Chapter 20 Electrochemistry
Oxidation-Reduction Reactions
Loss of an electron - oxidation (LEO) charge increases
Gain of an electron - reduction (GER)
   2H
    Zn
2     H
2
  charge decreases
Oxidation - the value of the charge increases
  
Zn
2
  
Reduction - the charge decreases (it is reduced in size)
2 H
   
H
2
Balancing Equations
Must balance the mass and the charge
Use the method of half reactions .
Treat oxidation as a separate reaction.
Balance its mass and charge then treat reduction as a separate reaction, balancing its mass and charge, then add the reactions together.
1
Electrochemistry 1

Problem solving Tips
Step 1.
Verify that the reaction is an oxidation-reduction process
Step 2..
Separate the overall reaction into half-reactions
Step 3. Balance each half-reaction for mass. a.
First balance the elements other than H and O b.
Next balance the O atoms by adding H
2
O. c.
Then balance the H atoms by adding H+ when in acidic solution (or by adding OH- when in basic solution)
Step 4.
Balance the half reactions for charge.
Step 5.
Multiply the balanced half-reactions so the number of electrons lost in one half-reaction equals the number of electrons gained in the other.
Step 6.
Add the balanced half-reactions and simplifying by canceling species appearing on both sides of the equation.
Step 7.
Check the final result for mass and charge balance.
The half-reactions can be balanced initially as if they occurred in acidic solution, followed by "neutralization" by adding an equal number of OH- to each side.
2
Electrochemistry 2

Chemical Change Leading to an Electrical Current
Voltaic cell : a spontaneous redox reaction to do electrical work.
Direction of electron flow electrode (-) electrode (+) e
reduced species oxidized species cations salt bridge oxidized species e
anions reduced species
electrode - a device for conducting electrons into and out of solutions in electrochemical cells.
Electric neutrality is maintained by migration of ions through a porous glass disc or through a salt bridge. salt bridge allows flow of non-reacting charged species between compartments.
3
Electrochemistry 3

Electrochemical Cells and Potentials electromotive force ( emf ) : electrons move from an electrode of higher potential to one of lower potential (potential difference, in Volts).
Electrical work = charge x potential difference charge in coulombs, C charge of one electron : 1.6022 x 1019 C
Volt = Ошибка!
or Work (Joule) = 1 volt x 1 coulomb standard conditions ,
• pure solid, or
• 1.0 M solution, or
• gas at 1 bar., give standard electrode potential, E°
Cell potentials for product-favored electrochemical reactions have positive values.
4
Electrochemistry 4

E cell

E red

E ox standard half-cell, standard hydrogen electrode (SHE)
2 H
3
O
  aq, 1 M
 
2e


H
2
g, 1 bar
 
H
2
Tables list standard reduction potentials
 
E

0.00 V
Look at Zn, Cu cell
  
Cu
2
  aq, 1M
 
Zn
2
   
Cu s E cell
 
1.10
V
Zn is oxidized, use opposite of reduction potential in table
Zn
2
   
2e

  
E cell
  
Zn
2
   
2e

E cell
 
0.76 V
 
0.76 V
Cu is reduced, use reduction potential from table
Cu
2
   
2e

  
E cell
 
0.34 V
E cell

0.76 V
  
0.34 V
  
1.10 V
5
Electrochemistry 5

Oxidizing and Reducing Agents
The more positive the E red
value for a half-reaction, the greater the tendency for the reactant of the half-reaction to be reduced, and therefore, to oxidize another species.
Look at the half-reactions in the table of reduction potentials.
Most positive values of E red strongest oxidizing agent
F
2
(g) + 2e
–
2F
–
(aq)
•
•
•
•
• •
2H
+
(aq)
•
•
•
Li
+
(aq) +e
–
H
2
(g)
•
•
•
Li(s)
Most negative values of E red strongest reducing agent
F
2
is most easily reduced and consequently the strongest oxidizing reagent.
F
2
  
2e


2F
  
E red

2.87 V
6
Electrochemistry 6

Lithium ion, Li
+
, is the most difficult to reduce and is therefore the poorest oxidizing reagent.
  
e

  
E red
 
3.05 V
However, the reverse reaction is favorable, and makes Li(s) the strongest reducing agent on the table.
F
–
(aq) is the weakest reducing agent on the table.
7
Electrochemistry 7


G rxn
  nFE

G rxn
 
RT ln K
Faraday's constant, F, 96,485 C/mol (J/V mol) , charge on one mole of electrons
F


6.022

10
23 mol

1
 
1.6022

10

19
C


96485 C mol

1 n , number of electrons transferred in the reaction
Effect of Concentration on Cell EMF : The Nernst Equation

G
 
G

RTlnQ

 

 



E

E

RT ln Q nF

G
  nFE
 
RTlnK
E

RT lnK nF
E

RT

nF
K
Q


8
Electrochemistry 8

E

E

0.0257
V ln Q n
Q will be concentration of ions on right side over concentration of ions on left side. Must be given.
E

RT ln K nF
K
 nE
V
RT

V
F
K
 e nFE RT  e nE 0.0257_
V
9
Electrochemistry 9

Batteries and Fuel Cells
Primary Batteries (not easily reversed; at equilibrium "dead")
Secondary Batteries (easily reversed or "recharged") car battery
Fuel Cells (reactants supplied from external source)
Corrosion
Oxidation of most metals by oxygen is favorable.
Iron:
Cathode:
2
+

2
E red
 
V
Anode:

2+
-
E red
 
V anodic inhibition - coat with inert material, paint or layer of oxide of metal protecting (Al
2
O
3
) cathodic inhibition - make Fe the cathode, use Zn or Mg as the anode
(sacrificial metal)
10
Electrochemistry 10

Electrolysis - use of an electrical current to bring about chemical change.
Direction of electron flow
electrode (+) electrode (–) e
reduced species salt bridge oxidized species e
oxidized species cations anions reduced sp ecies
11
Electrochemistry 11

Current, I (amperes, A) = Ошибка!
• Current times time, in seconds, gives charge in coulombs, Q

It
• Divide charge by Faraday's constant to get moles of e-
• Divide moles of electrons by moles of electrons per mole of metal reduced
• Divide moles of metal by the atomic weight of the metal to get grams of metal current
  s time ( s )

1_ Faraday
96485_ C

1_ mol _ e

Faraday
 mol _ metal mol _ e

 g _ metal mol _ metal
 g _ metal m

QM nF
Ошибка!
Calculate the mass of aluminum produced in 1.00 hr by the electrolysis of molten
AlCl 3 if the electrical current is 10.0 A.
12
Electrochemistry 12