VII. IONIC REACTIONS IN AQUEOUS SOLUTIONS
A. Electrolytes
1. Electrolyte: Substances that dissociate, or ionize, in water to produce “free” ions.
They include salts, acids and bases.
a) Strong electrolyte: ionize completely (or nearly completely) in water. They
include soluble salts, strong acids and strong bases.
NOTE: Soluble and insoluble are not quantitative terms, they are practical or relative terms. Nothing is
completely insoluble, and the solubility of most substances has a limit.
Example(1): NaCl in water
Example(2): MgCl2
Example(3): Na2SO4
b) Weak electrolytes: ionize only partially in water. They include insoluble salts,
weak acids and weak bases.
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IONIC REACTIONS / Electrolytes
Example(4): AgCl
c) Chemical equilibrium: The amounts of products and reactants don’t change,
but the reaction is still taking place.
Example(5): Mg(OH)2
2. Acids: Substances that dissolve in water and ionize to give “H+” and an anion.
Example(6): HCl
a) Strong acids: ionize completely (or nearly) in water.
b) Common strong acids:
c) Weak acids: ionize only partially in water.
77
IONIC REACTIONS / Electrolytes
Example(7): HF
d) Common weak acids:
3. Bases: Substances that dissociate or ionize in water to give OH- and a cation.
a) Strong bases: ionize completely (or nearly) in water.
b) Common strong bases:
c) Weak bases: dissociate or ionize only partially in water.
d) Common weak bases:
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IONIC REACTIONS / Double Replacement
B. Metathesis (Double Replacement) Reactions
Example(1): AgNO3 +
NaCl
 AgCl + NaNO3
1. Ionic and net ionic equations
a) Ionic equation: All strong electrolytes are written as separate ions, weak
electrolytes and nonelectrolytes are written in molecular form.
b) Net ionic equation: Only the things that change are shown. The ions that do
not change (spectator ions) are eliminated from the equation.
Example(2): Write the ionic and net ionic equations for:
NaCl  AgCl + NaNO3
AgNO3 +
Example(3): Write the ionic and net ionic equations for:
2Na3PO4
+
3CaCl2  Ca3(PO4)2 + 6NaCl
79
IONIC REACTIONS / Double Replacement
Example(4): Write the ionic and net ionic equations for:
CaCO3
2HCl  CaCl2 + CO2 + H2O
+
2. Precipitation reactions with salts
a) criteria for a ppt reaction:
soluble + soluble  insoluble + soluble
Example(5): Will the reaction proposed below occur? Write the ionic and net ionic equations.
Pb(NO3)2
+
2NaI 
PbI2
+
2NaNO3
Example(6): Will the reaction proposed below occur? Write the ionic and net ionic equations.
Al(NO3)3
+
3NaC2H3O2  Al(C2H3O2)3
+
3NaNO3
b) determining the formulas for the products
The formulas of the products are determined by the CHARGES on the ions, NOT
by making the equation balance.
Example(7): Give the products of the reaction below.
K2CO3
+
BaCl2 
Example(8): Give the products of the reaction below.
Na2S
+
Fe(NO3)3 
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IONIC REACTIONS / Double Replacement
3. Neutralization reactions:
acid
+
base 
salt
+
H2O
NOTE: Unless told otherwise, assume complete neutralization.
Example(9): Give the products of the reaction below, then write the ionic and net ionic equations.
HCl
+
NaOH 
Example(10): Give the products of the reaction below, then write the ionic and net ionic equations.
H2SO4
+
KOH 
Example(11): Give the products of the reaction below, then write the ionic and net ionic equations.
H3PO4
+
Ca(OH)2 
4. Formation of gases
We will assume that the three compounds below will decompose immediately after being made by a
double replacement reaction:
H2CO3  CO2 
+ H2O
H2SO3  SO2 
+ H2O
NH4OH  NH3 
+ H2O
Example(12): Give the products of the reaction below, then write the ionic and net ionic equations.
HCl
+
Na2CO3 
Example(13): Give the products of the reaction below, then write the ionic and net ionic equations.
NaOH
+
(NH4)2SO4 
81
IONIC REACTIONS / Oxidation-Reduction
C. Simple Oxidation-Reduction Reactions (Redox)
1. Single replacement reactions
Single replacement reactions are actually one type of redox reaction.
element + electrolyte
 new element + new electrolyte
Example(1): Give the products for the following reaction:
Al + HCl 
Example(2): Give the products for the following reaction:
F2 + NaCl 
Example(3): Give the products for the following reaction:
Al + CuSO4 
2. Oxidation and reduction
a) Oxidation: the loss of electrons.
b) Reduction: the gain of electrons.
c) Oxidizing agent: the substance that causes something to become oxidized;
the electron taker.
d) Reducing agent: the substance that causes something to become reduced:
the electron donor.
Example(4): a) Write the following equation in net ionic form. b) With an arrow show the transfer of the electron(s).
c) Identify which substance is oxidized, which is reduced, the oxidizing agent and the reducing agent.
Zn + CuSO4  ZnSO4 + Cu
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IONIC REACTIONS / Oxidation-Reduction
Example(5):
Mg + 2HCl  MgCl2 + H2
Example(6):
2Fe + 3SnCl4  2FeCl3 + 3SnCl2
3. Half-reactions
Example(7): Write the half-reactions (including the electrons) for: Zn + CuSO4
 ZnSO4 + Cu
Example(8): Write the half-reactions (including the electrons) for:
 MgCl2 + H2
Mg + 2HCl
Example(9): Write the half-reactions (including the electrons) for: 2Fe + 3SnCl4
 2FeCl3 + 3SnCl2
4. Properties of elements as oxidizing and reducing agents
Compounds can also be oxidizing and reducing agents. You will see this next semester.
a) The higher the EN of an element, the better the oxidizing agent (There are other factors,
but this is the primary correlation.)
Example(10):
F2
Cl2
Br2
I2
Example(11):
H2
O2
F2
83
IONIC REACTIONS / Oxidation-Reduction
Example(12): Will a reaction occur between these reactants?
Cl2 + NaBr

Example(13): Will a reaction occur between these reactants?
I2 + NaBr

b) The lower the IE or EN, the better the reducing agent
(There are other factors, but this is the
primary correlation.)
F-
Example(14):
ClBrI-
Example(15):
Na
Mg
Al
Example(16): Will a reaction occur between these reactants?
Mg +
AlCl3

Example(17): Will a reaction occur between these reactants?
Al
+ MgCl2

c) The better a metal is as a reducing agent, the more “active” it is
84
IONIC REACTIONS / Ion-Electron Method
D. The Ion-Electron Method
The ion-electron method has two goals, 1st it shows what takes place during the reaction, 2nd it balances
the equation.
Example(1): Apply the ion-electron method to the equation below:
Al + HCl  AlCl3 + H2
Step 1: Find and write the half-reactions
This can be done by checking for changes in oxidation states or by writing the equation in net ionic form.
Al  Al3+
H+  H2
Step 2: Balance the atoms in the half-reactions
Al
 Al3+
2H+  H2
Step 3: Balance the charges in the half-reactions by adding electrons
Al  Al3+
+ 3e-
2 e- + 2H+  H2
Step 4: Make the number of electrons lost = the number gained by multiplying one or
both half-reactions by a stoichiometric factor
2(Al  Al3+
3(2 e- + 2H+  H2)
+ 3e-)
85
IONIC REACTIONS / Ion-Electron Method
Step 5: Add the half-reaction to arrive at the balanced net ionic equation
2Al  2Al3+
+ 6e-
6 e- + 6H+  3H2
_____________________________________________
2Al + 6H+  2Al3+ + 3H2
Step 6: Add the spectator ion back in and associate the ion to arrive at the balanced
molecular equation
2Al + 6H+  2Al3+ + 3H2
6Cl6Cl2Al + 6HCl  2AlCl3 + 3H2
Example(2): Apply the ion-electron method to the equation below:
Fe + SnCl4  FeCl3 + SnCl2
Step 1: Find and write the half-reactions
Step 2: Balance the atoms in the half-reactions
Step 3: Balance the charges in the half-reactions by adding electrons
Step 4: Make the number of electrons lost = the number gained by multiplying one or
both half-reactions by a stoichiometric factor
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IONIC REACTIONS / Ion-Electron Method
Step 5: Add the half-reaction to arrive at the balanced net ionic equation
Step 6: Add the spectator ion back in and associate the ion to arrive at the balanced
molecular equation
Example(3): Apply the ion-electron method to the equation below:
F2 + NaCl  NaF + Cl2
Step 1: Write the half-reactions
Step 2: Balance the atoms
Step 3: Balance the charges
Step 4: Make the number of electrons lost = the number gained
Step 5: Add the half-reaction
Step 6: Add the spectator ion back in