1. Chapter 21: Electrochemistry

2. Electrochemical Cells
The branch of chemistry that deals with electricity-related applications of oxidation-reduction reactions is called electrochemistry
Redox reactions involve the loss of electrons, which is accompanied by a transfer of energy as heat. If the oxidation and reduction reactions are separated from each other by a salt bridge, or porous barrier, the transfer of energy can be of electrical energy instead of heat

3. Electrochemical Cells
An electrode is a conductor used to establish electrical contact with a nonmetallic part of a circuit

4. Electrochemical Cells
A single electrode immersed in a solution of its ions is called a half-cell
1. The electrode where oxidation occurs is called the anode (negative electrode)
2. The electrode where reduction occurs is called the cathode (positive electrode)

5. Electrochemical Cells
The electrochemical cell is represented by the following notation:
anode electrode ǀ anode solution ǁ cathode solution ǀ cathode electrode
The double line represents the salt bridge between the two half-cells
F. The cell notation of Zn + Cu+2 → Zn+2 + Cu is:
Zn(s) ǀ Zn+2(aq) ǁ Cu+2(aq) ǀ Cu(s)

6. Voltaic Cells
Voltaic cells convert chemical energy from spontaneous oxidation-reduction reactions into electrical energy; another name for these cells is called galvanic cells
Dry voltaic cells include: (illustrations on pages 732)

7. Voltaic Cells
Zinc-carbon dry cells: batteries used in flashlights; zinc serves as the anode and carbon serves as the cathode
ox: Zn(S) → Zn2+(aq) + 2e –
red: 2MnO2(s) + H2O(l) + 2e– → Mn2O3(s) + 2OH –(aq)

8. Voltaic Cells
Alkaline Batteries: used in many common small devices; anode is a Zn metal and KOH paste
ox: Zn(S) + 2OH –(aq) → Zn(OH)2(aq) + 2e –
red: 2MnO2(s) + H2O(l) + 2e– → Mn2O3(s) + 2OH –(aq)

9. Voltaic Cells
Mercury Batteries: tiny batteries found in hearing aids, calculators and camera flashes
ox: Zn(S) + 2OH –(aq) → Zn(OH)2(aq) + 2e –
red: HgO(s) + H2O(l) + 2e– → Hg(l) + 2OH –(aq)

10. Voltaic Cells
A fuel cell is a voltaic cell that could, in principle, work forever because the reactants are being continuously supplied and the products are being continuously removed

11. Voltaic Cells x varies, which affects the color of rust
Prevention of corrosion: Iron (Fe) rust forms by the following reaction:
4Fe + 3O2 + xH2O → 2Fe2O3∙xH2O
x varies, which affects the color of rust
ox: Fe(S) → Fe3+(aq) + 3e –
red: O2(g) + 2H2O(l) + 4e– → 4OH –(aq)

12. Voltaic Cells
Prevent corrosion of steel by coating it with zinc , which is called galvanizing
Another name for coating the steel with a sacrificial anode is called cathodic protection
Two metals that can protect Fe from rusting are Zn and Mg

13. Electrical Potential
The tendency for a half-reaction to occur as reduction in an electrochemical cell is called the reduction potential
The difference in potential between an electrode and its solution is called its electrode potential
The potential of a half-cell under standard conditions measured relative to the standard hydrogen electrode (SHE) is called the standard electrode potential (E˚); the SHE voltage is 0

14. E˚cell = E˚cathode - E˚anode
Electrical Potential
Electrode potentials are listed as potentials of reduction
Oxidizing agents have postitive E˚ values; reducing agents have negative E˚ values
E˚cell = E˚cathode - E˚anode
E˚ values come from the standard reduction potentials chart

15. Electrical Potential
Example Problem: Calculate the cell potential for the following redox reaction.
3Ag+ + Fe → 3Ag + Fe+3

16. Electrolytic Cells
If electrical energy is required to produce a redox reaction and bring about chemical change in an electrochemical cell, it is an electrolytic cell
Two differences between voltaic and electrolytic cells:
The anode and cathode in an electrolytic cell are connected to a power source (battery), whereas a voltaic cell serves as a source of electrical energy
Spontaneous redox reactions produce electricity in voltaic cells, whereas an external electrical energy source is needed to produce nonspontaneous redox reactions in electrolytic cells.

17. Electrolytic Cells
In an electrolytic cell, electrical energy is converted to chemical energy.
In a voltaic cell, chemical energy is converted to electrical energy

18. Electrolytic Cells
An electrolytic process in which a metal ion is reduced and a solid metal is deposited on a surface is called electroplating

19. Electrolytic Cells
Rechargeable Cells: both voltaic and electrolytic cells; the standard 12V car battery is a set of 6 rechargeable cells
ox: Pb(S) + SO4 2–(aq) → PbSO4(s) + 2e –
red: PbO2(s) + 4H+ + SO4 2–(aq) + 2e– → PbSO4(s) + 2H2O(l)

20. Electrolytic Cells
Electrolysis: is the process of passing a current through a cell for which the cell potential is negative, causing an oxidation-reduction reaction to occur
ox: 6H2O(l) → O2(g) + 4H3O +(aq) + 4e –
red: 4H2O(l) + 4e– → 2H2 (g) + 4OH –(aq)