Chapter 11
Chemical Reactions
Chemical Equation
Describes chemical reaction.
 Chemical equation: reactants yield products
 Reactants  Products


Much easier to write symbols and formulas
instead of words
Examples



Solid Iron reacts with oxygen gas to form the solid
IronIIIoxide.
iron(s) + oxygen(g)  ironIIIoxide(s)
Fe(s) + O2(g)  Fe2O3(s)

Carbon tetrahydride gas BURNS to form carbon dioxide
gas and water vapor.

Carbon tetrahydride(g) + oxygen(g)  carbon dioxide(↑) + water(↑)



CH4(g) + O2(g)  CO2(↑) + H2O(↑)
Skeleton Equation: chemical equation that tells you what the
reactants and products are but NOT how much of each you have.
First step in writing a chemical equation.
Symbols Used






(s)
(l)
(g)
(↑)
(aq)
()

D
Pt



solid
liquid
gas
gas as a product
aqueous (in water solution)
ppt (precipitate) solid product from 2 aqueous
reactants
means with heat
means with Platinum catalyst: speeds up a
reaction without being used.
reversible reaction
Balancing Chemical Equations
Balanced equations have:
 the same # of atoms of each element on
BOTH sides of the equation.


Law of Conservation of Mass – atoms can
neither be created nor destroyed, simply
rearranged.
Rules for Balancing Equations

Get the correct formulas for reactants and
products.

(USE ION CHART AND DON”T FORGET
DIATOMIC ELEMENTS!)

Write reactants on left, products on right.
Use plus signs to separate compounds and
yield sign to separate the reactants from
products.
Rules Continued
Count the # of atoms of each element in
reactants and products. (Polyatomic atoms
on both sides count as one.)
 Balance # of each element using
coefficients.
 Coefficient – small whole # in front of a
formula.
 NEVER CHANGE FORMULA
SUBSCRIPTS

Rules for Balancing Equations
Balance elements appearing 3 or more
places LAST.
 Check each element to make sure equation
is balanced.
 Make sure all coefficients are in the lowest
whole number ratio.

Do not change subscripts!!!
Diatomic Molecules
Diatomic Molecules- a molecule made up two
atoms of the same element. They are only
diatomic when they are alone.
-There are 7 naturally occurring
diatomic molecules.
H O N Cl
Br
I F
Balancing Examples

___ C(s) + ___ O2(g)  ___ CO2 (g)

___ C(s) + ___ O2(g)  ___ CO (g)

___ AgNO3(aq) + ___Cu(s)  ___ Cu(NO3)2(aq) + ___ Ag(s)

___ Al(s) + ___ O3(g)  ___ Al2O3(s)

*___ C2H6(g) + ___ O2(g)  ___ CO2(g) + ___ H2O(g)

*___ H3PO3  ___ H3PO4 + PH3
5 Types of Chemical Reactions

Combination Reaction – elements combine to
form a compound.
A

+B
AB
element + element  compound
Ex. Sodium + chlorine  sodium chloride
2
2 NaCl(s)
 ___Na(s)
+ ___ Cl2(g)  ___

5 Types of Reactions

Decomposition Reaction – compound
breaks down into its element.
D
A+B
compound D element + element
 AB

Ex: MercuryII oxide  mercury + oxygen
2 HgO  ___Hg
2
 ___
+ ___O2

5 Types of Reactions - 3
Single Replacement Reaction – one element
replaces another element in a compound.
+ -
+
+
AB + C
+ -
-
AB + D
+ -
A + CB
or
+ -
-
AD + B
Examples of Single Replacement
Reactions
• Must use Activity Series to see if reaction works
• Zinc + sulfuric acid  zinc sulfate + hydorgen
• Zn(s) + H2SO4(aq)  ZnSO4(aq) + H2(↑)
• Periodic table is activity series for halogens
•
Sodium bromide + chlorine  sodium chloride + bromine
•
2
2
___NaBr(s)
+ ___Cl2(g)  ___NaCl(s)
+ ___Br2(↑)
5 Types of Reactions
Double Replacement Reaction – two compounds react and
exchange positive ions to form two new compounds.
+ + + + -
AB + CD
AD + CB

Barium Chloride(aq) + potassium carbonate(aq) 
barium carbonate() + potassium chloride(aq)

2 KCl(aq)
BaCl2(aq) + K2CO3(aq)  BaCO3() + ___
5 Types of Reactions
Combustion Reaction – oxygen reacts with a
compound composed of C and H.
C x Hy + O 2
CO2 + H20
Also called Burning (exothermic)
The products are always CO2 and H2O.
Examples of Combustion Reactions
1.
C6H6 + 7½ O2
2 C6H6 + 15 O2
2.
CH3OH + 1½ O2
2 CH3OH + 3 O2
6 CO2 + 3 H2O
12 CO2 + 6 H2O
CO2 + 2 H2O
2 CO2 + 4 H2O
Special Decomposition Reactions:
 Decomposition

of a Carbonate:
Metal carbonate  metal oxide + carbon dioxide
XCO3
ex. Na2CO3
XO + CO2
Na2O + CO2
Special Decomposition Reactions:
 Decomposition

of a Hydroxide:
Metal hydroxide  metal oxide + water
XOH
ex. 2NaOH
XO + H2O
Na2O + H2O
Special Decomposition Reactions:
 Decomposition


ex.
of a Chlorate: (ClO3)
Metal chlorate  metal chloride + oxygen
XClO3
2
___NaClO
3
XCl + O2
___NaCl + ___O
3 2
Special Decomposition Reactions: 4
 Special
single Replacement Reaction:
»Group IA or IIA metal and H2O
X + HOH
ex. 2Na + 2HOH
XOH + H2
2NaOH + H2
How to ID types of reactions.
Combination Reactions – given 2 items that form 1
new compound.
Decomposition Reactions – given a single
compound that breaks into parts.
Single Replacement – given a single element plus a
single compound, forms a new compound a a
different element.
Double Replacement – given two compounds (+’s
change places).
Combustion Reaction – given CH compound with
Oxygen, always forms water and carbon dioxide.