AP Chapter Five Outline
Chemical Reactions
I.
Exchange Reactions: Precipitation and Net Ionic Equations
 When ionic compounds dissolve in water, they dissociate into ions surrounded by
water molecules. An ionic compound that completely dissolves into ions is a
strong electrolyte.
A. Exchange Reactions: AB + CD  AD + CB
1. If both reactants and products are water-soluble compounds, then
no overall reaction will take place.
2. If, however, one of the potential products of the reaction removes
ions from the solution, a reaction will occur.
a) formation of an insoluble ionic compound
AgNO3 (aq) + KCl (aq)  KNO3 (aq) + AgCl (s)
b) Formation of a molecular compound that remains in
solution. Commonly water is produced in acid-base
neutralization reactions
H2SO4 (aq) + 2NaOH (aq)  Na2SO4 (aq) + 2H2O (l)
c) formation of a gaseous molecular compound that evolves
from the solution
2 HCl (aq) + Na2S (aq)  2 NaCl (aq) + H2S (g)
B. Precipitation Reactions: when in an exchange reaction an insoluble
product – a precipitate – is formed
BaCl2 (aq) + Na2SO4 (aq)  BaSO4 (s) + 2NaCl (aq)
C. Net Ionic Equations: Equations that includes only the symbols or
formulas of ions in solution or compounds that undergo change. Spectator
ions, or ions that are present in the reaction but to not undergo any change
are ignored.

To Write a Net Ionic Equation:
1) Write the overall balanced equation
2) Determine solubilities of reactants and products
3) All soluble ionic compounds dissociate into their component ions in
aqueous solutions
4) Write a complete ionic equation
5) Cancel out spectator ions from both sides of the equation
6) Check that the sum of the charges is the same on each side of the equation
Example: write a net ionic equation for the reaction between aqueous
magnesium sulfate and aqueous sodium oxalate
MgSO4 (aq) + Na2C2O4 (aq)  MgC2O4 (s) + Na2SO4 (aq) (steps I and II)
Mg2+(aq) + SO42- (aq) + 2Na+(aq) C2O42- (aq)  MgC2O4 (s) + 2 Na+ (aq) +
SO42- (aq) (step IV)
Mg 2+ (aq) + C2O42-(aq)  MgC2O4 (s)
II.
Acids, bases, and Acid-Base Exchange Reactions
A. Acid: any substance that increases the concentration of hydrogen ions,
H+, when dissolved in pure water. The H+ ions combines with H2O to form
H2O+, the hydronium ion.
1. Acids that completely ionize in water are strong electrolytes and
stron acids.
Example: HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq)
2. Acids that ionize only slightly are weak electrolytes and weak
acids
Example: CH3COOH (aq) + H2O (l)  H3O+ (aq) + CH3COO- (aq)
B. Base: a substance that increases the concentration of the hydroxide ion,
OH-, when dissolved in water
C. Electrolytes
1. Strong electrolytes are either ionic compounds (salts or strong
bases) or molecular compounds that are strong acids and ionize
completely in aqueous solutions.
2. Weak electrolytes are molecular compounds that are weak acids
or bases and establish equilibrium with water.
3. Nonelectrolytes are molecular compounds that do not ionize in
aqueous solutions.
D. Neutralization Reactions: When a strong acid and a strong base are
mixed, they neutralize each other. The acid donates a H+ ion, the base
donates a OH- ion, and together they form water, a molecular compound.
Example: HX (aq) + MOH(aq)  HOH (l) + MX (aq)
Acid
+ Base
Water
Salt
E. Gas-Forming Exchange Reactions: Formation of a gas removes a
molecular product from a solution
Example: NaHCO3 (s) + CH3COOH (aq)  H2O (l) + CO2 (g) +
CH3COONa(s)
III.
Oxidation and Reduction Reactions
 Oxidation and Reduction reactions involve transfer of electrons from one reactant
to another.
 When a substance accepts electrons, it is reduced
 When a substance loses electrons it is oxidized
 LEO : Losing electrons = oxidation
 GER: Gaining electrons = reduction
 OIL RIG: Oxidation is loss, reduction is gaining
 The substance that is reduced causes the other species to be oxidized and is called
the oxidizing agent.
 The substance that is oxidized causes the other species to be reduced and is called
the reducing agent.
A. Oxidation Numbers and Redox Reactions
 Oxidation numbers compare the charge of an uncombined atom with
its actual charge in a compound. All neutral atoms have an equal
number of protons and electrons and thus have no net charge.
 Oxidation numbers of atoms in molecular compound are assigned as
though electrons were completely transferred to form ions.
B. Rules for Determining Oxidation Numbers:
1. The oxidation number of an atom of a pure element is zero.
2. The oxidation of a monatomic ion equals its charge.
3. Hydrogen’s oxidation number is +1 except when it is an anion ( a
hydride) in which case it is -1.
4. Fluorine’s oxidation number is -1.
5. Halogens other than fluorine have an oxidation number of -1
except in interhalogen compounds.
6. Oxygen’s oxidation number is -2 except in peroxides when it is 1.
7. In binary compounds atoms of group 6A have an oxidation of -2
except when combined with oxygen or halogens.
8. The sum of the oxidation numbers in a neutral compound is zero.
9. The sum of the oxidation numbers of a polyatomic ion equals the
charge on the ion.
IV. Displacement Reactions, Redox, and the Activity Series