1. Electrochemistry
Chapter 18

2. Oxidation State and Oxidation-Reduction Reactions
Redox reactions are one of the reaction types covered in Ch. 4
For the rules associated with assigning oxidation numbers see Pg. 149
Analogous to acid-base reactions
Oxidant
Reductant
Ex:
Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)

3. Identifying Oxidizing and Reducing Agents
Determine the oxidation numbers of the elements in the following reaction. Identify the oxidant and reductant.
5H2C2O4(aq) + 6H3O+(aq) + 2MnO4-(aq) → CO2(g) + 2Mn+(aq) + 14H2O(l)

4. Balancing Oxidation-Reduction Reactions
When balancing typical chemical reactions, mass must be conserved
However, charge must be conserved as well
Redox reactions must be broken up into two half reactions which must be balanced separately
Two half reactions: One for oxidation; one for reduction

5. Balancing Redox Reactions (cont.)
Two types of solutions: acidic and basic
Rules are only slightly different
Divide equation into two half reactions
Balance each half reaction
All elements other than O and H
Balance O by adding H2O
Balance H by adding H+
Balance charge by adding electrons, e-
Multiply each half reactions with integer values in order to balance the # of electrons
Combine both half reactions

6. Redox Reactions in Acidic Solution
Balance the following redox reaction which takes place in acidic solution:
Fe2+(aq) + Cr2O72-(aq) → Cr3+(aq) + Fe3+(aq)

7. Redox Reactions in Basic Solution
Exactly the same as acidic solution
Only difference is that OH- must be used to “neutralize” H+ added to balance H’s
No appreciable amount of H+ can exist in basic solution
Ex:
MnO4-(aq) + Br-(aq) → MnO2(s) + BrO3-(aq)

8. Galvanic Cells Section 18.1
Spontaneous redox reactions are used to perform electrical work using a voltaic cell
Voltaic cells (sometimes called galvanic cells) require an oxidant and reductant and some sort of medium for the electrons to travel through in order to produce current

9. Voltaic Cell Diagram
Current flow: anode to cathode
Each half of the voltaic cell is referred to as a half-cell
Cathode
Anode
Overall Rxn: Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s)

10. Describing a Voltaic Cell
Consider the cell shown below.
a.) Write the half-reaction that occurs
in each half cell
b.) Write the equation for the overall
chemical change taking place
c.) Identify cathode and anode
d.) Identify the signs of the electrodes
e.) Indicate the direction that the nitrate
ions flow through the salt bridge

11. Cell EMF Under Standard Conditions Section 18.2
In a voltaic cell, there is a potential difference between the two electrodes (electromotive force or cell potential)
For a spontaneous reaction Ecell will have a positive value
Under standard conditions (1 M or 1 atm) it is referred to as standard cell potential, Ecell

12. Standard Reduction Potentials
In theory, we could experimentally determine the Ecell values for every possible redox combination (dumb idea—too much work)
More practically, we can determine the voltage required for reduction to occur at an electrode
Referred to as standard reduction potentials and are referenced to the standard hydrogen electrode (SHE)
Ered = 0

13. Ecell = Ered(cathode) - Ered(anode)
Calculating Ecell
The cell potential, Ecell, can be calculated using tabulated standard reduction potentials:
Ecell = Ered(cathode) - Ered(anode)
For all reactions, the reaction is spontaneous if Ecell is positive
The more positive the value, the more driving force there is for the reaction to occur

14. Standard Reduction Potentials

15. Calculating Ecell from Standard Reduction Potentials
Calculate the standard potential and state the direction in which the reaction proceeds spontaneously for:
2Cr+3(aq) + 2Br-(aq) → 2Cr+2(aq) + Br2(l)
See Interactive Example 18.1 (Pg. 711)

16. Determining Half-Reactions at Electrodes and Calculating Cell EMF
A voltaic cell is based on a Co2+/Co half cell and an AgCl/Ag half cell.
a.) What half-reaction occurs at the anode?
b.) What is the standard cell potential?
See Interactive Example 18.2 (Pg. 713)

17. Relative Strengths of Oxidizing Agents

18. Free Energy and Redox Reactions Section 18.3
We know that a negative G value indicates a spontaneous reaction, and we also know that a positive emf, E, also indicates spontaneity
Therefore there must be a link between the two:
G = -nFE or G = -nFE (std. cond.)
F = Faraday’s constant = 96,485 J/V mol
n = # of electrons transferred in the reaction

19. AgCl(s) + Fe+2(aq)  Ag(s) + Fe+3(aq) + Cl-(aq)
Determining G
Calculate the standard free energy change for the reaction below:
AgCl(s) + Fe+2(aq)  Ag(s) + Fe+3(aq) + Cl-(aq)
See Interactive Example 18.3 (Pg. 716)

20. Sn2+(aq) + Ni(s)  Sn(s) + Ni2+(aq)
Determining G and K
Calculate the equilibrium constant for the reaction below:
Sn2+(aq) + Ni(s)  Sn(s) + Ni2+(aq)

21. Cell emf Under Nonstandard Conditions Section 18.4
Cell potential under nonstandard conditions can be calculated using the Nernst equation which is derived from the equation describing free energy change:
G = G + RT ln(Q)
Nernst equation:

22. Voltaic Cell EMF under Nonstandard Conditions
A voltaic cell consists of a half-cell of iron ions in a solution with [Fe2+] = 2.0 M and [Fe3+] = 0.75 M, and a half-cell of copper metal immersed in a solution containing Cu2+ at a concentration of 3.6 x 10-4 M. What is the voltage of this cell?
See Interactive Example 18.7 (Pg. 720)

23. Calculating Concentration in a Voltaic Cell
What is the pH of the solution in the cathode compartment of the reaction shown below when PH2 = 1.0 atm, [Zn2+] in the anode compartment is 0.10 M, and the cell emf (Ecell) is V?
Zn(s) + 2H+(aq) → Zn2+(aq) + H2(g)

24. Electrolysis Section 18.7
Voltaic cells take advantage of spontaneous electrochemical reactions to produce electrical current
The exact opposite occurs during electrolysis

25. Quantitative Aspects of Electrolysis
Consider the reactions below:
Li+ → Li(s)
Mg2+ → Mg(s)
If the amount of charge applied to an electrolytic cell is known, the mass of the electroplated material can be calculated
Coulombs = amperes x seconds

26. Relating Electrical Charge and Quantity of Electrolysis
A constant current of A passes through a silver nitrate solution for 90.0 minutes. What mass of silver metal is deposited at the anode?
See Interactive Example 18.9 (Pg. 730)