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Stoichiometry Practice Calculations
15. For each of the following unbalanced equations, indicate how many moles of the second reactant would be
required to react exactly with 0.275 mol of the first reactant. State clearly the mole ratio used for the conversion.
a. Cl2(g) + KI(aq)  I2(s) + KCl(aq)
b. Co(s) + P4(s)  Co3P2(s)
c. Zn(s) + HNO3(aq)  ZnNO3(aq) + H2(g)
d. C5H12(l) + 02(g)  CO2(g) + H20(g)
16. For each of the following unbalanced equations, indicate how many moles of the first product are produced if
0.625 mol of the second product forms state clearly the mole ratio used for each conversion.
a. KO2(s) + H2O(l)  O2(g) + KOH(s)
b. SiO2(g) + H2Si(g)  Si(s) + H2O(g)
c. CH3CH2OH(l) + O2(g)  CH3CHO(aq) + H2O(l).
d, Fe203(s) + Al(s) Fe(l) + Al2O3(s)
Mass Calculations
17. What quantity serves as the conversion factor between the mass of a sample and how many moles the sample
contains?
18. What does it mean to say that the balanced chemical equation for a reaction describes the stoichiometry of the
reaction?
19. Using the average atomic masses, calculate how many moles of each substance the following masses represent.
a. 14.2 g of silver
b. 729 g of cobalt(II) chloride
c. 4.16g of lithium carbonate.
d. 1.25 mg of lead (II) sulfate
e. 0.104 g of potassium
f. 1.21 lb of iron
g. 62.4 g of magnesium oxide
20. Using the average atomic masses, calculate the number of moles of each substance contained in the following
masses.
a. 2.01 x 102 g of silver
b. 45.2 mg of ammonium sulfide
c. 61.7 g of titanium
d. 5.23 kg of sulfur dioxide
e. 272 g of iron(III) nitrate
f. 12.7 mg of iron(II) sulfate
g. 6.91 x 103 g of lithium hydroxide
21. Using the average atomic masses, calculate the mass in grams of each of the following samples.
a. 1.25 mol of platinum(IV) chloride
b. 0.00255 mol of copper(II) oxide
c. 1.89 x 10-4 mol of ethane, C2H6
d. 55.56 mol of beryllium
e. 2.6 x 107 mol of diboron trioxide
f. 0.45 mol of sodium fluoride
g. 0.00115 mol of calcium nitrate
22. Using the average atomic masses, calculate the mass in grams of each of the following samples.
a. 2.21 x 10-4 mol of calcium carbonate
b. 2.75 rnol of helium
c. 0.00975 mol of oxygen gas
d. 7.21 mmol of carbon dioxide
e. 0.835 mol of iron(II) sulfide
f. 4.01 mol of potassium hydroxide
g. 0.0219 mol of. hydrogen gas
23. For each of the following unbalanced equations, indicate how many moles of each product could be produced
by complete reaction of 1.00 g of the reactant indicated in boldface.
a. LiOH(s) + CO2(g)  Li2CO3(s) + H2O(l)
b. Ba(OH)2(s)  BaO(s) + H20(g)
c. C2H4(g) + Cl2(g)  C2H4Cl2(l)
d. H2SO4(aq) + NaOH(aq)  Na2SO4(aq) + H2O(l)
24. For each of the following unbalanced equations, calculate how many moles of the second reactant would be
required to react completely with exactly 25.0 g of the first reactant.
a. Mg(s) + CuCl2(a)  MgCl2(aq) + Cu(s)
b. AgNO3(aq) + NiCl2(aq)  AgCl(s) + Ni(NO3)2(aq)
c. NaHSO3(aq) + NaOH(aq)  Na2SO3(aq) + H2O(l)
d. KHCO3(aq) + HCl(aq)  KCl(aq) + H2O(l) + CO2(g) ,
25. For each of the following unbalanced equations, calculate how many grams of each product would be produced
by complete reaction of 12.5 g of the reactant indicated in boldface.
a. TiBr4(g) + H2(g)  Ti(s) + HBr(g)
b. SiH4(g) + NH3(g)  Si3N4(s) + H2(g)
c. NO(g).+ H2(g)  N2(g) + 2H2O(l)
d. Cu2S(s)  Cu(s) + S(g)
26. For each of the following unbalanced equations, calculate how many grams of each product would be produced
by complete reaction of 10.0 g of the reactant indicated in boldface.
a. FeSO4(aq) + K2CO3(aq)  FeCO3(s) + K2SO4(aq)
b. Cr(s) + SnCl4(l)  CrCl3(s) + Sn(s)
c. Fe(s) + S8(s)  Fe2S3(s)
d. Ag(s) + HNO3(aq)  AgNO3(aq) + H2O(l) + NO(g)
27. Although mixtures of hydrogen and oxygen are highly explosive, pure elemental hydrogen gas itself burns
quietly in air with a pale blue flame, producing water vapor.
Calculate the mass (in grams) of water vapor produced when 56.0 g of pure hydrogen gas burns in air.
28. Given the information in Problem 27, calculate the mass of oxygen gas that would be necessary to burn 0.0275
mol of hydrogen gas.
29. When elemental carbon is burned in the open atmosphere, with plenty of oxygen gas present, the product is
carbon dioxide.
C(s) + O2(g) — CO2(g)
However, when the amount of oxygen prescnt during the burning of the carbon is restricted, carbon monoxide is
more likely to result.
2C(s) + 02(g) — 2C0(g)
What mass of each product is expected when a 5.00-g sample of pure carbon is burned under each of these
conditions?
30. The more reactive halogen elements are able to replace the less reactive halogens from their compounds. For ex
ample, if chlorine gas is bubbled through a potassium iodide solution, elemental iodine is produced.
C12(g) + KI(aq)  I2(s) + KCl(aq)
Calculate the mass of iodine produced when 2.55 g of chlorine gas is bubbled through an excess amount of
potassium iodide solution.
31. Although we usually think of substances as “burning” only in oxygen gas, the process of rapid oxidation to
produce a flame may also take place in other strongly oxidizing gases. For example, when lion is heated and placed
in pure chlorine gas, the iron “burns” according to the following (unbalanced) reaction:
How many milligrams of iron(III) chloride result when 15.5 mg of iron is reacted with an excess of chlorine gas?
Fe(s) + Cl2(g)  FeCl(s)
32. You have probably seen images of a chef preparing a “flaming” dessert or entree. The flame is usually the result
of the combustion of ethyl alcohol, CH5OH, that has been added to the food (perhaps as cognac or rum).
C2H5OH(l) + O2(g)  CO2(g) + H2O(l)
If 25.0 g of ethyl alcohol is burned in air (excess oxygen), calculate the mass of carbon dioxide produced.
33. Small quantities of oxygen gas can be generated in the laboratory by the decomposition of hydrogen peroxide.
The unbalanced equation for the reaction is
H2O2(aq)  H2O(l) + O2(g)
Calculate the mass of oxygen produced when 10.00 g of hydrogen peroxide decomposes.
34. Elemental fluorine and chlorine gases are very reactive. For example, they react with each other to produce an
interhalogen compound, chlorine monofluoride.
Cl2(g) + F2(g)  2ClF(g)
Calculate how many grams of chlorine gas are required to react completely with 500 mg of fluorine gas.
35. With the news that calcium deficiencies in many women’s diets may contribute td the development of
osteoporosis (bone weakening), the use of dietary calcium supplements has become increasingly common. Many of
these calcium supplements consist of nothing more than calcium carbonate, CaCO3. When a calcium carbonate
tablet is ingested, it dissolves by reaction with stomach acid, which contains hydrochloric acid, HCl. The
unbalanced equation is
CaCO3(s) + HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
What mass of HCl is required to react with a tablet containing 500. mg of calcium carbonate?
36. A tried-and-true introductory chemistry experiment involves heating finely divided copper metal with sulfur to
determine the proportions in which the elements react to form copper(II) sulfide. The experiment works well,
because any excess sulfur beyond that required to react with the copper may be simply boiled away from the
reaction container.
Cu(s) + S(s)  CuS(s)
If 1.25 g of copper is heated with an excess of sulfur, how many grams of sulfur will react?