Chapter 4 – Types of
Chemical Reactions and
Solution Stoichiometry
4.9 – 4.10 Notes
AP Chemistry
Oxidation-Reduction Reactions
Reactions in which one or more electrons
are transferred are called oxidationreduction reactions.
 Examples:
Mg(s) + S(s)  MgS(s)
CH4(g) + O2(g)  CO2(g) + 2H2O(l) + E

Oxidation States

Imaginary charges atoms would have if
covalently bound atoms transferred electrons or
actual charges on ions in ionic compounds.
(Assume the more electronegative atom takes
on negative charge)

Examples:
CO2: C=4+, O=2O2: O = 0 (equally shared)
PCl5: P=5+, Cl=1-
Rules for Assigning Oxidation States
1.
2.
3.
4.
5.
6.
An atom in elemental form receives an oxidation state
of 0. (O2, Na, Fe, Cl2 = 0)
Oxidation states for monatomic ions are the same as
their charge. (Na+=1+, O2-=2-)
In its compounds, fluorine always has an oxidation
state of 1-. (HF, F=1-)
In covalent compounds with nonmetals, hydrogen is
assigned an oxidation state of 1+. (HCl, NH3, H=1+)
Oxygen is usually assigned 2-. Exceptions include
peroxides (O2 group O2=2-) and compounds containing
fluorine (OF2. O=2+).
Sum of the oxidation numbers must equal zero, unless
substance is an ion.
Characteristics of Redox Reactions

Oxidation and reduction occur
simultaneously in a redox reaction.
– If a species is being oxidized, another species
must be reduced.
Oxidation – increase in oxidation state
(lose of electrons)
 Reduction – decrease in oxidation state
(gain of electrons)

Oxidizing and Reducing Agents
Oxidizing Agent – substance being
reduced
 Reducing Agent – substance being
oxidixed

Summary

Oxidized
– Loses electrons
– Oxidation state
increases
– Reducing agent

Reduced
– Gains electrons
– Oxidation state
decreases
– Oxidizing agent
Example

Identify the oxidizing and reducing agent
in the following redox reaction.
2Al(s) + 3I2(s)  2AlI3(s)
Al: Al0  Al3+, loses e-, is oxidized, and is
the reducing agent
 I2: I20  I-, gains e-, is reduced, and is the
oxidizing agent

Balancing Oxidation-Reduction
Reactions



Half-Reaction Method.
Different for acidic and basic media.
Involves analyzing oxidation and
reduction as two separate half reactions.
Half-Reactions
Consider the following redox reaction:
PbO(s) + CO(g)  Pb(s) + CO2(g)
 Oxidation Half-Reaction:
C2+  C4+ + 2e Reduction Half-Reaction:
Pb2+ + 2e-  Pb0

Steps for Balancing Redox
Reactions in an Acidic Solution
1.
2.
Write separate equations for the oxidation and
reduction half-reactions.
For each half-reaction:
1.
2.
3.
4.
3.
4.
5.
Balance
Balance
Balance
Balance
all the elements except hydrogen and oxygen
oxygen using H2O
hydrogen using H+
the charges using electrons
If necessary, multiply each half-reaction by an integer
to equalize the number of electrons transferred in the
two half-reactions.
Add the half-reactions, and cancel identical species.
Check that elements and charges are balanced.
Balancing Redox Equations in an
Acidic Solution
Example:
Balance the following redox equation.
Cu(s) + NO3-(aq)  Cu2+(aq) + NO(g)


Solution:
3Cu(s) + 8H+(aq) + 2NO3-(aq) 
3Cu2+(aq) + 2NO(g) + 4H2O
Steps for Balancing Redox
Reactions in an Basic Solution
1.
2.
3.
4.
Use half-reaction method as if balancing
in an acidic solution with H+ ions.
Add OH- ions to both sides in order to
change H+ ions to H2O.
Cancel out water molecules that appear
on both sides of the equation.
Check that elements and charges are
balanced.
Balancing Redox Equations in a
Basic Solution
Example:
Balance the following redox equation.
Cr(s) + CrO42-(aq)  Cr(OH)3(s)


Solution:
4H2O(l) + Cr(s) + CrO42-(aq) 
2Cr(OH)3(s) +2OH-(aq)