22.1 Electrochemical Processes

(Write ionic equation for: Zn + CuSO₄ ----> ZnSO₄ + Cu.
Write 1/2 reactions)
1. Zn metal dipped into CuSO₄ solution will become copper plated.

2. Spontaneous redox reaction:

3. Zn + Cu²⁺---->Zn²⁺ + Cu

4. Electrons are transferred from the Zn to the Cu.

5. See Table 22.1. Elements higher on table are more easily oxidized than those lower, thus Zn is more easily oxidized than Cu. Plating of Cu on Zn is spontaneous. Reverse would not occur and is nonspontaneous.

6. The conversion of chemical energy into electrical energy and visa versa are electrochemical processes.

7. A redox reaction is a source of electrical energy. Electricity is the flow of electrons in a wire. 2 half reactions must be physically separated, and electrons must pass through an external circuit.

8. An electrochemical cell is any device that converts chemical energy into electrical energy or electrical energy into chemical energy.

22.2 Voltaic Cells

1. Alessandro Volta.

2. Voltaic cells are electrochemical cells that are used to convert chemical energy into electrical energy. Spontaneous redox reactions.

3. A half-cell is one part of a voltaic cell in which either oxidation or reduction occurs.
4. Typical 1/2 cell has a metal like Zn in a solution of it's ions like zinc sulfate.
5. The other 1/2 cell has a metal like Cu in a solution of it's ions like copper sulfate.
6. A salt bridge allows only the sulfate anion to pass. The solutions don't mix completely.
7. An external wire carries the electrons from Zn to Cu.

22.3 Dry Cells

1. A dry cell is a voltaic cell in which the electrolyte is a paste.

2. Zn container is anode. Zn is oxidized. Graphite rod as cathode acts only as a conductor.

3. Thick paste of Manganese(IV) oxide (MnO₂), Zinc chloride(ZnCl₂), ammonium chloride(NH₄Cl) and water. 4. Manganese is reduced (MnO₂₎.
 

5. Voltage is 1.5 V

22.4 Lead Storage Batteries

1. A battery is a group of cells that are connected together.

2. In a car battery each cell produces 2 V.

3. Anode is spongy lead grid. Cathode is lead(IV)oxide grid (PbO₂). Electrolyte is sulfuric acid.

4. Half reactions:

Pb + SO₄^2- ----->PbSO₄ +2e^-(oxi)

PbO₂ + 4H⁺ + SO₄^2- + 2e^- ------> PbSO₄ +2H₂O (red)

5. The overall spontaneous redox reaction is the sum of the oxidation and reduction reactions.

Pb + PbO₂ + 2H₂O₄ --------> 2PbSO₄ + 2H₂O

6. During discharge lead sulfate builds up on plates and sulfuric acid concentration decreases.

7. During charging a current is passed through cell in reverse direction causing the reaction to go in the nonspontaneous direction. This reverse reaction is nonspontaneous and requires energy.

22.5 Fuel Cells

1. Fuel cells are voltaic cells in which a fuel substance undergoes oxidation and from which electrical energy is obtained continuously.

2. Do not have to be recharged, but require a constant source of fuel.

5. Half reactions:

Oxidation: 2H₂ + 4OH- ---> 4H₂O +4e^-
Reduction: O₂ + 2H₂O +4e^- ----->4OH^-
Overall reaction: 2H₂ + O₂ ------>2H₂O
6. Other fuels - methane, ammonia. Oxidizers - chlorine, ozone.
22.6 Half-Cells - Honors
1.The electrical potential of a voltaic cell is the ability of the cell to produce an electric current.
2. Usually measured in volts.
3. Can't measure half-cells separately, but can measure difference in pontential when 2 half-cells are connected. Eg., 1.0 M Zn-Cu cell is +1.10 V.
4. The reduction potential of a half-cell is a measure of the tendency of a given half-reaction to occur as a reduction. The difference between the reduction potentials of the two half-cells is acalled the cell potential.
4.
5. The standard potential (E^o_cell) is the measured cell potential when the ion concentrations in the half-cells are 1.00 M, gases are at a pressure of 101.3 kPa, and the temperature is 25^oC.
6. Half-cell potentials cannot be measured directly, so scientists use hydrogen as a reference.
7. The standard hydrogen electrode is used with other electrodes so that the reduction potentials of those cells can be measured.
8. 
9. Whether this half-cell reaction occurs as a reduction or as an oxidation is determined by the reduction potential of the half-cell to which it is connected.
22.7 Standard Reduction Potentials
1. By connecting a Zn half-cell to a hydrogen half-cell we can calculate the cell potential (E^o_cell) of Zn with a voltmeter = +.76 V.
2. E^o_cell = E^o_red - E^o_oxid
3. +.76V = 0.00V - E^o_Zn²⁺
4. E^o_Zn²⁺ = -.76 V (standard reduction potential for Zn)
5. Reduction potential for Zn is negative because it's tendancy is to be oxidized.
6. Another example is Cu half cell. If compared to hydrogen half-cell it it's reduction potential is +.34 V. Value is positive because tendancy is to be reduced.
7. See Table 22.2, page 636.
22.8 Calculating Cell Potentials
1. Using reduction potentials, we can predict in which of half-cells reduction and oxidation will occur.
2. See examples 1, 2; problems 17., 18.
22.9 Electrolytic Cells

1. The process in which electrical energy is used to bring about a chemical change is called electrolysis.

2. Electrolytic cells are electrochemical cells used to cause a chemical change through the application of electrical energy. Energy is used to make nonspontaneous reactions go to products.

3. Unlike voltaic cells, cathode is negative electrode, anode is positive. This is because the cathode is connected to anode (-) of battery and anode is connected to cathode (+) of battery.
4. Anode is site of oxidation. Cathode is site of reduction

5. See Figure 22.12, page 642
6. Electrolysis of Water:
Water is reduced to hydrogen at cathode:

Water is oxidized at anode.