Vii. Ionic Reactions in Aqueous Solutions

IONIC REACTIONS / Electrolytes

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.

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.

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:

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:

AgNO3 + NaCl AgCl + NaNO3

Example(3): Write the ionic and net ionic equations for:

2Na3PO4 + 3CaCl2 Ca3(PO4)2 + 6NaCl

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

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

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

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: Mg + 2HCl MgCl2 + H2

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

Example(11): H2 O2 F2

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.)

Example(14): F-

Cl-

Br-

I-

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

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+ + 3e-)

3(2 e- + 2H+ H2)

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

6Cl- 6Cl-

2Al + 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

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