Let's first start with a complete chemical equation and see how the net ionic equation is derived. Take for example the reaction of lead(II) nitrate with sodium chloride to form lead(II) chloride and sodium nitrate, shown below:

Pb(NO3)2(aq) + 2 NaCl(aq) --> PbCl2(s) + 2 NaNO3(aq)

This complete equation may be rewritten in ionic form by using the solubility rules. Lead(II) nitrate is soluble and therefore dissociated. The same for NaCl. As products, sodium nitrate is predicted to be soluble and will be dissociated. The lead(II) chloride, however, is insoluble. The above equation written in dissociated form is:

Pb2+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> PbCl2(s) + Na+(aq) + NO3-(aq)

At this point, one may cancel out those ions which have not participated in the reaction. Notice how the nitrate ions and sodium ions remain unchanged on both sides of the reaction.

Pb2+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) --> PbCl2(s) + Na+(aq) + NO3-(aq)

What remains is the net ionic equation, showing only those chemical species participating in a chemical process:

Pb2+(aq) + 2 Cl-(aq) --> PbCl2(s)

Diatomic Molecules must always be written as 2 atoms of the element:

#### H2, Cl2, F2, I2, N2, Br2, O2

There are five basic types of reactions:
1. Decomposition
2. Synthesis
3. Single Replacement
4. Double Replacement
5. Combustion

• Remember, a correctly written equation must have all the correct formulas and it must be balanced.

• Always make sure that you have the oxidation numbers of the elements within a compound balanced to zero.

• If an element has more than one oxidation number, the state in which it exsts as a reactant is the same state it will exist as a reactant unless otherwise stated.

• Reactants ----> Products

• #### 1. Decomposition Reaction

• When energy in the form of heat, electricity, light, or mechanical shock is supplied, a compound may decompose to form simpler substances.
compound ---> 2 or more substances
• Rule 1: Acids decompose to form nonmetallic oxides and water
H2CO3 ---> CO2 + H2O

• Rule 2: Metallic hydroxides decompose to form metallic oxides and water
Ba(OH)2 ---> BaO + H2O

• Rule 3: Metallic carbonates decompose to form metallic oxides and carbon dioxide
MgCO3 ---> MgO + CO2

• Rule 4: Metallic chlorates decompose to form metallic chlorides and oxygen gas (do not forget, oxygen is a diatomic molecule)
NaClO3 ---> NaCl + O2

• Rule 5: Metallic oxides decompose to form the metal and oxygen gas
MgO ---> Mg + O2

• Rule 6: Some compounds just decompose to for their individual elements.
NaCl ---> Na + Cl2

##### Sample Problems
NaOH --->
Fe(OH)3 --->
CaCO3 --->

#### 2. Synthesis Reaction

• In a synthesis reaction, 2 or more simple substances are combined to form one new and more complex substance. (The reverse of a decomposition reaction)

element/compound + element/compound ---> new compound

• Rule 1: 2 single elements form 1 substance
Na + Cl2 ---> NaCl

• Rule 2: Nonmetallic oxides combined with water to form an acid with one more oxygen than the nonmetallic oxide
SO2 + H2O ---> H2SO3

• Rule 3: Metallic hydroxides combine with water to form hydroxides
BaO + H2O ---> Ba(OH)2

##### Sample Problems
Mg + O2 --->
NO2 + H2O --->
Fe + O2 ---> (iron (III) is formed)

#### 3. Single Replacement Reaction

• One element displaces another element in a compound
single element + compound ---> new single element + new compound
• Rule 1: Replacement of a metal in a compound by a more active metal (see chart on page 159 of the textbook)
Na + Zn(NO3)2 ---> NaNO3 + Zn

Al + Pb(NO3)2 ---> Al(NO3)3 + Pb

• Rule 2: Replacement of hydrogen in water by metals resulting in hydrogen gas as a product.
Al + H2SO4 ---> ZnSO4 + H2

• Rule 3: An active nonmetal will displace a less active nonmetal (a nonmetal that is lower on the on the periodic table)
Cl2 + NaBr ---> NaCl + Br2
##### Sample Problems
Na + BaCl2 --->
Mg + HCl --->

#### 4. Double Replacement Reaction

• The positive and negative ions of 2 compounds are interchanged in a double replacement reaction
compound + compound ---> new compound + new compound

KOH + H2SO4 ---> K2SO4 + H2O
• (watch for the combination of hydrogen and the hydroxide ion--it will form water)
AgNO3 + NaCl ---> AgCl + NaNO3
Ca(OH)2 + HCl ---> CaCl2 + H2O
##### Sample Problems
NH4S + Fe(NO3)2 --->
AgNO3 + KCl --->
Mg(OH)2 + H3PO4 --->

#### 5. Combustion Reaction

• The reaction of oxygen with a compound that contains both hydrogen and carbon (hydrocarbon)
• The products of this type of reaction depends on the amount of oxygen present as a reactant.
• Limited supply of oxygen as a reactant with a hydrocarbon will produce carbon monoxide and water.
• Excess supply of oxygen as a reactant with a hydrocarbon will produce carbon dioxide and water.
CH2 + O2(exc) ---> CO2 + H2O
C6H6 + O2(lim) ---> CO + H2O

• As always, there are exceptions to the rules in chemistry.
• When H2CO3 or NH4OH are products that are formed in a chemical reaction, they will continue to decompose.
• Carbonic acid, H2CO3, will continue to decompose to H2O and CO2.
(NH4)2CO3 + HCl --> NH4Cl + H2CO3

carbonic acid will continue to decompose to form:

(NH4)2CO3 + HCl ---> NH4Cl + H2O + CO2
• Ammonium hydroxide will continue to decompose into ammonia and water
(NH4)2CO3 + NaOH ---> Na2CO3 + NH4OH

ammonium hydroxide will continue to decompose:

(NH4)2CO3 + NaOH ---> Na2CO3 + NH3 + H2O

Extra Problems

Write the correct formulas, equation, and balance the equation:

• Strontium carbonate decomposes into:
• Lithium hydroxide decomposes into:
• Potassium chlorate decomposes into:
• Lead (II) oxide decomposes into:
• Fluorine gas reacts with potassium iodide to form:
• Ammonium nitrate reacts with sodium hydroxide to form: