Stoichiometry
Outline
1. Molar Calculations
2. Limiting Reactants
3. Empirical and Molecular Formula Calculations
Review
1. Molar Calculations
moles (n) =
% Composition =
mass (g)
PV
=
= 𝑀V
−1
Molar Mass (g mol ) RT
mass A
Molar Mass A
x 100% =
= 100%
total mass
Total Molar Mass
% Yield =
Actual
x 100%
Theoretical
6.02 x 1023 particles (atoms or molecules) / 1 mole
2. Limiting Reactants

Find moles of one product – the least amount shows the limiting reactant. The other reactant is in excess.
3. Empirical and Molecular Formula Calculations
# % Composition or # grams |
= moles⁄smallest moles = molar ratio
| Molar Mass
Note: if molar ration is 0.5, multiply all molar ratios by 2; if 0.333, multiply all molar ratios by 3

Molecular Formula Calculation
True Molar Mass
= multiplier for the molar ratio to get molecular formula
Empirical Formula Molar Mass

Empirical Formula Calculation by Combustion
# g − CO2 | mol | 1 mol C | 12 g
= g−C
|12 g |1 mol CO2 | mol
# g − H2 O | mol | 2 mol H | 1 g
=g−H
|18 g |1 mol H2 O| mol
# g Cx Hy Oz − (g of C + g of H) = g − O
o
Stoichiometry
Then follow the Empirical Calculation directions.
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Stoichiometry Problem Set
Molar Calculations
1. Calculate the following quantities:
a. Mass, in grams, of 0.50 mol of Ca(OH)2
b. Number of moles of in 50.0 g of NH4Cl
c. Number of molecules of 0.25 mol of NH3
d. Number of moles of carbon in 2.50 mol of C4H10
e. Molarity of 0.05 mol of KMnO4 in 400 mL of water.
f. Mass, in grams, of NaOH in 500mL of 0.05 M solution.
g. Volume, in mL, of 11.5 M HCl solution needed to make a 500mL solution of 0.5 M.
__ NaN3 (s)  __ Na (s) + __ N2 (g)
2. Balance the reaction and calculate the following quantities:
a. Number of moles of nitrogen gas when 1.50 mol of NaN3.
b. How many grams of NaN3 are needed to produce 10.0 g of nitrogen gas.
c. If NaN3 has a density of 1.25 g/L, how many grams of solid sodium are produced when 750 mL of NaN3
is used?
__ NaHCO3 (s) + __H2SO4 (aq)  __Na2SO4 (aq) + __ H2O (l) + __ CO2 (g)
3. Balance the reaction and calculate the following quantities:
a. If 0.25 moles of NaHCO3 is used in this reaction, how many total moles of products are formed?
b. If excess NaHCO3 is added to 0.500 liters of 2.00 M H2SO4, how many moles of carbon dioxide are
formed?
c. If 100mL of 6.0 M H2SO4 was spilled, what is the minimum mass of NaHCO3 that must be used to
neutralize this spill? Why would we use NaHCO3 and not water to neutralize the spill?
2 NaOH (s) + CO2 (g)  Na2CO3 (s) + H2O (l)
4. Calculate the following quantities based on the reaction above:
a. If 1.50 moles of NaOH and 1.00 moles of CO2 are allowed to react, which is the limiting reactant?
b. How many maximum moles of Na2CO3 are produced? What is the maximum mass in grams produced of
Na2CO3?
c. How many moles of excess reactant remain after the completion of the reaction?
5. A hydrocarbon sample with a mass of 8 grams underwent combustion, producing 22 grams of carbon dioxide. If
all of the carbon initially present in the compound was converted to carbon dioxide, what was the percent of
carbon, by mass, in the hydrocarbon sample?
(A) 25%
(B) 33%
(C) 50%
(D) 66%
(E) 75%
Stoichiometry
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6. What is the mass of oxygen in 37 grams of calcium hydroxide (Ca(OH)2) (MW=74)?
(A) 16 grams
(B) 24 grams
(C) 32 grams
(D) 40 grams
(E) 48 grams
C3H8 + 5 O2  3 CO2 + 4 H2O
7. If approximately 30 liters of CO2 were produced at STP and all of the carbon in the CO2 came from the propane,
what was the mass of the propane sample?
(A) 10 grams
(B) 20 grams
(C) 35 grams
(D) 40 grams
(E) 60 grams
8. A pure sample of KClO3 is found to contain 71 grams of chlorine atoms. What is the mass of the sample?
(A) 122 grams
(B) 170 grams
(C) 209 grams
(D) 245 grams
(E) 293 grams
C3H7OH + …O2  …CO2 + …H2O
9. One mole of C3H7OH underwent combustion as shown in the reaction above. How many moles of oxygen were
required for the reaction?
(A) 2 moles
(B) 3 moles
(C) 7/2 moles
(D) 9/2 moles
(E) 5 moles
CH4 + 2 O2  CO2 + 2 H2O
10. If 16 grams of CH4 reacts with 16 grams of O2 in the reaction shown above, which of the following will be true?
(A) The mass of H2O formed will be twice the mass of CO2 formed.
(B) Equal mass of H2O and CO2 will be formed.
(C) Equal numbers of moles of H2O and CO2 will be formed.
(D) The limiting reagent will be CH4.
(E) The limiting reagent will be O2.
Stoichiometry
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2 ZnS (s) + 3 O2(g)  2 ZnO (s) + 2 SO2 (g)
11. If the reaction above took place at standard temperature and pressure, what was the volume of O2 (g) required to
produced 40.0 grams of ZnO (s)?
(A)
(40.0)(2)
(81.4)(3)(22.4)
(40.0)(3)
𝐿
(B) (81.4)(2)(22.4) 𝐿
(C)
(D)
(E)
(40.0)(2)(22.4)
(81.4)(3)
(40.0)(3)(22.4)
(81.4)(2)
(40.0)(2)(22.4)
(81.4)(2)
𝐿
𝐿
𝐿
Empirical and Molecular Formula Calculations
12. The formula of a specific hydrate is Na2CO3 ∙ x H2O. When a 2.558 gram sample of this hydrate is heated at
25°C, all of the water of hydration is lost, leaving 0.948 g of Na2CO3. Determine the molecular formula of this
hydrate.
Mass (g)
Initial Compound
5.061 g
Compound after heating
2.950 g
Compound after more heating
2.471 g
Compound after more heating
2.472 g
Compound after more heating
2.472 g
13. The formula for an unknown hydrate is MgSO4 ∙ x H2O is decomposed in the laboratory as shown in the table
above. What is the mass of the compound after heating? Explain. What are the moles of water present in the
hydrate?
14. Determine the empirical formula for the following compounds given:
a. 0.0130 mol C, 0.0390 mol H, and 0.0065 mol O
b. 11.66 g of Fe and 5.01 g of oxygen
c. 40% C, 6.67% H, 53.3% O by mass.
d. 0.104 mol K, 0.052 mol C, and 0.156 mol O
e. 5.28 g Sn and 3.37 g F
f. 87.5% N and 12.5% H by mass.
15. Determine the empirical and molecular formula for the following compounds given:
g. Hydrocarbon with 92.3% C by mass and a molar mass of 104 g/mol.
h. 49.5% C, 5.15% H, 28.9%, N, and O by mass with a molar mass of 195 g/mol.
i. 35.51% C, 4.77% H, 37.85% O, 8.29% N, and 13.60% Na with a molar mass of 169 g/mol.
j. 75.69% C, 8.80% H, and O by mass with a molar mass of 206 g/mol.
k. 58.55% C, 13.81% H, 27.40% N by mass with a molar mass of 102.2 g/mol.
l. 59.0% C, 7.1% H, 26.2% O, and 7.7% N by mass with a molecular weight of 180.
Stoichiometry
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16. A 10.00 gram sample of an organic compound is combusted and produces 28.204 of CO2 and 11.538 g of H2O.
Determine the empirical and molecular formula of this compound if its molar mass is 156.0 g/mol.
17. A 2.78 g sample of an organic compound is combusted and produces 6.32 g of CO2 and 2.58 g of H2O.
Determine the empirical formula of this compound and without any calculations, show how to determine the
molecular formula.
18. If a 1 M sample of Fe2O3 ∙ x H2O is found to contain 108 grams of water, what is the molecular formula for the
sample?
(A) Fe2O3 ∙ H2O
(B) Fe2O3 ∙ 3 H2O
(C) Fe2O3 ∙ 6 H2O
(D) Fe2O3 ∙ 10 H2O
(E) Fe2O3 ∙ 12 H2O
19. Rutile, an oxide of titanium has 60% titanium and 40% oxygen in its chemical formula. What is the empirical
formula for rutile?
(A) TiO
(B) TiO2
(C) Ti2O
(D) Ti2O3
(E) Ti3O2
20. A compound containing only sulfur and oxygen is 40% sulfur by weight. What is the empirical formula for the
compound?
(A) SO
(B) SO2
(C) SO3
(D) S2O
(E) S3O
21. A compound containing carbon, hydrogen, and oxygen has 60% carbon, 8% hydrogen, and 32% oxygen by mass.
What is the empirical formula for the compound?
(A) C60H8O32
(B) C7H1O4
(C) C3H8O4
(D) C5H8O2
(E) C4H6O2
Stoichiometry
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Free Response:
22. A student is assigned the task of determining the mass percent of silver in an alloy of copper and silver by
dissolving a sample of the alloy in excess nitric acid and then precipitating silver as AgCl.
First the student prepares 50. mL of 6 M HNO3.
(a) The student is provided with a stock solution of 16 M HNO3, two 100 mL graduated cylinders that can be read
to ±1 mL, a 100 mL beaker that can be read to ±10 mL, safety goggles, rubber gloves, a glass stirring rod, a
dropper, and distilled H2O.
(i)
(ii)
(iii)
(iv)
Calculate the volume, in mL, of 16 M HNO3 that the student should use for preparing 50. mL of 6 M
HNO3.
Briefly list the steps of an appropriate and safe procedure for preparing the 50. mL of 6 M HNO3.
Only materials selected from those provided to the student (listed above) may be used.
Explain why it is not necessary to use a volumetric flask (calibrated to 50.00 mL ± 0.05mL) to
perform the dilution.
During the preparation of the solution, the student accidentally spills about 1 mL of 16 M HNO3 on
the bench top. The student finds three bottles containing liquids near the spill: a bottle of distilled
water, a bottle of 5 percent NaHCO3 (aq), and a bottle of saturated NaCl (aq). Which of the liquids is
best to use in cleaning up the spill? Justify your choice.
Then the student pours 25 mL of the 6 M HNO3 into a beaker and adds a 0.6489 g sample of the alloy. After the
sample completely reacts with the acid, some saturated NaCl (aq) is added to the beaker, resulting in the
formation of an AgCl precipitate. Additional NaCl (aq) is added until no more precipitate is observed to form.
The precipitate is filtered, washed, dried, and weighed to constant mass in a filter crucible as shown in the table
below.
Mass of sample of copper-silver alloy
Mass of dry filter crucible
Mass of filter crucible and precipitate (first weighing)
Mass of filter crucible and precipitate (second weighing)
Mass of filter crucible and precipitate (third weighing)
0.6489 g
28.7210 g
29.3587 g
29.2599 g
29.2598 g
(b) Calculate the number of moles of AgCl precipitate collected.
(c) Calculate the mass percent of silver in the alloy of copper and silver.
23. The table below shows three common forms of copper ore.
Iron Ore #
Empirical Formula
Percent by Weight
Copper
Sulfur
Iron
1
Cu2S
?
?
0
2
?
34.6
34.9
30.5
3
?
55.6
28.1
16.3
(a) What is the percent by weight of copper in Cu2S?
(b) What is the empirical formula of Iron Ore #2?
(c) If a sample of Iron Ore #3 contains 11.0 grams of iron, how many grams of sulfur does it contain?
(d) Cu can be extracted from Cu2S by the following process:
3 Cu2S + 3 O2  3 SO2 + 6 Cu
If 3.84 grams of O2 are consumed in the process, how many grams of Cu are produced?
Stoichiometry
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