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Types of Chemical
Reactions
Classifying Chemical
Changes
 The products of a chemical reaction may often be
predicted by applying known facts about common
reaction types
 Five general types of reactions:
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Single displacement
Double displacement
Decomposition
Synthesis
Combustion
Single Displacement
 (aka single replacement)
 One element displaces another element in a compound
 Metal replaces metal (+)
 Nonmetal replaces nonmetal (-)
 General form:
 element + compound  element + compound
A + BC  B + AC
 General types of single displacement reactions:
 An active metal will replace the metallic ion in a compound of a less
active metal.
 Ex: Fe + Cu(NO3)2  Fe(NO3)2 + Cu
 Some active metals such as sodium and calcium will react with
water to give a metallic hydroxide and hydrogen gas.
 Ex: Ca + 2H2O  Ca(OH)2 + H2
 Active metals such as zinc, iron, and aluminum will displace the
hydrogen in acids to give a salt and hydrogen gas.
 Ex: Zn + 2HCl  ZnCl2 + H2
 An active nonmetal will displace a less active nonmetal.
 Ex: Cl2 + 2NaBr  2NaCl + Br2
Double Displacement
 (aka double replacement)
 Ions in two compounds “change partners”
 Cation of one compound combines with the anion of the
other compound
 General form:
AB + CD  AD + CB
 compound + compound  compound + compound
 General types of double displacement reactions:
 A reaction between an acid and a base yields a salt and water.
Such a reaction is a neutralization reaction.
 Ex: 2KOH + H2SO4  K2SO4 + 2H2O
 Reaction of a salt with an acid forms a salt of the acid and a
second acid that is volatile.
 Ex: 2KNO3+ H2SO4  K2SO4 + 3HNO4
 Reactions of some soluble salts produce an insoluble salt and a
soluble salt.
 Ex: AgNO3 + NaCl  AgCl + NaNO3
Decomposition
 A compound breaks down into 2 or more simpler
substances
 Only one reactant
 General form:
 compound  two or more substances
AB  A + B
 General types of decomposition reactions:
 When some acids are heated, they decompose to form water and an acidic
oxide.
 Ex: H2CO3  CO2 + H2O
 When some metallic hydroxides are heated, they decompose to form metallic
oxide and water.
 Ex: Ca(OH)2  CaO + H2O
 When some metallic carbonates are heated, they decompose to form a
metallic oxide and carbon dioxide.
 Ex: Li2CO3  Li2O + CO2
 Most metallic oxides are stable, but a few decompose when heated.
 Ex: 2HgO  2Hg + O2
 Some compounds cannot be decomposed by heat, but can be decomposed
into their elements by electricity.
 Ex: 2NaCl  2Na+ Cl2
Synthesis
 The combination of 2 or more substances to form a compound
 Only one product
 General form:
 element or compound + element or compound 
compound
A + B  AB
 General types of synthesis reactions:
 Two or more elements combine to form a compound.
 Ex: Fe + S  FeS
 An acid anhydride, nonmetallic oxide, combines with water to give
an acid.
 Ex: SO2 + H2O  H2SO3
 A basic anhydride, metallic oxide, combines with water to form a
base.
 Ex: Na2O + H2O  2NaOH
 A basic oxide combines with a nonmetallic oxide to form a salt.
 Ex: CO2 + Na2O  Na2CO3
Combustion
 Contains oxygen (O2)
 Usually contains hydrocarbons (H and C molecules)
combine with oxygen to produce a large amount of
energy, which produce CO2 + H2O
 General form:
 Hydrocarbon + oxygen  carbon dioxide + water
CxHy + O2 CO2 + H2O
 Ex: The burning of propane
C3H8 + 5O2  3CO2 + 4H2O
You Try!
 Balance and classify the following chemical reactions:
1 Na3PO4 + __
3 KOH  ___
3 NaOH + __
1 K3PO4
1. __
double displacement
1 P4 + __
3 O2  __
2 P2O3
2. __
synthesis
2 NO  __
2 O +1__ N
3. __
2
2
2
decomposition
2 AgNO3 + 1
1 Cu(NO3)2 + __
2 Ag
4. __
__ Cu  __
single displacement
1 C3H6O + __
4 3
3 2 + __ H2O
5. __
O2  __ CO
combustion