Stoichiometry
Calculations with Chemical Formulas and Equations
Chapter 3 BLB 12th
Expectations
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Balance chemical equations.
g ↔ moles ↔ molecules ↔ atoms
Find empirical and molecular problems.
Calculate amounts of reactants and products.
Calculate theoretical and percent yield.
Stoichiometry
Quantity relationships based on chemical
equations
3 Main Concepts:
1. Chemical formula – molar ratio of atoms
2. Chemical equations – molar ratio of
compounds
3. Law of Conservation of Mass:
mass of reactants = mass of products
3.1 Chemical Equations
Components:
 reactants → products
 Physical states (s, l, g, aq)
 Reaction conditions (heat Δ, light, solvents, etc.)
 Coefficients determine molar ratios. The number
of moles of each type of atom must be the same
on each side.
Balancing:
 By inspection
 Use coefficients; don’t change chemical formulas
Coefficients vs. Subscripts
Fe2S3(s)
+
HCl(aq)
KClO3(s)
HNO3(l)
+
→
P4O10(s)
→
FeCl3(s)
KCl(s)
→
+
+
H2S(g)
O2(g)
(HPO3)3(l)
+
N2O5(g)
3.2 Some Simple Patterns of
Chemical Reactivity
Combination and Decomposition
combination: A + B → C
4 Fe(s) + 3 O2(g) → 2 Fe2O3(s)
decomposition: C → A + B
2 NaN3(s) → 2 Na(s) + 3 N2(g)
3.2 Some Simple Patterns of
Chemical Reactivity
Combustion
 burning of a fuel in the presence of oxygen
 products of complete combustion: CO2, H2O
 exothermic (produces heat)
Molecular view of methane combustion
3.2 Some Simple Patterns of
Chemical Reactivity
Combustion
C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g)
2 CH3OH(g) + 3 O2(g) → 2 CO2(g) + 4 H2O(g)
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Each C atom in fuel produces 1 mol CO2
Each H atom in fuel produces ½ mol H2O
3.3 Formula Weights
Formula and Molecular Weights (amu)
formula weight – general
molecular weight – molecules
formula unit weight – ionic compound
- sum of the atomic masses of each atom
in chemical formula
% Composition
part
percent 
100%
whole
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% composition by mass
Mass of one type of atoms over mass of all atoms
3.4 Avogadro’s Number and the Mole
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Word association:
pair –
dozen –
case –
ream –
3.4 Avogadro’s Number and the Mole
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amu impractical for lab use (too small)
Avogadro’s number: 6.0221421 x 1023 mol-1
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The number of atoms in exactly 12 g of 12C
For conversions: 6.022 x 1023 ?/mol, where ?
can equal atoms, molecules, ions, etc.
1 mole = Avogadro’s number of anything
molar mass – mass in grams of one mole of
a substance, which is equal to the atomic
mass in amu; g/mol
3.4 Avogadro’s Number and the Mole

Atoms & compounds have different masses,
thus the mass of 1 mole of atoms &
compounds are different.
3.4 Avogadro’s Number and the Mole
Conversions:
 g → mol divide by molar mass
 mol → g multiply by molar mass
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Abbreviations:
mole – mol
molarity – M
Practice with Avogadro’s # & the Mole
3.5 Empirical Formulas from Analyses
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Empirical formula – smallest whole number
ratio of atoms
Molecular formula – actual ratio of atoms in
a compound; multiple of the empirical
formula; must know molar mass of
compound
Use % composition to find formula
Problems
A once-used gasoline additive contains 49.5% C,
3.2% H, 22.0% O, and 25.2% Mn. Determine
the emipirical formula of this compound.
Azulene, a hydrocarbon, contains 93.71% C. Its
molar mass is ~128 g/mol. Determine the
emipirical and molecular formulas for azulene.
3.5 Empirical Formulas from Analyses
Summary:
 % data → grams
 Grams → moles
 Moles → molar ratio → empirical formula
 Empirical formula → molecular formula
3.5 Empirical Formulas from Analyses
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Combustion analysis
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1 mol C in fuel → 1 mol CO2
2 mol H in fuel → 1 mol H2O
Problems
The combustion of propane, a hydrocarbon,
produces 2.641 g CO2 and 1.442 g H2O.
Determine the emipirical formula of propane.
3.6 Quantitative Information from
Balanced Equations
3 Main Concepts:
1. Chemical formula – molar ratio of atoms
2. Chemical equations – molar ratio of
compounds
3. Law of Conservation of Mass:
mass of reactants = mass of products
What’s balanced in a balanced equation?
Stoichiometry Problems
Use these 4 steps as a guide: (p. 97)
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Write & balance chemical equation.
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Convert to moles.
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Apply molar ratio.
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Convert from moles to quantity desired
(mass, volume, etc.)
Stoichiometry Problems
How many grams of CaCl2 is produced from taking 2
antacid tablets, each containing 500. mg of CaCO3?
CaCO3(s) + 2 HCl(aq) → CO2(g) + H2O(l) + CaCl2(s)
From Sample Exercise 3.16, p. 98
How many grams of HCl are needed to react with
1000 mg of CaCO3?
CaCO3(s) + 2 HCl(aq) → CO2(g) + H2O(l) + CaCl2(s)
3.7 Limiting Reactants
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Limiting reactant – reactant that is completely
consumed; limits the amount of product that can
be formed
Theoretical yield – calculated yield of a product
based on limiting reactant
Percent yield
experimental yield
% yield 
100%
theoretical yield
78. Calculate the theoretical yield (in grams) of
NO when 2.00 g of NH3 react with 2.50 g of O2.
NH3(g) + O2(g) → NO(g) + H2O(g)
Silver metal reacts with elemental sulfur according to the
reaction below. If 2.0 g each of silver and sulfur react,
what is the theoretical yield (in grams) of silver(I)
sulfide? How many grams are left over?
16 Ag(s) + S8(s) → 8 Ag2S(s)
How many grams are left over?
16 Ag(s) + S8(s) → 8 Ag2S(s)
84. When hydrogen sulfide gas is bubbled into a solution of
sodium hydroxide, the reaction forms sodium sulfide and water.
How many grams of sodium sulfide are formed if 1.25 g of
hydrogen sulfide is bubbled into a solution containing 2.00 g of
sodium hydroxide, assuming that the sodium sulfide is made in
92.0% yield?
When 0.750 g iron(III) chloride hydrate is heated,
0.300 g of steam is produced. What is the value of x ?
Δ FeCl (s) + x H O(g)
FeCl3·x H2O(s) →
3
2