1. Principles of Reactivity: Electron Transfer Reactions
Chemistry and Chemical Reactivity 6th Edition
John C. Kotz Paul M. Treichel Gabriela C. Weaver
Principles of Reactivity: Electron Transfer Reactions
Lectures written by John Kotz
© 2006 Brooks/Cole Thomson

2. ELECTROCHEMISTRY Chapter 19

3. TRANSFER REACTIONS Atom/Group transfer Electron transfer
HCl + H2O ---> Cl H3O+
Electron transfer
Cu(s) Ag+(aq) ---> Cu2+(aq) Ag(s)

4. Electron Transfer Reactions
Electron transfer reactions are oxidation-reduction or redox reactions.
Redox reactions can result in the generation of an electric current or be caused by imposing an electric current.
Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

5. Review of Terminology for Redox Reactions
OXIDATION—loss of electron(s) by a species; increase in oxidation number.
REDUCTION—gain of electron(s); decrease in oxidation number.
OXIDIZING AGENT—electron acceptor; species is reduced.
REDUCING AGENT—electron donor; species is oxidized.

6. OXIDATION-REDUCTION REACTIONS
Direct Redox Reaction
Oxidizing and reducing agents in direct contact.
Cu(s) Ag+(aq) ---> Cu2+(aq) Ag(s)

7. Cu + Ag+ --give--> Cu2+ + Ag
Balancing Equations
Cu + Ag give--> Cu Ag

8. Balancing Equations
Step 1: Divide the reaction into half-reactions, one for oxidation and the other for reduction.
Ox Cu ---> Cu2+
Red Ag+ ---> Ag
Step 2: Balance each for mass. Already done in this case.
Step 3: Balance each half-reaction for charge by adding electrons.
Ox Cu ---> Cu e-
Red Ag+ + e- ---> Ag

9. Balancing Equations
Step 4: Multiply each half-reaction by a factor so that the reducing agent supplies as many electrons as the oxidizing agent requires.
Reducing agent Cu ---> Cu e-
Oxidizing agent 2 Ag e- ---> 2 Ag
Step 5: Add half-reactions to give the overall equation.
Cu Ag > Cu Ag
The equation is now balanced for both charge and mass.

10. Balancing Equations for Redox Reactions
Some redox reactions have equations that must be balanced by special techniques.
MnO Fe H+---> Mn Fe H2O
Mn = +7
Fe = +2
Mn = +2
Fe = +3

11. Reduction of VO2+ with Zn

12. Add H2O on O-deficient side and add H+ on other side for H-balance.
Balancing Equations
Balance the following in acid solution—
VO Zn ---> VO Zn2+
Step 1: Write the half-reactions
Ox Zn ---> Zn2+
Red VO > VO2+
Step 2: Balance each half-reaction for mass.
Red
2 H+ +
VO > VO H2O
Add H2O on O-deficient side and add H+ on other side for H-balance.

13. Balancing Equations Step 3: Balance half-reactions for charge.
Ox Zn ---> Zn e-
Red e H+ + VO > VO H2O
Step 4: Multiply by an appropriate factor.
Red 2e H VO > 2 VO H2O
Step 5: Add balanced half-reactions
Zn H VO > Zn VO H2O

14. Tips on Balancing Equations
Never add O2, O atoms, or O2- to balance oxygen ONLY add H2O or OH-.
Never add H2 or H atoms to balance hydrogen ONLY add H+ or H2O.
Be sure to write the correct charges on all the ions.
Check your work at the end to make sure mass and charge are balanced.
PRACTICE!

16. Potential Ladder for Reduction Half-Reactions
Figure 20.14
Best oxidizing agents
Best reducing agents
Potential Ladder for Reduction Half-Reactions

17. TABLE OF STANDARD REDUCTION POTENTIALS
oxidizing
ability of ion E o
(V) Cu 2+
+ 2e Cu
+0.34
2 H +
+ 2e H
0.00 Zn
+ 2e Zn
-0.76
reducing ability
of element 2

18. Using Standard Potentials, Eo Table 20.1
Which is the best oxidizing agent: O2, H2O2, or Cl2? _________________
Which is the best reducing agent: Hg, Al, or Sn? ____________________

19. Standard Redox Potentials, Eo
Any substance on the right will reduce any substance higher than it on the left.
Zn can reduce H+ and Cu2+.
H2 can reduce Cu2+ but not Zn2+
Cu cannot reduce H+ or Zn2+.

20. Standard Redox Potentials, Eo
Ox. agent Cu 2+
+ 2e- --> Cu
+0.34 +
2 H
+ 2e- --> H2
0.00 Zn
+ 2e- --> Zn
-0.76
Red. agent
Any substance on the right will reduce any substance higher than it on the left.
Northwest-southeast rule: product-favored reactions occur between
reducing agent at southeast corner
oxidizing agent at northwest corner

21. 2 H+(aq, 1 M) + 2e- <----> H2(g, 1 atm)
CELL POTENTIALS, Eo
Can’t measure 1/2 reaction Eo directly. Therefore, measure it relative to a STANDARD HYDROGEN CELL
2 H+(aq, 1 M) e- <----> H2(g, 1 atm)
Eo = 0.0 V

22. Calculating Cell Voltage
Balanced half-reactions can be added together to get overall, balanced equation.
Zn(s) ---> Zn2+(aq) + 2e-
Cu2+(aq) + 2e- ---> Cu(s)
Cu2+(aq) + Zn(s) ---> Zn2+(aq) + Cu(s)
If we know Eo for each half-reaction, we could get Eo for net reaction.

23. Uses of Eo Values
Organize half-reactions by relative ability to act as oxidizing agents
Use this to predict direction of redox reactions and cell potentials.
Cu2+(aq) + 2e- ---> Cu(s) Eo = V
Zn2+(aq) + 2e- ---> Zn(s) Eo = –0.76 V
Note that when a reaction is reversed the sign of E˚ is reversed!

24. Using Standard Potentials, Eo
In which direction do the following reactions go?
Cu(s) Ag+(aq) ---> Cu2+(aq) Ag(s)
Goes right as written
2 Fe2+(aq) + Sn2+(aq) ---> 2 Fe3+(aq) + Sn(s)
Goes LEFT opposite to direction written
What is Eonet for the overall reaction?

25. Eo and Thermodynamics ∆Go = -nFEo
Eo is related to ∆Go, the free energy change for the reaction.
∆G˚ is proportional to –nE˚
∆Go = -nFEo
where F = Faraday constant = x 104 J/V•mol of e-
(or x 104 coulombs/mol)
and n is the number of moles of electrons transferred

26. Eo and ∆Go ∆Go = - n F Eo For a product-favored reaction
Reactants ----> Products
∆Go < 0 and so Eo > 0
Eo is positive
For a reactant-favored reaction
Reactants <---- Products
∆Go > 0 and so Eo < 0
Eo is negative

27. Eo and Equilibrium Constant
DGo = -RT ln K
DGo = -nFEo