1. Redox Reactions
Chapter 18
+ O2 

2. Oxidation-Reduction (Redox) Reactions
“redox” reactions: rxns in which electrons are transferred from one species to another
oxidation & reduction always occur simultaneously
we use OXIDATION NUMBERS to keep track of electron transfers

3. Rules for Assigning Oxidation Numbers:
1) the ox. state of any free (uncombined) element is zero.
Ex: Na, S, O2, H2, Cl2, O3

4. Rules for Assigning Oxidation Numbers:
2) The ox. state of an element in a simple ion is the charge of the ion.
Mg2+  oxidation of Mg is +2

5. Rules for Assigning Oxidation Numbers:
3) the ox. # for hydrogen is +1
(unless combined with a metal, then it has an ox. # of –1)
Ex: NaOH (H bonded to O) v. NaH (H bonded to Na)

6. Rules for Assigning Oxidation Numbers:
4) the ox. # of fluorine is always –1.

7. Rules for Assigning Oxidation Numbers:
5) the ox. # of oxygen is usually –2.
Why USUALLY? Not -2 when it’s in a peroxide, such as hydrogen peroxide:
H2O2

8. Rules for Assigning Oxidation Numbers:
6) in any neutral compound, the sum of the oxidation #’s = zero.

9. Rules for Assigning Oxidation Numbers:
7) in a polyatomic ion, the sum of the oxidation #’s = the overall charge of the ion.

10. Rules for Assigning Oxidation Numbers:
**use these rules to assign oxidation #’s; assign known #’s first, then fill in the #’s for the remaining elements:

11. Examples: Assign oxidation #’s to each element:
a) NaNO3

12. Examples: Assign oxidation #’s to each element:
b) SO32-

13. Examples: Assign oxidation #’s to each element:
c) HCO3-

14. Examples: Assign oxidation #’s to each element:
d) H3PO4

15. Examples: Assign oxidation #’s to each element:
e) Cr2O72-

16. Examples: Assign oxidation #’s to each element:
f) K2Sn(OH)6

17. Definitions
Oxidation: the process of losing electrons (ox # increases)
Reduction: the process of gaining electrons (ox # decreases)
Oxidizing agents: species that cause oxidation (they are reduced)
Reducing agents: species that cause reduction (they are oxidized)

18. To help you remember…
OIL RIG
Oxidation Is Loss
Reduction Is Gain

19. Are all rxns REDOX rxns?
a reaction is “redox” if a change in oxidation # happens; if no change in oxidation # occurs, the reaction is nonredox.

20. Examples
MgCO3  MgO + CO2

21. Examples
Zn + CuSO4  ZnSO4 + Cu

22. Examples
NaCl + AgNO3  AgCl + NaNO3

23. Examples
CO2 + H2O  C6H12O6 + O2

24. Balancing Redox Equations

25. Balancing Redox Equations
In balancing redox equations, the # of electrons lost in oxidation (the increase in ox. #) must equal the # of electrons gained in reduction (the decrease in ox. #)
There are 2 methods for balancing redox equations:
Change in Oxidation-Number Method
The Half-Reaction Method

26. 1. Change in Oxidation-Number Method:
based on equal total increases and decreases in oxidation #’s
Steps:
Write equation and assign oxidation #’s.
Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.
Connect the atoms that change ox. #’s using a bracket; write the change in ox. # at the midpoint of each bracket.
Choose coefficients that make the total increase in ox. # = the total decrease in ox. #.
Balance the remaining elements by inspection.

27. Example
S HNO3  SO2 + NO H2O

28. If needed, reactions that take place in acidic or basic solutions can be balanced as follows:
add H2O to the side needing oxygen
balance as if in acidic sol’n
then add H+ to balance the hydrogen
add enough OH- to both sides to cancel out each H+ (making H2O) & then cancel out water as appropriate

29. Example: Balance the following equation, assuming it takes place in acidic solution.
ClO4- + I-  Cl- + I2

30. 2. The Half-Reaction Method:
separate and balance the oxidation and reduction half-reactions.
Steps:
Write equation and assign oxidation #’s.
Determine which element is oxidized and which is reduced, and determine the change in oxidation # for each.
Construct unbalanced oxidation and reduction half reactions.
Balance the elements and the charges (by adding electrons as reactants or products) in each half-reaction.
Balance the electron transfer by multiplying the balanced half-reaction by appropriate integers.
Add the resulting half-reaction and eliminate any common terms to obtain the balanced equation.

31. Example: Balance the following using the half-reaction method:
HNO3 + H2S  NO + S + H2O