2
H ( )
Zn(s) Zn ( ) 2
g
aq e
 



 
2
2
2H ( ) Zn(s) H ( ) Zn ( )
aq g aq
 
 
 

18.4 Cell Potential, Gibbs Energy, and the
Equilibrium Constant
Spontaneous
• ΔrG° < 0
• E°cell > 0
• K > 1
Non-spontaneous
• ΔrG° > 0
• E°cell < 0
• K < 1

The Relationship Between ΔG and E°
cell

The Relationship Between E°
cell and K (1 of 2)
r
r
– ln [18.5]
– cell
G RT K
G nFE
  
   
– – ln
ln [18.6]
cell
cell
nFE RT K
RT
E K
nF
 
 
0.0257 V
at 25 C ln [18.7]
cell
E K
n
  

The Relationship Between E°
cell and K (2 of 2)

18.6 Cell Potential and Cell Concentration

Cell Potential and Concentration
Figure 18.12 Cell Potential and Concentration

The Nernst Equation (1 of 2)
The Nernst Equation
r r
– ln [18.8]
G G RT Q
   
r
– cell
G nFE
  
– – – ln
cell cell
nFE nFE RT Q
 
– ln [18.9]
cell cell
RT
E E Q
nF
 
0.0257V
– ln [18.10]
cell cell
E E Q
n
 

The Nernst Equation (2 of 2)
0.0257 V
– log
cell cell
E E Q
n
 
• When a redox reaction within a voltaic cell occurs under
standard conditions, Q = 1; therefore E°cell = Ecell.
• When a redox reaction within a voltaic cell occurs under
conditions in which Q < 1 the greater concentration of reactants
relative to products drives the reaction to the right, resulting in
Ecell > E°cell.
• When a redox reaction within an electrochemical cell occurs
under conditions in which Q > 1 the greater concentration of
products relative to reactants drives the reaction to the left,
resulting in Ecell < E°cell.
• When a redox reaction reaches equilibrium, Q = K. The redox
reaction has no tendency to occur in either direction and Ecell = 0.

Concentration Cells (1 of 3)
Figure 18.13 Cu/Cu2+ Concentration Cells

2
Cu 2 Cu 0
( ) ( ) .34V
aq e s E
 
 
   
Cathode
2
( )
Cu 2 Cu 0.34V
( )
aq e s E
 
 
   
Anode
Standard Cell Potential
–
0.00V
cell cathode anode
E E E
  


18.7 Batteries: Using Chemistry to Generate
Electricity
Dry-Cell Batteries
Figure 18.16 Alkaline Batteries
[Dorling Kindersley, Ltd.]
–
2
Zn 2 O
( ) ( )
Zn(OH
H
) ( ) 2
s aq
s e




( )
Oxidat de
i o
n An
o
2 2
2MnO ( ) 2H O( )+ 2
s l e

( )
Reduct Catho
i de
on
2MnO(OH)( )+2 OH ( )
s aq



2 2
2
( ) (
Zn 2MnO H O
2MnO(OH)( ) Zn
) (
)
O
)
H (
( )
s s l
s s
 


 
Overall

Lead-Acid Storage Batteries (1 of 2)
Figure 18.17 Lead-Acid Storage Battery

Lead-Acid Storage Batteries (2 of 2)

Other Rechargeable Batteries (1 of 4)
The Nickel–Cadmium (NiCad) Battery
2
Cd( ) 2OH ( ) Cd(OH) ( ) 2
s aq s e
 
 
 
( )
A
Oxidatio no
n de
–
2 2
2 NiO(OH)( ) 2H O 2 2 Ni O
( ) ( ) ( ) 2OH ( )
H
s l e a
s q

   

( )
Reduc Catho
ti n d
o e
2 2 2
Cd( ) 2 NiO(OH) 2H O
( ) ( ) Cd(OH) ( ) 2 NiO(OH) ( )
s
s s s
l
  
 
Overall
produces about 1.3 V

The Nickel–Metal Hydride (NiMH) Battery
– –
2
( ) ( )
MH OH M( ) H O )
2 (
s aq s l e
 
  
( )
Oxidatio Anode
n
–
2 2
NiO(OH) 2H
( ) ( ) ( ) ( )
O Ni OH OH ( )
s l e aq
s 

  

( )
Reduc Cathode
tion
2
MH NiO(OH)
( ) ( ) M( ) NiO(OH) ( )
s s s s
 
 
Overall

Figure 18.18 Lithium-Ion Battery

Table 18.2 Energy Density and Overcharge Tolerance of
Several Rechargeable Batteries
Battery Type Energy Density (W h kg−1) Overcharge Tolerance
NiCad 45–80 Moderate
NiMH 60–120 Low
Li ion 110–160 Low
Pb storage 30–50 High

Fuel Cells (1 of 2)
Figure 18.19 Hydrogen-Oxygen Fuel Cell

Fuel Cells (2 of 2)

18.7 Electrolysis: Driving Nonspontaneous
Chemical Reactions With Electricity
Figure 18.20 Electrolysis of Water

Electrolysis (1 of 2)
Figure 18.21 Silver Plating

Electrolysis (2 of 2)
–0.76V 0.34V
cell cat
– 0.34V – (–0.76V) 1.10V
an
E E E
  
Figure 18.22 Voltaic Versus Electrolytic Cells

Summarizing, Characteristics of Electrochemical Cell
Types:
In all electrochemical cells:
• Oxidation occurs at the anode.
• Reduction occurs at the cathode.
In voltaic cells:
• The anode is the source of electrons and has a negative charge (anode −).
• The cathode draws electrons and has a positive charge (cathode +).
In electrolytic cells:
• Electrons are drawn away from the anode, which must be connected to
the positive terminal of the external power source (anode +).
• Electrons are forced to the cathode, which must be connected to the
negative terminal of the power source (cathode −).

Predicting the Products of Electrolysis (1 of 8)
anion
oxidized
at
anode
cation
reduced
at
cathode
Figure 18.23 Electrolysis of Molten NaCl

Mixtures of Cations
–
–
Na Na –2.71 V
( ) ( )
( )
K K –2.9
( ) 2 V
aq e s E
aq e s E


 
  
 
  
or Anions
– –
2
– –
2
Cl 2 2 Cl 1.36 V
B
( ) ( )
( ) ( )
r 2 2 Br º 1.09 V
g e aq E
g e aq E
 
   
 
  

Aqueous Solutions Water can be either oxidized or reduced
(water is oxidized)
Anode
–
2 2
O 4 H ( ) 2H O
( (
0.82V (pH 7)
) ) 1.23V
g aq 4e l E
E

  
   
  
(water is reduced)
Cathode
–
2 2
2H O H ( ) 2OH ( ) 0.83V
0.41V
(
( 7)
)
pH
2e g aq E
E
l 
 
   
 


– –0.41V – ( 0.82V) –1.23V
cell cat an
E E E
   

Pure water is a poor conductor of electrical current, but the addition of an electrolyte allows
electrolysis to take place, producing hydrogen and oxygen gas in a stoichiometric ratio.
[Charles D. Winters/Science Source]

Aqueous Solutions
At the cathode it is possible to reduce either Na+ or H2O
+ –
2 Na ( ) 2 2 Na( ) 2.71V
aq e s E
 
   
Reduction
–
2 2
2H O( ) 2 H ( ) 2OH ( )
–0.41V (pH7)
l e g aq
E

 
 

Reduction
At the anode it is possible to oxidize either H2O or I−
–
2
I ( ) 2 2I ( ) 0.54V
s e aq E

 
   
Oxidation
+ –
2 2
O ( ) 4H ( ) 4 2H O( )
0.82V(pH7)
g aq e l
E
  

 
Oxidation

Figure 18.24 Electrolysis of Aqueous NaI

Overpotential
+ –
2 Na ( ) 2 2 Na( ) 2.71V
aq e s E
 
   
Reduction
–
2 2
2H O( ) 2 H ( ) 2OH
–0.41V (pH7)
l e g
E

 
 

Reduction
–
2
Cl ( ) 2 2Cl ( ) 1.36V
s e aq E

 
  
Oxidation
+ –
2 2
O ( ) 4H ( ) 2H O( )
0.82V(pH7)
g aq 4e l
E
  

 
Oxidation

Figure 18.25 Electrolysis of Aqueous NaCl: The Effect of Overpotential

Stoichiometry of Electrolysis
Figure 18.26 Electrolytic Cell for Copper Plating

18.9 Corrosion: Undesirable Redox
Reactions
–
2 2
O ( ) 4H O 4OH
( ) ( ) 0.40V
g l 4e aq E

  
   
+ –
2 2
O ( ) 4H 2H
( ) O( ) 1.23V
g aq 4e l E
  
   
The cationic form of a metal must
usually be reduced to extract the metal
from its ore. In corrosion, the metal is
oxidized back to its more natural state.
[StockPhotosArt/Shutterstock]
Aluminum is stable because its oxide
forms a protective film over the underlying
metal, preventing further oxidation.
[Donovan Reese/The Image Bank/Getty Images]

Corrosion (1 of 2)
A scratch in paint often
allows the underlying
iron to rust.
[Stokkete//Fotolia]
2 –
Fe ( ) 2 F ( )
e –0.45V
g e s E

 
  
+ –
2 2
O ( ) 4H ( ) 4 2H O 1.23V
( )
g aq e l E
  
   
2 2
2 2
2 Fe O 4 H 2
( ) ( ) ( ) ( )
H O 2 Fe 1
( ) .68 V
aq g aq l aq E
  
    
  

2
2 2 2 3 2
( ) ( ) ( ) (
4 Fe O 4 2 H O 2 Fe O
) ( • H O 8
( H
) ) ( )
aq g n l s n s aq
 
 


  

Corrosion (2 of 2)
Figure 18.27 Corrosion of Iron: Rusting
[Jovan Nikolic/Fotolia]

Preventing Corrosion
In galvanized nails, a layer of zinc
prevents the underlying iron from rusting.
The zinc oxidizes in place of the iron,
forming a protective layer of zinc oxide.
[CDV, LLC/Pearson Science]
If a metal more active than iron, such as
magnesium or aluminum, is in electrical contact
with iron, the metal rather than the iron will be
oxidized. This principle underlies the use of
sacrificial electrodes to prevent the corrosion of
iron.

End of Chapter 18

chapter 18.pptx

Recommended

Recommended

More Related Content

Similar to chapter 18.pptx

Similar to chapter 18.pptx (20)

Recently uploaded

Recently uploaded (20)

chapter 18.pptx