Stoichiometry
Chemistry 6.0
The Mathematics of Chemical Reactions:
STOICHIOMETRY
I. Balanced Chemical Equations
A. Provide qualitative and quantitative information
B. Supports the Law of Conservation of Matter
2H2 + O2  2H2O
The above equation is interpreted in terms of particles as
follows:
1. 2 molecules of H2 react with 1 molecule of O2
to produce 2 molecules of water.
The ratio of H2 to O2 is 2:1.
or 2 moles of H2 react with 1 mole of O2 to produce
2 moles of water.
The ratio of H2 to O2 is 2:1.
2. It is more convenient to interpret the coefficients as number of
moles, because we measure amounts of substances by massing.
C. Stoichiometry
1. The study of the quantitative relationships
that exist in a formula or a chemical reaction.
2. Importance
a. Provides for the safe, economical and reproducible
manufacture of chemicals.
b. Provides for the safe administration of pharmaceuticals.
D. Proof of the conservation of matter in a balanced equation
1. Convert all reactants and products to their mass equivalents.
2. Sum up the mass of reactants and compare the sum of the
mass of products.
II. Stoichiometry Problems
A. Steps to Solve Problems
1. Write a balanced equation.
2. Identify the given ( ) and the
unknown or required substance (?).
3. Convert mass of given into moles.
4. Use the mole (molar) ratio to convert
from given to required substance.
5. If needed, convert moles of required
into mass of required substance.
B. Examples
1. How many moles of oxygen are required to react with
16 moles of hydrogen in the production of water?

?
2 H2 + O2  2 H2O
16 moles H2 1 mole O2
2 moles H2
=
8.0 moles O2
Mole ratio links 1 substance to another in a reaction.
Required in problem solving.
2. Antimony reacts with water to produce antimony(III) oxide
and hydrogen. How many moles of hydrogen are produced
from 7.5 moles of antimony?
2Sb + 3H2O  Sb2O3 + 3 H2
11 mol H2
3. What mass of aluminum oxide can be prepared by the
reaction of 67.5 g of aluminum in a synthesis reaction?
4 Al + 3O2  2Al2O3
128 mol Al2O3
4. Sodium bicarbonate, a.k.a. baking soda, can be used to
extinguish a fire. When heated, it decomposes to give
carbon dioxide gas which smothers the fire. It also
produces sodium carbonate and water. If a sample contains
4.0 g of sodium bicarbonate, what mass of carbon dioxide
is produced?
2 NaHCO3  Na2CO3 + H2O + CO2
III. Percent Yield
A. Expected Yield: the amount of product that
should be produced (theoretical)
B. Actual Yield: the amount of product that is
actually produced (experimental)
C. Percent Yield: percent of expected yield
that was obtained
% Yield = (actual yield/expected yield) x 100
D. Steps to Solving Percent Yield
Problems
1. Write a balanced equation
2. Identify the given () which is the mass of
reactant, and identify the actual yield.
3. Solve for the expected mass of product using
the given mass of reactant.
4. Calculate the % yield.
% Yield = actual yield x 100
expected yield
E. Examples
1. A reaction between 2.80 g aluminum nitrate and
excess sodium hydroxide produced 0.966 g of
aluminum hydroxide in this double replacement
reaction. Calculate the % yield.
Al(NO3)3 + 3NaOH  Al(OH)3 + 3NaNO3
1.03 g Al(OH)3
2. Determine the % yield for the reaction between 3.74 g of
sodium and excess oxygen if 4.24 g of sodium oxide is
recovered in the direct combination reaction.
4Na + O2  2Na2O
IV. Limiting Reactants
A. Definition: the reactant that determines, or limits,
the amount of product(s) formed in a chemical
reaction
B. Problem Solving Tips
1. The limiting reactant is not necessarily the reactant
present in the smallest amount
2. When you are given the amounts of 2 or more
reactants, you should suspect that you are dealing with
a limiting reactant problem.
C. Steps
1. Write a balanced equation
2. Calculate the number of moles of each
reactant
3. Compare the mole ratios of the reactants as
available ratio (from the given masses) and
the required ratio (from the coefficients)
4. Identify the limiting reactant, and use it to
calculate the mass of product formed.
D. Examples
1. What mass of CO2 could be formed by the
combustion of 16.0 g CH4 with 48.0 g O2?
CH4 + 2O2  CO2 + 2H2O
33.0 g CO2
Oxygen is the limiting reactant
2. What is the maximum mass of nickel(II) hydroxide that
could be prepared by mixing 25.9 g nickel(II) chloride
with 10.0 g sodium hydroxide?
NiCl2 + 2NaOH  Ni(OH)2 + 2NaCl
11.6 g Ni(OH)2
Sodium Hydroxide is the LR
V. Solution Stoichiometry
A. Many reactants are introduced to a reaction chamber
as a solution.
B. The most common solution concentration is molarity.
molarity =
mol/liter
C. Examples
1. Excess lead(II) carbonate reacts with 27.5 mL of 3.00M nitric
acid. Calculate the mass of lead(II) nitrate formed
PbCO3 + 2HNO3  Pb(NO3)2 + H2CO3
2. Calculate the volume, in mL, of a 0.324 molar solution of
sulfuric acid required to react completely with 2.792 g of
sodium carbonate according to the equation below.
H2SO4 + Na2CO3  Na2SO4 + CO2 + H2O
VI: Acid-Base Titration
1. An acid-base titration is a carefully controlled
neutralization reaction or redox which can
determine concentration of an unknown solution.
2. To determine the concentration of an unknown
substance, a standard solution is needed. This
solution has a known concentration.
3. Titration curve: graph that
shows how pH changes
during a titration.
4. An indicator, usually
phenolphthalein, is
used in a titration.
•
Colorless in an acid,
pink in a base.
5. The point at which
enough standard
solution is added to
neutralize the unknown
solution is called the
equivalence point.
6. The point at which the indicator
changes color is called the endpoint.
7. Therefore: [H+] = [OH-] at the
equivalence point
22
Ex) Solutions of sodium hydroxide are used to unclog
drains. A 43.0 mL volume of sodium hydroxide was
titrated with 32.0 mL of 0.100 M HCl. What is the
molarity of the sodium hydroxide solution?
HCl + NaOH  NaCl + HOH
Note: Mole ratio of the acid and base is 1:1
HCl (Acid) =
NaOH (Base)
0.0744 M = Mb
Ex) A volume of 25.0 mL of 0.120M sulfuric acid
neutralizes 40.0mL of a sodium hydroxide solution.
What is the concentration of the sodium hydroxide
solution?
H2SO4 + 2 NaOH 
Na2SO4 + 2 HOH
Note: Mole ratio between acid an base in not 1:1
Mb = 0.150M
VII. Stoichiometry and Heat
Changes
1. How much heat is released when 22.0g of
propane is burned?
C3H8 + 5O2 → 3CO2 + 4H2O
ΔH = -2.22x103 kJ
-1.11x103 kJ (released)
2. How much carbon dioxide is produced, in
grams, when 2,500 kJ of energy is released?
150 g CO2
VIII. Determining ∆H Using Heats of
Formation
The standard enthalpy change of a
reaction is equal to the sum of the
standard molar enthalpies of formation
of the products multiplied by its
coefficient, n, in the balanced equation,
minus the corresponding sum of
standard molar enthalpies of formation
of reactants.
Hrxn = ∑ nHf,products - ∑ nHf,reactants
ΔH from ΔHf Problem:
Using the Heats of Formation Table,
calculate the H for the following
reaction:
SF6(g) + 3H2O(l)  6HF(g) + SO3(g)
45.1 kJ
Write the thermochemical equation for
this reaction:
SF6 + 3H2O + 45.1kJ  6HF + SO3
ΔH from ΔHf Problem:
Using the Heats of Formation Table, calculate
the standard heat of combustion for
propane.
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
-2043.9 kJ
Write the thermochemical equation for this
reaction:
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g) + 2043.9kJ