Chapter 18 Oxidation – Reduction Reactions
18.1 Defining Oxidation and Reduction
(p. 713 – 720)
OIL RIG
Oxidation Is Loss of one or more electrons by a substance
Oxidation number increases (toward +10)
Substances are oxidized when they react with oxygen
Rusting  iron + oxygen
Combustion  organic molecules + oxygen
Burning Magnesium  2Mg + O2  MgO + light
Fruit (apples, bananas) change colour and texture when oxidized
Ethanol forms ethanoic acid (vinegar) when it’s oxidized
Reduction Is Gain of one or more electrons by a substance
Oxidation number is reduced (toward –10)
Oxygen is often reduced
O20 + 4 e–  2 O2–
Redox Reactions (Oxidation – Reduction Reactions)
-
must involve the transfer of electrons and changes in oxidation number
The electrons lost by one substance = the number of electrons gained by another
Redox reactions can be written as two half-reactions  oxidation (OX) and reduction (RED)
Oxidizing Agent = the substance that helps the oxidation half-reaction to happen
 gains (accepts) electrons and gets reduced. (usually oxygen)
Reducing Agent = the substance that helps the reduction half-reaction to happen
 loses (donates) electrons and gets oxidized.
ex.
2Mg(s) + O2(g)  2MgO(s)
OX
2MgO(s) = 2Mg2+ + 2O2--
2 Mg  2 Mg2+ + 4 e–
Reducing agent: Mg
Practice Problems:
#1–8
RED O2 + 4e–  2 O2–
Oxidizing agent: O2
(p. 715, 716)
Reaction Types (p. 31 – 33)
Combination (synthesis)
one product
2 H2(g) + O2 (g)  2 H2O(g) + energy
Decomposition
one reactant
2 H2O(g) + energy  2 H2(g) + O2 (g)
Single Replacement
1 element + 1 compound
Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g)
Double Replacement
2 compounds
Pb(NO3)2(aq) + 2NaOH(aq)  2NaNO3 (aq) + Pb(OH)2(s)
Combustion
X + O2  CO2 + H2O
All of these are REDOX, EXCEPT Double Replacement  ions change partners, not charges (oxidation numbers).
Any reaction that has an elemental form is a redox reaction.
Activity Series
More Reactive Elements
Less Reactive Elements
 form compounds; form ions
 form elemental products
ex.1
Mg (s) + 2 HCl (aq)  MgCl2 (aq) + H2 (g)
Magnesium changes to ion (forms the compound)  more reactive than hydrogen
Hydrogen ions change to hydrogen gas (elemental form)  less reactive than magnesium
ex. 2
NaF (aq) + Cl2 (g)  no reaction
Fluorine stays as fluoride ion in the compound  more reactive than chlorine
Chlorine stays in its elemental form  less reactive than fluorine
Activity Trends
Noble metals – Gold, Platinum , Silver
– tend to remain in elemental form and resist oxidation
– do not react with acids  Au is less reactive than H
exception: Gold will react with aqua regia (royal water)
Aqua regia = 1 part concentrated nitric acid + 3 parts concentrated hydrochloric acid
Base Metals
– iron, lead, Group 1A, and 2A
– tend to combine react easily with oxygen (rust)
– tend to react with other elements to form compounds (form ions)
– react with acids  more reactive than hydrogen
Reactivity Series can be determined using a series of single replacement reactions
If there is no reaction, the elemental form must be less reactive.
If there is a reaction, the elemental form must have been more reactive.
See Reactivity Series on p. 718
•
Given a set of reactions, you are expected to be able to determine the order of reactivity.
Sample Problem
Reaction
A + BX  No reaction
A + DX  AX + D
C + BX  No reaction
AX + C  A + CX
Conclusions
B is more reactive than A
A is more reactive than D
B is more reactive than C
C is more reactive than A
B>C>A>D
18.2 Oxidation Numbers
Oxidation Numbers –
–
(p. 721 – 729)
can be either the actual charge (ions) or the hypothetical charge of each atom present.
used to determine which element is oxidized and which is reduced
Rules for Assigning Oxidation Numbers (see Table 18.1 Oxidation Number Rules p. 724)
Practice Problems
Assigning Oxidation Numbers
# 9 – 12
p. 725 –726
Practice Problems
Identifying Redox Reactions
# 13 – 16
p. 727 –728
Section Review
# 2, 7, 8
18.3 The Half-Reaction Method for Balancing Equations
• Write and balance half-reactions and net reactions
The number of electrons lost in the oxidation half-reaction must equal the number of electrons gained in the
reduction half-reaction.
For each half-reaction, multiply each term by a factor that produces equal numbers of electrons lost and gained.
Ex. Co(NO3)3 + Cd  Co(NO3)2 + Cd(NO3)2
1. remove spectator ions and assign oxidation numbers
Co 3+ + Cd 0  Co 2+ + Cd 2+
2. Write half-reactions
a. Oxidation half-reactions
b. Reduction half-reactions
Cd 0  Cd 2+ + 2 eCo 3+ + 1 e-  Co 2+
3. Balance electrons
OX
RED
1 x OX
2 x RED
Cd 0  Cd 2+ + 2 e2 Co 3+ + 2e-  2Co 2+
4. Write net reaction (check total charge on each side)
2 Co 3+ + Cd 0  2 Co 2+ + Cd 2+
5. Write net equation with spectators (optional)
2Co(NO3)3 + Cd  2 Co(NO3)2 + Cd(NO3)2
Variations
Acid Solutions
Balance any atoms other than H and O
Add water (H20) to balance O
Add hydrogen ions (H+) to balance H
Basic Solutions
Balance any atoms other than H and O
Add water (H20) to balance O
Add hydrogen ions (H+) to balance H
Add the same number of hydroxide ions (OH-) to both sides
Change (H+ + OH-) to H2O
Write net number of water molecules
Practice Problems
Balancing a Half-Reaction in Acidic Solution
# 17 – 21
p. 732
Practice Problems
Balancing a Half-Reaction in Basic Solution
# 22 – 24
p. 733 – 734
Practice Problems
Balancing a Redox Reaction in Acidic Solution
# 25 – 28
p. 737 – 739
** Not responsible for redox titrations (p. 742 – 744)