Ch20

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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

ANODE compartment

OXIDATION occurs

CATHODE comparment

REDUCTION occurs

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

Calculating the Potential E° of an Electrochemical Cell

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

Free Energy Change

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

G

 

nFE

nFE

 

nFE

RTlnQ

E

E

RT lnQ nF

E

E

RT ln Q nF

G

  nFE

 

RTlnK

E

RT lnK nF

E

RT



ln

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° and the Equilibrium Constant

E

RT ln K nF

ln

K

 nE

0.0257

V

RT

0.0257_

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:

O

2

(g) + 4 H

+

(aq) + 4e

2 H

2

O(l)

E red

 

0.14_

V

Anode:

Fe

Fe

2+

(aq) + 2e

-

E red

 

0. 44_

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

+

Battery

electrode (+) electrode (–) e

reduced species salt bridge oxidized species e

oxidized species cations anions reduced sp ecies

11

ANODE compartment

OXIDATION occurs

CATHODE comparment

REDUCTION occurs

Electrochemistry 11

Counting Electrons

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

m =

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

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