Types of Chemical Reactions
Synthesis (combination) reactions – two or more substances combine to form a single substance.
A + B = AB
 Group A metals combine with nonmetals to form ionic compounds. (To get the correct formula you must
know the charges of the cations and anions that the metal and nonmetal form).
2K(s) + Cl2(g)  2KCl(s)

When two nonmetals react in a synthesis reaction, there is more than one possible product. – you must be
given the product name
S(s) + O2(g)  SO2(g) or
2S(s) + 3O2  2SO3(g)

When a transition metal and a nonmetal react in a synthesis reaction, there may be more than one possible
product because the transition metal could form more than one cation. – you must be given at least the
charge on the cation.
Fe(s) + S(s)  FeS(s) Iron (II) sulfide
2Fe(s) + 3S(s)  Fe2S3 Iron (III) sulfide.

Nonmetal oxides react with water to produce an acid (H+ compound).
SO2(g) + H2O(l)  H2SO3(aq) sulfurous acid

Metallic oxides react with water to give a base (OH- compounds). Use ionic charges to write formula of
product.
CaO(s) + H2O(l)  Ca(OH)2(aq)

A metal oxide and nonmetal oxide combine to form a salt.
CO2(g) + Na2O(cr)  Na2CO3(s)
Decomposition Reactions – a single compound is broken down into two or more products and usually require
energy (heat, light or electricity) to take place.

AB
A + B



When a binary (2 elements only) compound breaks down, the products will be those 2 elements.
electricit

y 

H2O(l)
H2(g) + O2(g)
.
When some acids are heated, they decompose to form water and nonmetal oxide

H2CO3(aq)
CO2(g) + H2O(l)



When some metal hydroxides are heated, they decompose to form a metallic oxide and water.

Ca(OH)2
CaO(s) + H2O(g)



When some metallic carbonates are heated, they decompose to form a metallic oxide and carbon dioxide.

Li2CO3(s)
Li2O(s) + CO2(g)



When metallic chlorates are heated, they decompose to form metallic chlorides and oxygen.

KClO3(s)
2KCl(s) + 3O2(g)


Single Replacement/Displacement Reactions – one element replaces a second element in a compound.
A + BC  B + AC (If A is a metal) or
A + BC  C + BA (If A is a nonmetal)

Whether one metal will displace another metal from a compound can be determined by the relative
reactivities of the two metals. An activity series lists the reactivities of some metals.
Decreasing Reactivity
Activity Series of Metals
Name
Symbol
Lithium
Li
Potassium
K
Calcium
Ca
Sodium
Na
Magnesium
Mg
Aluminum
Al
Zinc
Zn
Iron
Fe
Lead
Pb
(Hydrogen)
(H)*
Copper
Cu
Mercury
Hg
Silver
Ag
*Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only.

A nonmetal can also replace another nonmetal from a compound, usually a halogen. The activity of the
halogens decreases as you go down group 17 of the periodic table.
F
Cl
Br
I
At
Combustion Reactions – an element or compound reacts with oxygen, often producing energy as heat and
light.
 Commonly involves hydrocarbons (compounds that only contain H and C)


Complete combustion – forms carbon dioxide and water.
CxHy + O2  CO2 + H2O
Incomplete combustion – reaction runs out of oxygen, then elemental carbon and carbon monoxide
may be additional products.
 Combustion reactions between elements and oxygen also exist.
2Mg(s) + O2(g)  2MgO(s)
Double Replacement/Displacement Reactions – Involve an exchange of two anions between two reacting
ionic compounds.
AB + CD  AD + CB
 For a double- displacement reaction to occur, one of the following is usually true:
1. One product precipitates out of solution.
2. One product is a gas that bubbles out of the mixture.
3. One product is a molecular compound, usually water.
To describe double displacement reactions more clearly we use net ionic equations.
Net Ionic Equations
 Most ionic compounds when dissolved in water dissociate, or separate, into their anions and cations.
Molecular equation:
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)

When AgNO3 is dissolved in water, it separates into Ag1+ cations and NO3- anions. The other aqueous
compounds dissociate also.
Complete Ionic Equation:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)

Ions that appear on both sides of the reaction are not directly involved in the reaction and are called
spectator ions. These ions may be canceled out of the reaction.
Net Ionic Equation:
Ag+(aq) + Cl-(aq)  AgCl(s)
 Note: When writing a balanced net ionic equation, you must balance the charges as well as the atoms.
Precipitation Reactions
The solid that forms after a reaction is called a precipitate.
To decide if the product is a precipitate or not you must follow the solubility rules.
Solubility Rules for Ionic Compounds
Soluble
(dissolve, dissociate, separate)
Nitrate (NO3-) salts
Chlorate (ClO3-) salts
Alkali metal salts (Na+, K+) and ammonium salts (NH4+)
Chloride salts (Cl-) and Br-, ISulfate salts (SO42-)
Except ones with Na, K and Ca
*salt is used to mean ionic compound
Insoluble
Stays a solid in solution
Except ones with Ag, Hg, and Pb
Except ones with Pb, Ag, Hg, Ba, Sr, and
Ca
Sulfides (S2-)
Phosphates (PO43-)
Chromates(CrO42-)
Carbonates (CO32-),
Hydroxides (OH-)
Acid/ Base Reactions/Neutralization Reactions

Acids – a substance that produces H+ ions (protons) when it is dissolved in water. Strong acids completely
dissociate in water. The strong acids are HCl, HNO3, and H2SO4.

Bases/alkalis – a substance that produces hydroxide ions (OH-) in water. Strong bases completely
dissociate in water. The strong bases are NaOH and KOH.

In the reaction of a strong acid and a strong base, one product is always water and the other is always a salt
(ionic compound) that remains dissolved in the water. Therefore, the net ionic equation for all strong
acid/strong base reactions is always:
H+(aq) + OH-(aq)  H2O(l)
Oxidation-Reduction (Redox) Equations – involves the transfer of electrons.
Electrolyte – Carries a current through water. Strong – almost all of the molecules dissociate into ions, carries a
strong current. Weak – some may dissociate, but mostly the molecule stays intact in water, carries very little
current.
Oxidation - Reduction
Synthesis
Decomposition
Combustion
Double
displacement
Single
displacement
Activity series
of metals
Precipitation
Acid/Base
Activity Series
of halogens