Unit 4: Stoichiometry –
Calculations with Chemical Formulas and
Equations
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Reactions and Chemical Equations
Atomic and Molecular Weights
Moles
Stoichiometric Calculations
Limiting Reactants
Percent Composition and Empirical Formulas
Chemical Reactions
 When chemicals are mixed together, one of
two things can happen:
 Mixture forms
chemicals retain their own physical and
chemical properties
individual components can be separated
 Chemical reaction occurs
Chemical Reactions
 Chemical reaction:
 a process in which chemicals are mixed
under certain conditions and transformed
into new chemicals
atoms of the chemicals rearrange and
combine to form new substances
new substances formed often have
different physical and chemical properties
2 Na (s) + Cl2 (g)
2 NaCl (s)
Types of Chemical Reactions
 Combustion Reaction
 a rapid chemical reaction in which heat is
evolved
 Requires oxygen as a co-reactant
 Usually produces a flame
 Burning charcoal:
C (s) + O2 (g)
 Burning propane:
 CO2 (g)
C3H8 (g) + 5 O2 (g)  3CO2 (g) + 4 H2O(l)
Types of Chemical Reactions
 Combination Reactions
 two or more substances react to form one
product
A + B  C
 Producing ammonia:
N2 (g) + 3 H2 (g)  2 NH3 (g)
 Acid rain:
SO3 (g) + H2O (l)  H2SO4 (aq)
Types of Chemical Reactions
 Decomposition Reactions
 chemical reactions in which one substance
reacts to produce two or more other
substances
C  A + B
 Air bags:
2 NaN3 (s)  2 Na (s) + 3 N2 (g)
Signs of Chemical Reactions
 Visible Signs
 formation of a solid, liquid or gas
 color change
Signs of Chemical Reactions
 Other indications of a chemical reaction:
 Energy is released as heat
or light
 Energy is absorbed from
the environment
 Chemical analysis
reveals changes in molecular formula,
chemical properties, etc.
Chemical Equations
 Chemical equation: a representation of a
chemical reaction that uses the chemical
formulas for the compounds used and formed
during the reaction
Reactants
Products
General format for all chemical equations
Chemical Equations
 Reactants
 the starting substances in a chemical
reaction
 Products
 the new chemicals that are formed as a
result of a chemical reaction
CH4 (g) + 2 O2 (g)
reactants
CO2 (g) + 2 H2O (g)
products
Chemical Equations
 Chemical equations provide information about:
 formulas for reactants and products
 relative number of molecules (or moles) of
each reactant or product
 the physical state of each reactant or
product
(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)
2 N aH C O 3 (s)
+
H 2 S O 4 (aq)
N a 2 S O 4 (aq) + 2 H 2 O (l) + 2 C O 2 (g)
Chemical Equations
 Aqueous reactions (aq)
 water as a solvent
dissolved in water to do the reaction
Water does NOT participate in the
reaction
 solvent only
 can be recovered unchanged
AgNO3 (aq) + NaCl (aq)
AgCl (s) + NaNO3 (aq)
Chemical Reactions
 Chemical reactions occur on a molecular or
atomic level.
H2
+
O2

H 2O
H H
O
O
What happened to the
other oxygen atom???
Law of Conservation of Matter
 Atoms cannot be created or destroyed in an
ordinary chemical reaction.
H H
+
O
O
H H
2 H2 + O2 
The reaction to form water
requires 2 molecules of H2
for every molecule of O2.
2 H2 O
Balanced Chemical Equations
 Chemical equations must be balanced:
 They must follow the Law of Conservation of
Matter (Mass).
 They must account for every atom involved
in a chemical reaction.
Balancing Chemical Equations
 To balance a chemical equation:
 Write the correct formulas for reactants
and products.
H2
+
O2

H 2O
 Count the number of atoms of each type on
both sides of the equation.
H2
+
O2

2 H
2 O
H 2O
2 H
1 O
Balancing Equations
 If the number of any element differs from
side to side, add coefficients to make both
sides equal.
H2
+
O2

H 2O
2 H
2 O
H2
2 H
1 O
+ O2
2 H2 + O2

2 H2O
 2 H2O
Rules for Balancing Equations
 NEVER change a formula or subscript in a
formula.
H2 +
O2

H 2O
H2 +
O2

H 2O 2
2H2 +
O2
 2 H2O
Rules for Balancing Equations
 Don’t insert coefficients inside a formula.
H2 +
O2

H 2O
H2 +
O2

H22O
2 H2 + O2
 2 H 2O
Rules for Balancing Equations
 Don’t add extra formulas to the equation.
H2 +
O2

H 2O
H2 +
O2

H 2O + O
2 H2 + O2
 2 H2O
Suggestions for Balancing Equations
 Work with one element at a time.
 Start with an element that is present in
only one compound on each side if
possible.
 Balancing one element may unbalance
another element.
 You many need to change the
coefficients again!!!
 Balance diatomic elements last.
Balancing Equations
Examples: Balance the following equations:
 Ca + N2  Ca3N2
 Ba + O2  BaO
 P + H2  PH3
 Fe + O2  Fe2O3
 CaCO3 + C  CaC2 + CO2
 Ca + NH3  CaH2 + Ca3N2
 Ba(OH)2 + Al(NO3)3  Ba(NO3)2 + Al
(OH)3