Prentice Hall © 2003 Chapter 20

Chapter 20Chapter 20ElectrochemistryElectrochemistry

Prentice Hall © 2003 Chapter 20

• Oxidation-reduction reactions• Oxidation numbers• Oxidation of metals by acids and salts• The activity series

ALL these to be done in class

Prentice Hall © 2003 Chapter 20

• In the reaction

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g).

• Which element is oxidised and which one is reduced?• Oxidation – loss of e-

• Reduction – gain of e-

Oxidation-Reduction Oxidation-Reduction ReactionsReactions

Prentice Hall © 2003 Chapter 20

• Law of conservation of mass: the amount of each element present at the beginning of the reaction must be present at the end.

• Conservation of charge: electrons are not lost in a chemical reaction.

Balancing Oxidation-Balancing Oxidation-Reduction ReactionsReduction Reactions

Prentice Hall © 2003 Chapter 20

Half ReactionsHalf-reactions are a convenient way of separating oxidation and reduction reactions.

Prentice Hall © 2003 Chapter 20

• The half-reactions for

Sn2+(aq) + 2Fe3+(aq) Sn4+(aq) + 2Fe3+(aq)

are………

Prentice Hall © 2003 Chapter 20

Balancing Equations by the Method of Half Reactions

The two incomplete half reactions are

MnO4-(aq) Mn2+(aq)

C2O42-(aq) 2CO2(g)

Balance the overall reaction equation in an acidic solution

Prentice Hall © 2003 Chapter 20

Balancing Equations for Reactions Occurring in Basic Solution

• We use OH- and H2O rather than H+ and H2O.

• The same method as for acidic solution is used, but OH- is added to “neutralize” the H+ used.

Prentice Hall © 2003 Chapter 20

• If a strip of Zn is placed in a solution of CuSO4, Cu is deposited on the Zn and the Zn dissolves by forming Zn2+.

Voltaic CellsVoltaic Cells

Prentice Hall © 2003 Chapter 20

• Zn is spontaneously oxidized to Zn2+ by Cu2+.• The Cu2+ is spontaneously reduced to Cu0 by Zn.

Prentice Hall © 2003 Chapter 20

“Rules” of voltaic cells:1. At the anode electrons are products. (Oxidation)

2. At the cathode electrons are reagents. (Reduction)

3. Electrons can’t swim!

Prentice Hall © 2003 Chapter 20

• Anions and cations move through a porous barrier or salt bridge.

• Cations move into the cathodic compartment to neutralize the excess negatively charged ions

• Anions move into the anodic compartment to neutralize the excess Zn2+ ions formed by oxidation

Prentice Hall © 2003 Chapter 20

A Molecular View of Electrode Processes

Prentice Hall © 2003 Chapter 20

• e- flow from anode to cathode because the cathode has a lower electrical potential energy than the anode.

• 1 V is the potential difference required to impart 1 J of energy to a charge of one coulomb:

Cell EMFCell EMF

C 1J 1

V 1

Prentice Hall © 2003 Chapter 20

• 1 V is the potential difference required to impart 1 J of energy to a charge of one coulomb:

C 1J 1

V 1

Prentice Hall © 2003 Chapter 20

• Electromotive force (emf) is the force required to push electrons through the external circuit.

• Cell potential: Ecell is the emf of a cell.

• For 1M solutions at 25 C (standard conditions), the standard emf (standard cell potential) is called Ecell.

Prentice Hall © 2003 Chapter 20

Standard Reduction (Half-Cell) Potentials

• Standard reduction potentials, Ered are measured relative to the standard hydrogen electrode (SHE).

Prentice Hall © 2003 Chapter 20

Prentice Hall © 2003 Chapter 20

• For the SHE, we assign

2H+(aq, 1M) + 2e- H2(g, 1 atm)

• Ered = 0.

anodecathode redredcell EEE

Prentice Hall © 2003 Chapter 20

• For Zn:

Ecell = Ered(cathode) - Ered(anode)

0.76 V = 0 V - Ered(anode).

• Therefore, Ered(anode) = -0.76 V.

• Standard reduction potentials must be written as reduction reactions:

Zn2+(aq) + 2e- Zn(s), Ered = -0.76 V.

Prentice Hall © 2003 Chapter 20

• Changing the stoichiometric coefficient does not affect Ered.

• Therefore,

2Zn2+(aq) + 4e- 2Zn(s), Ered = -0.76 V.

Prentice Hall © 2003 Chapter 20

• Reactions with Ered < 0 are spontaneous oxidations relative to the SHE.

• The larger the difference between Ered values, the larger Ecell.

Prentice Hall © 2003 Chapter 20

Oxidizing and Reducing Agents

• The more positive Ered the stronger the oxidizing agent on the left.

• The more negative Ered the stronger the reducing agent on the right.

Prentice Hall © 2003 Chapter 20

Prentice Hall © 2003 Chapter 20

• More generally, for any electrochemical process

processoxidation processreduction redredcell EEE

Prentice Hall © 2003 Chapter 20

Example: For the following cell, what is the cellreaction and Eo

cell?

Al3+(aq) + 3e- → Al(s); EoAl = -1.66 V

Fe2+(aq) + 2e- → Fe(s); EoFe = -0.41 V

Al(s)|Al3+(aq)||Fe2+(aq)|Fe(s)

Prentice Hall © 2003 Chapter 20

Example: When an aqueous solution of CuSO4 is electrolysed, Cu metal is deposited:

Cu2+(aq) + 2e- → Cu(s)

If a constant current was passed for 5.00 h and 404 mg of Cu metal was deposited, what was the current?

Ans: 6.81 x 10-2 A