Chemistry 6.0
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
2
H2 + O2  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
Problem Solving Tips
B.
1.
2.
The limiting reactant is not necessarily the reactant
present in the smallest amount
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.
2.
3.
4.
Write a balanced equation
Calculate the number of moles of each reactant
Compare the mole ratios of the reactants as available
ratio (from the given masses) and the required ratio
(from the coefficients)
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. Sequential Reactions
A. Definition: A chemical process in which several
reactions are required to convert starting
materials into product(s).
B. Main Concept: The amount of desired product
from each reaction is taken as the starting
material for the next reaction. Consideration is
given to steps that produce less than 100%.
C. Examples
1. Given the following sequence of equations, calculate
the mass of Ni(CO)4 produced from 75.0 g of
carbon. Assume 100% yields.
C + H2O
Ni + 4CO


CO + H2
Ni(CO)4
2. Hydrogen, obtained by the electrical decomposition
of water, is combined with chlorine to produce 84.2 g
of hydrogen chloride. Calculate the mass of water
decomposed. Assume 100% yield.
2H2O
H2 + Cl2


2H2 + O2
2 HCl
VI. 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
Problems
C3H8 + 5O2 → 3CO2 + 4H2O
1.
ΔH = -2.22x103 kJ
How much heat is released when 22.0g of propane
is burned?
-1.11x103 kJ (released)
2.
How much carbon dioxide is produced, in grams,
when 2,500 kJ of energy is released?
150 g CO2
ΔH from ΔHf
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
***By definition, the standard enthalpy of formation of
the most stable form of any element is zero because
there is no formation reaction needed when the
element is already in its standard state.***
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