Unit 3/4 – Moles / Stoichiometry
Chemistry Review
Unit 3/4 – Moles / Stoichiometry
Formula Writing, Naming & Writing Chemical Compound Formulas, Chemical Equations, Mole Interpretation, Stoichiometry
Moles and Stoichiometry
1. A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can only be broken down by chemical means.
2. Chemical compounds can be represented by a specific formula and assigned a name based on the IUPAC
system.
3. Types of chemical formulas include empirical, molecular, and structural.
 Empirical formulas show elements in their simplest whole number ratios. This may or may not be the same as the
molecular formula.
 Molecular formulas show the actual number of atoms per element in a single molecule.
 Structural formulas show the number of each type of atom as well as their physical arrangement.
4. All chemical reactions show a conservation of mass, energy and charge.
5. A balanced chemical equation represents conservation of atoms.
6. The coefficients in a balanced chemical equation can be used to determine mole ratios in the reaction.
7. The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram formula
mass) equals the mass of one mole of that substance.
8. The percent composition by mass of each element in a compound can be calculated mathematically.
9. Types of chemical reactions include synthesis, decomposition single replacement, and double replacement.
-1-
Unit 3/4 – Moles / Stoichiometry
June 2013
January 2013
June 2012
-2-
Unit 3/4 – Moles / Stoichiometry
10: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
37: 3.3viii Calculate the formula mass and gram-formula mass
51: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts
per million (ppm)
52: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
53: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction.
54: 3.2a A physical change results in the rearrangement of existing particles in a substance. A chemical change results in
the formation of different substances with changed properties.
55: 3.3ix Determine the number of moles of a substance, given its mass
3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
68: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically.
-3-
Unit 3/4 – Moles / Stoichiometry
January 2012
-4-
Unit 3/4 – Moles / Stoichiometry
17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles.
33: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
36: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting
various scales of measurement
37: 4.1d Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between the
potential energy of the products and potential energy of the reactants
66: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system.
74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction.
76: 3.3ix Determine the number of moles of a substance, given its mass
77: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically.
78: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound.
June 2011
-5-
Unit 3/4 – Moles / Stoichiometry
7: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
9: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula
mass) of a substance equals one mole of that substance.
34: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
-6-
Unit 3/4 – Moles / Stoichiometry
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system.
35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
37: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound.
39: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
41: 3.2a A physical change results in the rearrangement of existing particles in a substance. A chemical change results in
the formation of different substances with changed properties.
46: 3.2d An oxidation-reduction (redox) reaction involves the transfer of electrons (e-).
49: 5.3c Energy released during nuclear reactions is much greater than the energy released during chemical reactions.
66: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction.
January 2011
-7-
Unit 3/4 – Moles / Stoichiometry
10: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula
mass) of a substance equals one mole of that substance.
39: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
47: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
55: Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
69: Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
76: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that
oxidation and reduction have occurred.
August 2010
-8-
Unit 3/4 – Moles / Stoichiometry
7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction.
9: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically.
18: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles
36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
70: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
72: 3.3viii Calculate the formula mass and gram-formula mass
June 2010
-9-
Unit 3/4 – Moles / Stoichiometry
17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles.
34:
36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
56: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording
experimental data and use data in calculations such as using appropriate equations and significant digits
73: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
75: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
- 10 -
Unit 3/4 – Moles / Stoichiometry
January 2010
9: 3.3e The formula mass of a substance is the sum of the atomic masses of its atoms. The molar mass (gram-formula
mass) of a substance equals one mole of that substance
10: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
22: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles
23: 5.2m Intermolecular forces created by the unequal distribution of charge result in varying degrees of attraction
between molecules. Hydrogen bonding is an example of a strong intermolecular force
36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
- 11 -
Unit 3/4 – Moles / Stoichiometry
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
68: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
August 2009
- 12 -
Unit 3/4 – Moles / Stoichiometry
35: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
40: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts
per million (ppm)
52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
55: 3.3vi Determine the mass of a given number of moles of a substance
56: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
57: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
79: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
June 2009
- 13 -
Unit 3/4 – Moles / Stoichiometry
7: 3.3vi Determine the mass of a given number of moles of a substance
33: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
34: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
79: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting
various scales of measurement
80: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
81: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
83: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
January 2009
- 14 -
Unit 3/4 – Moles / Stoichiometry
- 15 -
Unit 3/4 – Moles / Stoichiometry
- 16 -
Unit 3/4 – Moles / Stoichiometry
11: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
17: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles
22: 3.3b In a redox reaction the number of electrons lost is equal to the number of electrons gained
36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
69: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
71: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
- 17 -
Unit 3/4 – Moles / Stoichiometry
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
76: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
78: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
August 2008
34: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can
be used to determine mole ratios in the reaction
- 18 -
Unit 3/4 – Moles / Stoichiometry
35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
63: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
74: 3.3vi Determine the mass of a given number of moles of a substance
75: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
76: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
June 2008
- 19 -
Unit 3/4 – Moles / Stoichiometry
7: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
19: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles
35: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording
experimental data and use data in calculations such as calculating percent error
36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
42: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
47: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
- 20 -
Unit 3/4 – Moles / Stoichiometry
compound
56: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
67: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
3.3vi Determine the mass of a given number of moles of a substance
3.3viii Calculate the formula mass and gram-formula mass
3.3ix Determine the number of moles of a substance, given its mass
74: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
January 2008
34: 3.3viii Calculate the formula mass and gram-formula mass
54: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
77: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
78: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
79: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
- 21 -
Unit 3/4 – Moles / Stoichiometry
August 2007
10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
12: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system.
35: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
70: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
71: 3.3viii Calculate the formula mass and gram-formula mass
72: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
- 22 -
Unit 3/4 – Moles / Stoichiometry
June 2007
7: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
36: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
55: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
59: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
60: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that
oxidation and reduction have occurred.
- 23 -
Unit 3/4 – Moles / Stoichiometry
January 2007
10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
17: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
18: 3.1t The proportions of components in a mixture can be varied. Each component in a mixture retains its original
properties.
35: 3.3viii Calculate the formula mass and gram-formula mass
36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
54: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording
experimental data and use data in calculations such as calculating percent error
- 24 -
Unit 3/4 – Moles / Stoichiometry
August 2006
- 25 -
Unit 3/4 – Moles / Stoichiometry
- 26 -
Unit 3/4 – Moles / Stoichiometry
6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system.
9: 5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds,
and ions
37: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
38: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
44: 4.1b Chemical and physical changes can be exothermic or endothermic.
82: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
83: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
84: 3.3ix Determine the number of moles of a substance, given its mass
85: 3.2x Use an activity series to determine whether a redox reaction is spontaneous
June 2006
- 27 -
Unit 3/4 – Moles / Stoichiometry
6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
- 28 -
Unit 3/4 – Moles / Stoichiometry
33: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
35: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
51: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
52: 3.2x Use an activity series to determine whether a redox reaction is spontaneous
69: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
70: 3.4ii Solve problems, using the combined gas laws
January 2006
8: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
10: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
33: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
36: 3.3viii Calculate the formula mass and gram-formula mass
57: 3.3ix Determine the number of moles of a substance, given its mass
- 29 -
Unit 3/4 – Moles / Stoichiometry
75: 3.3viii Calculate the formula mass and gram-formula mass
76: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts
per million (ppm)
77: 3.1xxx Describe the preparation of a solution, given the molarity
August 2005
9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
31: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
- 30 -
Unit 3/4 – Moles / Stoichiometry
36: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
38: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
48: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
51: 3.3viii Calculate the formula mass and gram-formula mass
52: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
68: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
69: 3.3ix Determine the number of moles of a substance, given its mass
70: Density equation – General math standard
June 2005
9: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
36: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
39: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
54: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
January 2005
- 31 -
Unit 3/4 – Moles / Stoichiometry
8: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
9: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
19: 3.4e Equal volumes of gases at the same temperature and pressure contain an equal number of particles
34: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
35: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can
be used to determine mole ratios in the reaction
68: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
69: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
August 2004
- 32 -
Unit 3/4 – Moles / Stoichiometry
- 33 -
Unit 3/4 – Moles / Stoichiometry
8: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
29: 3.1zz Titration is a laboratory process in which a volume of a solution of known concentration is used to determine
the concentration of another solution.
37: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
50: M1.1B – Use algebraic and geometric representations to describe and compare data by measuring and recording
experimental data and use data in calculations such as using appropriate equations and significant digits
52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
53: 3.1ppThe concentration of a solution may be expressed in molarity (M), percent by volume, percent by mass, or parts
per million (ppm)
61: 3.3viii Calculate the formula mass and gram-formula mass
62: 3.3vi Determine the mass of a given number of moles of a substance
66: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
67: 4.1c Energy released or absorbed during a chemical reaction can be represented by a potential energy diagram
- 34 -
Unit 3/4 – Moles / Stoichiometry
68: 3.2i Oxidation numbers (states) can be assigned to atoms and ions. Changes in oxidation numbers indicate that
oxidation and reduction have occurred
69: M1.1C – Use algebraic and geometric representations to describe and compare data by recognizing and converting
various scales of measurement
June 2004
7: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
8: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
38: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
51: 3.2b Types of chemical reactions include synthesis, decomposition, single replacement, and double replacement
- 35 -
Unit 3/4 – Moles / Stoichiometry
52: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can
be used to determine mole ratios in the reaction
53: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
81: 5.2d Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements,
compounds, and ions
82: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
83: 3.1xxix Calculate solution concentration in molarity (M), percent mass, and parts per million (ppm)
84: S3.4 Using results of the test and through public discussion, revise the explanation and contemplate additional
research
January 2004
- 36 -
Unit 3/4 – Moles / Stoichiometry
6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
36: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
39: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge.
3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation can
be used to determine mole ratios in the reaction
51: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a balanced chemical equation
can be used to determine mole ratios in the reaction
52: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
75: 3.3viii Calculate the formula mass and gram-formula mass
76: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
- 37 -
Unit 3/4 – Moles / Stoichiometry
August 2003
6: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
8: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
10: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
39: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
42: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
June 2003
- 38 -
Unit 3/4 – Moles / Stoichiometry
8: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of the elements in a
compound. It may be different from the molecular formula, which is the actual ratio of atoms in a molecule of that
compound
10: 3.3f The percent composition by mass of each element in a compound can be calculated mathematically
19: 3.1cc A compound is a substance composed of two or more different elements that are chemically combined in a fixed
proportion. A chemical compound can be broken down by chemical means. A chemical compound can be represented by
a specific chemical formula and assigned a name based on the IUPAC system
20: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
59: 3.3a In all chemical reactions there is a conservation of mass, energy, and charge
January 2003
- 39 -
Unit 3/4 – Moles / Stoichiometry
8: 3.3f The percent composition by mass of each element in a compound can be calculated
mathematically
13: 3.3d The empirical formula of a compound is the simplest whole-number ratio of atoms of
the elements in a compound. It may be different from the molecular formula, which is the actual
ratio of atoms in a molecule of that compound
22: 3.3f The percent composition by mass of each element in a compound can be calculated
mathematically
42: 3.3iv Calculate simple mole-mole stoichiometry problems, given a balanced equation
43: M1.1B – Use algebraic and geometric representations to describe and compare data by
measuring and recording experimental data and use data in calculations such as calculating
percent error
48: 3.3c A balanced chemical equation represents conservation of atoms. The coefficients in a
balanced chemical equation can be used to determine mole ratios in the reaction
- 40 -
Unit 3/4 – Moles / Stoichiometry
1) How many moles of sodium atoms correspond to 1.56x1021 atoms of sodium?
2) How many moles of Al atoms are needed to combine with 1.58 mol of O atoms to
make aluminum oxide, Al2O3?
3) How many moles of Al are in 2.16 mol of Al2O3?
4) Aluminum sulfate, Al2(SO4)3, is a compound used in sewage treatment plants.
a. Construct a pair of conversion factors that relate moles of aluminum to
moles of sulfur for this compound
b. Construct a pair of conversion factors that relate moles of sulfur to moles of
Al2(SO4)3
c. How many moles of Al are in a sample of this compound if the sample also
contains 0.900 mol S?
d. How many moles of S are in 1.16 mol Al2(SO4)3?
5) How many moles of H2 and N2 can be formed by the decomposition of 0.145 mol
of ammonia, NH3?
6) What is the total number of atoms in 0.260 mol of glucose, C6H12O6?
7) What is the mass of 1.00 mol of each of the following elements?
a. Sodium
b. Sulfur
c. Chlorine
8) Determine the mass in grams of each of the following:
a. 1.35 mol Fe
b. 24.5 mol O
c. 0.876 mol Ca
d. 1.25 mol Ca3(PO4)2
e. 0.625 mol Fe(NO3)3
f. 0.600 mol C4H10
g. 1.45 mol (NH4)2CO3
9) Calculate the number of moles of each compound:
a. 21.5 g CaCO3
b. 1.56 g NH3
c. 16.8 g Sr(NO3)2
d. 6.98 g Na2CrO4
41
Unit 3/4 – Moles / Stoichiometry
Percent composition and empirical formulas
10) Calculate the percentage composition by mass of each element in the following
compounds:
a. NaH2PO4
b. NH4H2PO4
c. (CH3)2CO
11) Phencyclidine is C17H25N. A sample suspected of being this illicit drug was found
to have a percentage composition of 83.71% C, 10.42% H, and 5.61% N. Do
these data acceptably match the theoretical data for phencyclidine?
12) How many grams of O are combined with 7.14x1021 atoms of N in the compound
N2O5?
13) Quantitative analysis of a sample of sodium pertechnetate with a mass of 0.896g
found 0.111g Na and 0.477g technetium (Tc). The remainder was
oxygen. Calculate the empirical formula of sodium pertechnetate, NaxTcyOz.
14) A substance was found to be composed of 22.9% Na, 21.5% B, and 55.7%
O. What is the empirical formula of this compound?
15) When 0.684 g of an organic compound containing only C, H, and O was burned
in oxygen 1.312g CO2 and 0.805g H2O were obtained. What is the empirical
formula of the compound?
Balancing equations
16) Write the equation that expresses in acceptable chemical shorthand the following
statement: “Iron can be made to react with molecular oxygen (O2) to give iron
oxide with the formula Fe2O3”
17) Balance the following reactions:
a. Ca(OH)2 + HCl  CaCl2 + H2O
b. AgNO3 + CaCl2  Ca(NO3)2 +AgCl
c. Fe2O3 + C  Fe + CO3
d. NaHCO3 + H2SO4  Na2SO4 + H2O + CO2
e. C4H10 + O2  CO2 +H2O
f. Mg(OH)2 + HBr  MgBr2 + H2O
g. Al2O3 + H2SO4  Al2(SO4)3 + H2O
h. KHCO3 + H3PO4  K2HPO4 + H2O + CO2
i. C9H10O + O2  CO2 + H2O
Stoichiometry/limiting reactants
18) Chlorine is used by textile manufacturers to bleach cloth. Excess chlorine is
destroyed by its reaction with sodium thiosulfate, Na2S2O3:
Na2S2O3(aq) + 4Cl2(g) + 5H2O(aq)  2NaHSO4(aq) + 8HCl(aq)
42
Unit 3/4 – Moles / Stoichiometry
a.
b.
c.
d.
How many moles of Na2S2O3 are needed to react with 0.12mol of Cl2?
How many moles of HCl can form from 0.12mol of Cl2?
How many moles of H2O are required for the reaction of 0.12mol of Cl2?
How many moles of H2O react if 0.24mol HCl is formed?
19) The incandescent white of a fireworks display is caused by the reaction of
phosphorous with O2 to give P4O10.
a. Write the balanced chemical equation for the reaction.
b. How many grams of O2 are needed to combine with 6.85g of P?
c. How many grams of P4O10 can be made from 8.00g of O2?
d. How many grams of P are needed to make 7.46g P4O10?
20) In dilute nitric acid, HNO3, copper metal dissolves according to the following
equation:
3Cu(s) + 8HNO3(aq)  3Cu(NO3)2(aq) + 2NO(g) + 4H2O(aq)
How many grams of HNO3 are needed to dissolve 11.45g of Cu?
21) The reaction of powdered aluminum and iron(II)oxide,
2Al(s) + Fe2O3(s)  Al2O3(s) + 2Fe(l)
produces so much heat the iron that forms is molten. Because of this, railroads use
the reaction to provide molten steel to weld steel rails together when laying
track. Suppose that in one batch of reactants 4.20mol Al was mixed with 1.75mol
Fe2O3.
a. Which reactant, if either, was the limiting reactant?
b. Calculate the mass of iron (in grams) that can be formed from this mixture
of reactants.
22) Silver nitrate, AgNO3, reacts with iron(III) chloride, FeCl3, to give silver chloride,
AgCl, and iron(III) nitrate, Fe(NO3)3. A solution containing 18.0g AgNO3 was
mixed with a solution containing 32.4g FeCl3. How many grams of which
reactant remains after the reaction is over?
Theoretical and percent yield
23) Barium sulfate, BaSO4, is made by the following reaction:
Ba(NO3)2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaNO3(aq)
An experiment was begun with 75.00g of Ba(NO3)2 and an excess of
Na2SO4. After collecting and drying the product, 63.45g BaSO4 was
obtained. Calculate the theoretical yield and percent yield of BaSO4.
24) Aluminum sulfate can be made by the following reaction:
2AlCl3(aq) + 3H2SO4(aq)  Al2(SO4)3(aq) + 6HCl(aq)
It is quite soluble in water, so to isolate it the solution has to be evaporated to
dryness. This drives off the volatile HCl, but the residual solid has to be treated to
a little over 200C to drive off all the water. In one experiment, 25.0g of
AlCl3 was mixed with 30.0g H2SO4. Eventually, 28.46g of pure Al2(SO4)3 was
isolated. Calculate the percent yield.
43
Unit 3/4 – Moles / Stoichiometry
Answers
2.59x103mol Na atoms
1.05mol Al
4.32mol Al
a. 2mol Al/3mol S
b. 3mol S/1mol Al2(SO4)3
c. 0.600mol Al d.
3.48mol S
5) 0.0725mol N2 and 0.218mol H2
6) 3.76x1024 atoms
7) a. 23.0g Na
b. 32.1g S
c. 35.3g Cl
8) a. 75.4g Fe
b. 392g O
c. 35.1g Ca
d. 388g
Ca3(PO4)2
e. 151g Fe(NO3)2
f. 34.9g C4H10
g. 139g (NH4)2CO3
9) a. 0.215mol
b. 0.0916mol
c. 0.0794mol
d.

8
4.31x10 mol
10) a. 19.2% Na, 1.68% H, 25.8% P, 53.3% O
b. 12.2% N, 5.26% H, 26.9% P, 55.6%O
c. 62.0% C, 10.4% H, 27.6% O
11) Theoretical data (83.89% C, 10.35% H, 5.76% N) are consistent with
experimental results.
12) 0.474g O
13) NaTcO4
14) Na2B4O7
15) C2H6O
16) 4Fe + 3O2  2Fe2O3
17)
a. Ca(OH)2 + 2HCl  CaCl2 + 2H2O
b. 2AgNO3 + CaCl2  Ca(NO3)2 + 2AgCl
c. 2Fe2O3 + 3C  4Fe + 3CO3
d. 2NaHCO3 + H2SO4  Na2SO4 + 2H2O + 2CO2
e. 2C4H10 + 13O2  8CO2 + 10H2O
f. Mg(OH)2 + 2HBr  MgBr2 + 2H2O
g. Al2O3 + 3H2SO4  Al2(SO4)3 + 3H2O
h. 2KHCO3 + H3PO4  K2HPO4 + 2H2O + 2CO2
i. C9H10O + 14O2  9CO2 + 10H2O
18) a. 0.030mol Na2S2O3
b. 0.24mol HCl
c. 0.15mol H2O
d. 0.15mol H2O
19) a. 4P + 5O2  P4O10
b. 8.85g O2
c. 14.2g P4O10
d.
3.26g P
20) 30.31g HNO3
21) a. limiting reactant is Fe2O3
b. 195g Fe is formed
22) 26.7g of FeCl3 are left over
23) theoretical yield = 66.98g BaSO4, % yield = 94.73%
24) % yield = 88.74%
1)
2)
3)
4)
1.
Which unit is most analogous to a mole?
44
Unit 3/4 – Moles / Stoichiometry
a fluid ounce
a dozen
a mile
a degree Fahrenheit
a pound
2.
In order to balance the equation C2H6 + O2 ---> H2O + CO2, you
should
change the subscript of O in water to 2 to help balance the O.
add O2 to the product side to help balance the O in the equation.
change the coefficients.
add H2 to the products to balance H.
3.
Balance the following equation:
B10H18 + O2
4.
B10H18 + 7O2
5B2O3 + 9H2O
B10H18 + 9O2
5B2O3 + 9H2O
B10H18 + 19O2
10B2O3 + 9H2O
B10H18 + 12O2
5B2O3 + 9H2O
Balance the following equation and indicate whether it is a
combustion, combination, or decomposition reaction.
H2O2(l)
5.
B2O3 + H2O
H2O(l) + O2(g)
2H2O2(l)
2H2O(l) + O2(g), decomposition reaction
H2O2(l)
H2O(l) + O2(g), decomposition reaction
H2O2(l)
H2O(l) + (1/2)O2(g), decomposition reaction
H2O2(l)
H2O(l) + (1/2)O2(g), combination reaction
H2O2(l)
H2O(l) + O2(g), combustion reaction
Convert the following description into a balanced equation:
45
Unit 3/4 – Moles / Stoichiometry
When ammonia gas, NH3(g), is passed over hot sodium, hydrogen
gas is released and sodium amide, NaNH2, is formed as a solid
product.
Be sure to indicate the state of each element or compound.
2NH3 + Na
NH3 + Na
2NH3(g) + 2Na(s)
2NH3 + 2Na
2NH3(g) + 2Na(s)
6.
2NaNH2 + H2
NaNH2 + H2
2NaNH2(s) + H2(g)
2NaNH2 + H2
2NaNH2(s) + H2(l)
How many molecules of CH2O are in 30.0 g of CH2O?
1.00
6.02 x 1023
1.81 x 1025
5.32 x 10-23
7.
Calculate the mass in mg of Na+ in 10.0 g of sodium carbonate.
4.34 x 103 mg
46.0 mg
2.77 mg
4.34 mg
0.23 mg
8.
How many carbon atoms are there in 200 molecules of C3H8O?
1.20 x 1026
200
3.61 x 1026
600
9.
A sample of vitamin A, C20H30O, contains 4.0 x 1022 atoms of carbon.
How many atoms of hydrogen and how many molecules of vitamin A
does it contain?
46
Unit 3/4 – Moles / Stoichiometry
6.0 x 1022 atoms of H, 8.0 x 1023 molecules of vitamin A
4.0 x 1022 atoms of H, 4.0 x 1022 molecules of vitamin A
6.0 x 1022 atoms of H, 4.0 x 1022 molecules of vitamin A
6.0 x 1022 atoms of H, 2.0 x 1021 molecules of vitamin A
10
.
The element zinc consists of five isotopes with masses 63.929,
65.926, 66.927, 67.925, and 69.925 amu. The relative abundances of
these five isotopes are 48.89, 27.81, 4.110, 18.57, and 0.62 percent,
respectively. From these data calculate the average atomic mass of
zinc.
66.93 amu
65.39 amu
66.927 amu
65.389 amu
63.93 amu
11
.
What is the mass in grams of 0.257 mol of sucrose, C12H22O11?
12.5 g
88.0 g
7.51 x 10-4 g
8.80 g
342 g
12
.
Determine the approximate formula weight of the following:
Ca(C2H3O2)2
69
158
152
99
94
Problem Solving Center
Homework 2
1.
What numbers would properly balance the reaction ___ C12H26 + ___
O2 ---> ___ CO2 + ___ H2O ?
47
Unit 3/4 – Moles / Stoichiometry
1, 25, 12, 13
2, 37, 24, 26
2, 37, 24, 13
1, 38, 12, 26
2.
Complete the following statement: _____________ are in 10.0 moles
of C10H8.
6.022 x 1024 atoms of C
10.0 moles of C
8.00 moles of H
4.818 x 1024 atoms of H
4.818 x 1025 atoms of H
3.
Balance the following equation:
Al + Cr2O3
Al2O3 + Cr
2Al + Cr2O3
Al2O3 + 2Cr
2Al + Cr2O3
Al2O3 + Cr
Al + Cr2O3
Al2O3 + 2Cr
4Al + 2Cr2O3
4.
Balance the following equation:
C6H14O + O2
2C6H14O + 18O2
2C6H14O + 4O2
C6H14O + (19/2)O2
C6H14O + 9O2
2C6H14O + 19O2
5.
2Al2O3 + 4Cr
CO2 + H2O
12CO2 + 14H2O
2CO2 + 2H2O
6CO2 + 7H2O
6CO2 + 7H2O
12CO2 + 14H2O
Balance the following equation and indicate whether it is a
combustion, combination, or decomposition reaction.
48
Unit 3/4 – Moles / Stoichiometry
H2O2 + SO2
H2SO4
H2O2 + SO2
H2SO4, decomposition reaction
2H2O2 + SO2
H2SO4, decomposition reaction
2H2O2 + SO2
H2SO4, combination reaction
H2O2 + SO2
6.
H2SO4, combination reaction
Calculate the number of molecules in 6.2 g of formaldehyde, CH2O.
3.7 x 1024
1.2 x 1023
6.0 x 1023
2.4 x 1023
7.
Calculate the mass in grams of 0.0112 mol of
-fructose, C6H12O6.
180 g
1.12 g
2.02 g
0.0112 g
8.
A sample of glucose, C6H12O6, contains 4.0 x 1022 atoms of carbon.
How many atoms of hydrogen and how many molecules of glucose
does it contain?
8.0 x 1022 atoms of H, 6.7 x 1021 molecules of glucose
8.0 x 1022 atoms of H, 8.0 x 1022 molecules of glucose
8.0 x 1022 atoms of H, 2.4 x 1023 molecules of glucose
4.0 x 1022 atoms of H, 4.0 x 1022 molecules of glucose
8.0 x 1022 atoms of H, 4.0 x 1022 molecules of glucose
9.
The element oxygen consists of three isotopes with masses
15.994915, 16.999133, and 17.99916. The relative abundances of
these three isotopes are 99.7587, 0.0374, and 0.2039, respectively.
From these data calculate the average atomic mass of oxygen.
49
Unit 3/4 – Moles / Stoichiometry
15.9930
15.999377
15.994915
16.0
15.9994
15.9563
10
.
Calculate the number of atoms in 48.0 g glucose, C6H12O6.
1.60 x 1023
9.24 x 1025
3.85 x 1024
2.89 x 1025
1.64 x 1016
11
.
Calculate the molecular weight of xenon tetrafluoride, XeF4, a
colorless, crystalline compound at room temperature.
601.2
150.3
76
169.3
207.3
12
.
How many iron ions (Fe3+) are present in 43.6 g FeCl3?
2.63 x 1025
4.86 x 1023
0.807
3.72
1.62 x 1023
Problem Solving Center
Homework 3
50
Unit 3/4 – Moles / Stoichiometry
1.
Which of these samples contains the most atoms?
a gram of germanium
a gram of francium
a gram of americium
a gram of gallium
a gram of europium
2.
What is the percent yield of CaO in the reaction CaCO 3 ---> CaO + CO2 if 5.33 g of CaO
are obtained when 10.0 g of CaCO3 are used?
5.60 percent
64.7 percent
53.3 percent
5.33 percent
95.2 percent
3.
Calculate the percentage of carbon present in cadaverine, C5H14N2, a compound present
in rotting meat.
58.8 percent C
51
Unit 3/4 – Moles / Stoichiometry
68.2 percent C
67.4 percent C
51.7 percent C
4.
What is the empirical formula of a compound that contains 7.989 g of carbon and 2.011 g
of hydrogen?
C2H5
C8H2
CH3
C3H
C2H6
5.
Give the empirical formula of the following compound if a sample contains 40.0 percent C,
6.7 percent H, and 53.3 percent O by mass.
C2H4O2
CH20
C3H6O3
C6HO8
6.
Determine the empirical formula of a compound that contains 52.9 percent aluminum and
47.1 percent oxygen.
AlO
Al2O3
Al4O6
Al0.53O0.47
Al3O2
7.
In making H2O from hydrogen and oxygen, if we start with 4.6 mol of hydrogen and 3.1
mol of oxygen, how many moles of water can be produced and what remains unreacted?
2.3 mol of water would be produced, with 0.8 mol of O2 remaining.
4.6 mol of water would be produced, with 0.8 mol of O2 remaining.
7.7 mol of water would be produced, with 0.0 mol of O2 remaining.
4.6 mol of water would be produced, with 0.0 mol of O2 remaining.
3.1 mol of water would be produced, with 1.5 mol of O2 remaining.
52
Unit 3/4 – Moles / Stoichiometry
8.
Automotive airbags inflate when sodium azide, NaN 3, rapidly decomposes to its
component elements via the reaction
2NaN3
2Na + 3N2.
How many grams of sodium azide are required to form 5.00 g of nitrogen gas?
3.33 g
11.61 g
15.48 g
7.74 g
9.
What is the molecular formula of the following compound?
empirical formula C2H3, molar mass 54 g/mol
C2H3
C6H9
C4H6
C8H12
10
.
Aluminum and bromine react vigorously according to the following equation:
2Al(s) + 3Br2(l)
2AlBr3(s)
What mass of product can be made by reacting 5.0 g of aluminum and 25 g of bromine?
11 g
28 g
62 g
49 g
42 g
11
.
The alcohol in "gasohol" burns according to the following equation:
C2H5OH + 3O2
2CO2 + 3H2O
How many grams of CO2 are produced when 3.00 g of C2H5OH are burned in this way?
53
Unit 3/4 – Moles / Stoichiometry
5.74 g
88.0 g
6.00 g
2.87 g
0.130 g
12
.
How many moles of H2O are produced when 2.5 mol of O2 react according to the following
equation?
C3H8 + 5O2
3CO2 + 4H2O
3.0
2.5
4.0
2.0
Problem Solving Center
Homework 4
1.
You are setting up a reaction between two chemicals that react
according to the equation 3 A + 4 B ---> products. If you start with
1.00 mole each of both A and B, which chemical will be in excess at
the end, and by how much (assuming the reaction goes to
completion)?
Neither A nor B is in excess, because the reaction "goes to
completion."
B is in excess by 0.250 mol.
B is in excess by 0.333 mol.
A is in excess by 0.333 mol.
A is in excess by 0.250 mol.
2.
What is the molecular formula of the following compound?
empirical formula CH, molar mass 78 g/mol
C4H4
C3H3
C6H6
CH
54
Unit 3/4 – Moles / Stoichiometry
C2H2
3.
Give the empirical formula of the following compound if a sample
contains 57.8 percent C, 3.6 percent H, and 38.6 percent O by mass.
C4H3O2
C12H9O6
C2HO
C8H6O4
4.
Based on the following structural formula, calculate the percentage of
carbon present.
(CH2CO)2C6H3(COOH)
67.37 percent
66.67 percent
76.73 percent
64.70 percent
5.
Calculate the mass percent of nitrogen in HNO3.
45.2 percent
22.2 percent
20.0 percent
25.0 percent
none of these
6.
Which of the following cannot be an empirical formula?
C3H6
NO2
CH
H2N
CO2
7.
A manufacturer of bicycles has 5350 wheels, 3023 frames, and 2655
handlebars. How many bicycles can be manufactured using these
parts?
55
Unit 3/4 – Moles / Stoichiometry
2675 bicycles
5350 bicycles
3023 bicycles
2655 bicycles
8.
CO2 exhaled by astronauts is removed from the spaceship
atmosphere by reaction with KOH:
CO2 + 2KOH
K2CO3 + H2O
How many kg of CO2 can be removed with 1.00 kg of KOH?
0.392 kg
0.784 kg
1.57 kg
0.500 kg
9.
Aluminum and oxygen react according to the following equation:
4Al(s) + 3O2(g)
2Al2O3(s)
In a certain experiment 4.6 g Al was reacted with excess oxygen and
6.8 g of product was obtained. What was the percent yield of the
reaction?
63 percent
78 percent
74 percent
134 percent
68 percent
10
.
For the reaction 3NO2 + H2O
2HNO3 + NO, how many grams
of HNO3 can form when 1.00 g of NO2 and 2.25 g of H2O are allowed
to react?
1.37 g
0.913 g
0.667 g
56
Unit 3/4 – Moles / Stoichiometry
15.7 g
11
.
For the reaction Fe(CO)5 + 2PF3 + H2
Fe(CO)2(PF3)2(H)2 +
3CO, how many moles of CO are produced from a mixture of 5.0 mol
Fe(CO)5, 8.0 mol PF3, and 6.0 mol H2?
18 mol
5.0 mol
6.0 mol
12 mol
15 mol
12
.
If 4.0 moles of Li and 2.0 moles of O2 are used in the reaction 4Li +
O2 ---> 2Li2O, then the limiting reactant is _________ and the
theoretical yield of Li2O is ____________ g.
oxygen, 6.0 x 10
oxygen, 1.2 x 102
lithium, 3.0 x 10
lithium, 6.0 x 10
Problem Solving Center
Quiz 1
1.
Write the balanced equation for the reaction that occurs when solid potassium nitrate is
heated and decomposes to form solid potassium nitrite and oxygen gas.
2KNO3
2.
2KNO2 + O2
2KNO4(s)
2KNO3(s) + O2(g)
2KNO3(s)
2KNO2(s) + O2(g)
KNO3(s)
KNO2(s) + (1/2)O2(g)
What is the formula weight of (NH4)2SO4?
118 amu
100 amu
116 amu
132 amu
3.
Balance the following equation and indicate whether it is a combustion, combination, or
decomposition reaction:
57
Unit 3/4 – Moles / Stoichiometry
Li + N2
Li3N
Li + N2
3Li3N, decomposition reaction
6Li + N2
2Li3N, decomposition reaction
6Li + N2
2Li3N, combination reaction
Li + N2
6Li + N2
4.
3Li3N, combination reaction
2Li3N, combustion reaction
Suppose you are setting up a reaction that requires an iodide salt and are planning to use
sodium iodide. However, at the last minute you find that you are out of sodium iodide, so
you must use potassium iodide instead. Will you need to weigh out more, less, or the
same mass of potassium iodide in order to get the same number of moles of iodide ions?
less
same
more
5.
Convert the following to a balanced chemical reaction:
Gaseous hydrogen reacts with carbon monoxide to form methanol, CH3OH.
4H + CO
CH3OH
H2 + CO
CH3OH
2H2 + CO2
CH3OH
58
Unit 3/4 – Moles / Stoichiometry
2H2 + CO
6.
Balance the following equation:
Mg3N2 + H2O
7.
CH3OH
Mg(OH)2 + NH3
Mg3N2 + 6H2O
3Mg(OH)2 + NH3
Mg3N2 + 6H2O
3Mg(OH)2 + 2NH3
Mg3N2 + 2H2O
Mg(OH)2 + NH3
Mg3N2 + 3H2O
3Mg(OH)2 + 2NH3
Calculate the number of molecules in a tablespoon of table sugar, C12H22O11, weighing
10.5 g.
2.22 x 1023
3.01 x 1023
1.85 x 1022
6.32 x 1024
6.02 x 1023
8.
Potassium sulfate contains 44.9 percent potassium by mass. In a 50.0-g sample of
potassium sulfate, the number of moles of potassium is
1.28 mol.
2.00 mol.
1.74 mol.
0.287 mol.
0.574 mol.
9.
The reaction C7H8 + 3HNO3 ---> C7H5N3O6 + 3H2O can be used to make TNT. How many
grams of HNO3 are required to react with 10.0 g of C7H8?
20.5 g
6.81 g
30.0 g
2.28 g
10.1 g
10
.
Calculate the number of moles of water present in a 10.0-kg sample.
1.80 x 102 mol
55.5 mol
555 mol
1.80 x 105 mol
none of these
11
.
How many F- ions are present in 2.50 mol of BaF2?
5.00
3.01 x 1024
59
Unit 3/4 – Moles / Stoichiometry
1.51 x 1024
2.50
8.31 x 10-24
12
.
What mass of silver chloride can be made from the reaction of 4.22 g of silver nitrate with
7.73 g of aluminum chloride? (Be sure to balance the reaction.)
AgNO3 + AlCl3
Al(NO3)3 + AgCl
24.9 g
3.56 g
10.7 g
12.7 g
Practice Questions
Write the balanced equations for the following reactions.
1. C2H6 + O2→ CO2 + H2O
2. Na + H2O → NaOH + H2
3. Ammonium nitrate decomposes to yield dinitrogen monoxide and water.
4. Ammonia reacts with oxygen gas to form nitrogen monoxide and water.
5. Iron (III) oxide reacts with carbon (C) to yield iron metal and carbon monoxide.
6. Hydrogen gas reacts with carbon monoxide to yield methanol (CH3OH). How many grams of methanol are formed
when 15.6 g of hydrogen react with excess carbon monoxide?
7. How many moles of carbon dioxide are formed in the fermentation of 75 g of glucose?
8. The thermite reaction (Fe2O3 + Al → Fe + Al2O3) can be used to ignite solid-fuel rockets or bombs. How much
aluminum is needed to react with 10.0 g of Fe2O3?
9. Identify the limiting reactant and how much ammonia gas can be produced when 7.2 g of nitrogen gas react with 1.5
g of hydrogen gas by the use of the Haber process: 3H2 + N2→ 2NH3.
10. Identify the limiting reactant and how much carbon dioxide gas can be produced when 15.2 g of methane react with
18.5 g of oxygen gas to produce water and carbon dioxide.
11. Identify the limiting reactant and how much nitric acid can be produced when 60.0 g of nitrogen dioxide react with
18.5 g of water to produce nitric acid and nitrogen monoxide.
60
Unit 3/4 – Moles / Stoichiometry
12. Identify the limiting reactant and how much aspirin (C9H8O4) can be produced when 52.3 g of salicylic acid
(C8H6O3) react with 25.0 g of acetic acid (CH3CO2H): C8H6O3 + CH3CO2H → C9H8O4 + H2O.
Calculate the percent yield for practice problems 9 through 12 if
13. 6.3 g of ammonia were produced from problem 9.
14. 12.4 g of carbon dioxide were produced from problem 10.
15. 51 g of nitric acid were produced from problem 11.
16. 31.0 g of aspirin were produced from problem 12.
Answers
1. 2C2H6 + 7O2 → 4CO2 + 6H2O
2. 2Na + 2H2O → 2NaOH + H2
3. NH4NO3 → N2O + 2H2O
4. 4NH3 + 5O2 → 4NO + 6H2O
5. Fe2O3 + 3C → 2Fe + 3CO
6. 124 g
7. 37 g
8. 3.38 g
9. Limiting reactant = H2; theoretical yield = 8.4 g
10. Limiting reactant = O2; theoretical yield = 12.7 g
11. Limiting reactant = NO2; theoretical yield = 55 g
12. Limiting reactant = salicylic acid; theoretical yield = 62.8 g
13. 75%
14. 97.6%
15. 93%
61
Unit 3/4 – Moles / Stoichiometry
16. 56%
-
Match the Following
62
Unit 3/4 – Moles / Stoichiometry
63
Unit 3/4 – Moles / Stoichiometry
64
Unit 3/4 – Moles / Stoichiometry
65
Unit 3/4 – Moles / Stoichiometry
66
Unit 3/4 – Moles / Stoichiometry
Chemistry: Chapter 9 Stoichiometry and
Baking Soda (NaHCO3) as the Limiting
Reactant Activity
Purposes:
1. Calculate theoretical mass of NaCl based on a known mass of NaHCO 3.
2. Experimentally determine the actual mass of NaCl produced.
3. Calculate the percent yield for your experiment.
Reaction Equation:
NaHCO3(s) + HCl(aq)  NaCl(s) + CO2(g) + H2O(l)
Materials:
safety glasses
HCl and dropper
evaporating dish
and matches
watch glass
baking soda (NaHCO3)
concentrated
ring stand with ring
bunsen burner
wire gauze
tongs
Procedure:
1. Find the mass of the evaporating dish and watch glass. Record this mass in the Data Table.
2. Add 1/3 of a teaspoon of baking soda to the evaporating dish, and record the total mass in the
Data Table.
3. Cover the evaporating dish with the watch glass so that only the spout of the evaporating dish
is exposed.
4. Use the dropper to drip HCl down the spout and into the dish. Add HCl until the fizzing
ceases.
5. Leaving the watch glass in place, boil off the liquid until only table salt (NaCl) remains in the
dish.
6. Let the dish cool for five minutes, then weigh it again and record the mass in the Data Table.
7. Clean up by rinsing your equipment with water and wiping dry with a paper towel.
Data Table: (please include units)
Quantity Measured
Mass
evaporating dish, watch glass
evaporating dish, watch glass, NaHCO3
evaporating dish, watch glass, NaCl
67
Unit 3/4 – Moles / Stoichiometry
Calculations:
1. Find the theoretical mass of NaCl that would be produced if your experiment were perfect.
2. Find the actual mass of NaCl that you obtained.
3. Find the percent yield for your experiment. If your percent yield is greater than 100%, provide
at least one possible source of error that might have caused you to get more than 100%
yield.
Chemistry: Ch.9 Complete Stoichiometry Review
1. How many moles of O2 should be supplied to
burn 1 mol of C3H8 (propane) molecules in a camping stove?
2. How many moles of O2 molecules should be
supplied to burn 1 mol of CH4 molecules in a domestic furnace?
3. Sodium thiosulfate (Na2S2O3), photographer’s
“hypo” reacts with unexposed silver bromide in the film emulsion to form sodium bromide and a
compound of formula Na5[Ag(S2O3) 3]. How many moles of Na2S2O3 formula units are needed to
make 0.10 mol of AgBr soluble?
4. Calculate the mass of alumina (Al2O3)
produced when 100 g of aluminum burns in oxygen.
5. “Slaked lime,” Ca(OH) 2, is formed from
“quick-lime” (CaO) by adding water. What mass of water is needed to convert 10 kg of quicklime
to slaked lime? What mass of slaked lime is produced?
6. Camels store the fat tristearin (C57H110O6) in
the hump. As well as being a source of energy, the fat is a source of water, because when it is
used the reaction
2 C57H110O6(s) + 163 O2(g) 
114 CO2(g) + 110 H2O(l)
takes place. What mass of water is available from 1.0 kg of fat?
7. The compound diborane (B2H6) was at one
time considered for use as a rocket fuel. How many grams of liquid oxygen would a rocket have
to carry to burn 10 kg of diborane completely? (The products of the combustion are B2O3 and
H2O.)
8. Given the balanced chemical equation
Br2 + 2 NaI  2 NaBr + I2
How many moles of sodium bromide (NaBr) could be produced from 0.172 mol of bromine (Br 2)?
68
Unit 3/4 – Moles / Stoichiometry
9. How many formula units of calcium oxide
(CaO) can be produced from 4.9 x 105 molecules of oxygen gas (O2) that react with calcium (Ca)
according to this balanced chemical equation?
2 Ca(s) + O2 (g)  2 CaO(s)
10. Aluminum metal (Al) reacts with sulfur (S) to
produce aluminum sulfide (Al2S3) according to this balanced chemical equation:
2 Al(s) + 3 S(s)  Al2S3(s)
How many atoms of aluminum will react completely with 1.33 x 10 24 atoms of sulfur?
Name _____________________________ Hr ___
LIMITING REAGENTS
11. What is the maximum mass of methane (CH 4)
that can be burned if only 1.0 g of oxygen is available?
12. What is the maximum mass of glucose
(C6H12O6) that can be burned in 10 g of oxygen?
13. The solid fuel in the booster stage of the space shuttle is a mixture of ammonium
perchlorate
and aluminum powder, which react as follows:
6 NH4ClO4(s) + 10 Al(s)  5 Al2O3(s) +
3 N2(g) + 6 HCl(g) + 9 H2O(g)
What mass of aluminum should be mixed with 5.0
x 103 kg of ammonium perchlorate, if the reaction
proceeds as stated?
14. A solution containing 5.0 g of silver nitrate was mixed with another containing 5.0 g of
potassium
chloride. Which was the limiting reagent for the
precipitation of silver chloride?
15. Given the balanced chemical equation
2 Ag + I2  2 AgI
How many atoms of silver metal (Ag) are required to react completely with 531.8 g of iodine (I 2) to
produce silver iodide (AgI)?
16. The theoretical yield of ammonia in an industrial synthesis was 550 tons, but only 480 tons
was
obtained. What was the percentage yield of the
reaction?
17. Calculate the volume occupied by 16.3 moles of
nitrogen gas (N2) at STP.
18. How many moles of fluorine gas (F2) are
69
Unit 3/4 – Moles / Stoichiometry
contained in 0.269 dm3 container at STP?
19. Assuming that the gases are all at STP, find the
volume of nitrogen dioxide gas (NO2) that could be produced from 71.11 dm 3 of nitrogen gas (N2)
according to this balanced chemical equation.
N2(g) + 2 O2(g)  2 NO2(g)
20. How many moles of oxygen (O2) would be
needed to produce 79.60 moles of sulfur trioxide (SO3) according to the following balanced
chemical equation?
2 SO2 + O2  2 SO3
21. How many grams of water will be produced from 50 g hydrogen reacting with 50 g
oxygen?
Think Critically
22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 
2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the the above information?
23. According to the balanced chemical equation;
how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver
nitrate?
Cu(s) + 2 AgNO3(aq)  Cu(NO3) 2(aq) + 2 Ag(s)
24. According to the balanced chemical equation;
how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with
1000 dm3 of oxygen gas? Assume a 75% yield.
2 ZnS(s) + 3 O2 (g)  2 ZnO(s) + 2 SO2(g)
25. I need to produce 500 g of lithium oxide(Li2O)
a) how many grams of Lithium AND
b) how many liters of oxygen do I need
The balanced equation is: Li + O2  LiO2
26. How many grams of water will be produce
from 50 g hydrogen reacting with 50 g oxygen?
Think Critically:
22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O2(g) 
2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the the above information?
23. According to the balanced chemical equation;
how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver
nitrate?
Cu(s) + 2 AgNO3(aq)  Cu(NO3) 2(aq) + 2 Ag(s)
70
Unit 3/4 – Moles / Stoichiometry
24. According to the balanced chemical equation;
how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with
1000 dm3 of oxygen gas? Assume a 75% yield.
2 ZnS(s) + 3 O2 (g)  2 ZnO(s) + 2 SO2(g)
25. I need to produce 500 g of lithium oxide(Li2O)
a. how many grams of Lithium AND
b. how many liters of oxygen do I need
The balanced equation is: Li + O2  LiO2
26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g
oxygen?
Think Critically:
22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 
2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the the above information?
23. According to the balanced chemical equation;
how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver
nitrate?
Cu(s) + 2 AgNO3(aq)  Cu(NO3) 2(aq) + 2 Ag(s)
24. According to the balanced chemical equation;
how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with
1000 dm3 of oxygen gas? Assume a 75% yield.
2 ZnS(s) + 3 O2 (g)  2 ZnO(s) + 2 SO2(g)
25. I need to produce 500 g of lithium oxide(Li2O)
a. how many grams of Lithium AND
b. how many liters of oxygen do I need
The balanced equation is: Li + O2  LiO2
26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g
oxygen?
Think Critically:
22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 
2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the the above information?
23. According to the balanced chemical equation;
how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver
nitrate?
Cu(s) + 2 AgNO3(aq)  Cu(NO3) 2(aq) + 2 Ag(s)
71
Unit 3/4 – Moles / Stoichiometry
24. According to the balanced chemical equation;
how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with
1000 dm3 of oxygen gas? Assume a 75% yield.
2 ZnS(s) + 3 O2 (g)  2 ZnO(s) + 2 SO2(g)
25. I need to produce 500 g of lithium oxide(Li2O)
a. how many grams of Lithium AND
b. how many liters of oxygen do I need
The balanced equation is: Li + O2  LiO2
26. How many grams of water will be produced from 50 g hydrogen reacting with 50 g
oxygen?
Think Critically
22. The reaction of 1 mol of C to form carbon monoxide in the reaction 2 C(s) + O 2(g) 
2 CO(g) releases 113 kJ of heat. How much heat will be released by the combustion of 100 g of
C according the the above information?
23. According to the balanced chemical equation;
how many atoms of silver will be produced from combining 100 g of copper with 200 g of silver
nitrate?
Cu(s) + 2 AgNO3(aq)  Cu(NO3) 2(aq) + 2 Ag(s)
24. According to the balanced chemical equation;
how many moles of SO2(g) will be produced when 1.5 x 108 molecules of zinc sulfide react with
1000 dm3 of oxygen gas? Assume a 75% yield.
2 ZnS(s) + 3 O2 (g)  2 ZnO(s) + 2 SO2(g)
25. I need to produce 500 g of lithium oxide(Li2O)
a) how many grams of Lithium AND
b) how many liters of oxygen do I need
The balanced equation is: Li + O2  LiO2
26. A tin ore contains 3.5% SnO2. How much tin is produced by reducing 2.0 kg of the ore with
carbon?
SnO2 + C  Sn + CO2
27. If 36.5 g of HCl and 73 g of Zn are put together:
2 HCl + Zn  ZnCl2 + H2
a.
b.
c.
d.
Determine which reactant is the limiting reactant,
Find the mass of ZnCl2 formed,
Find the volume of H2 (@ STP) formed,
Determine which reactant is in excess and by how much.
28. Many plants synthesize glucose by photosynthesis as follows:
CO2(g) + H2O(l) + energy  C6H12O6(s) + O2(g)
a. Write a balanced equation for this process,
72
Unit 3/4 – Moles / Stoichiometry
b. How many molecules of water are needed to make one molecule of glucose?
c. How many liters of oxygen (@STP) are given off when 2.50 mol of glucose is
synthesized?
d. How many moles of CO2 are needed for a plant to make 2.50 mole of glucose?
e. How many carbon atoms are used to produce 2.50 mole of glucose?
f. How many dm3 of oxygen gas are produced from 9.32 dm 3 of CO2 (all @ STP)?
29. Assume that the human body requires daily energy that comes from metabolizing 816 g
of sucrose, C12H22O11, using the following reaction:
C12H22O11(s) + 12 O2(g)  12 CO2(g) +
11 H2O(l) + energy
How many dm3 of pure oxygen (@ STP) is consumed by a human being in 24 hours?
30. A student has a mixture of KClO3, K2CO3, and KCl. She heats 50 g of the mixture and
determines that 5 g O2 and 7 g CO2 are produced by these reactions:
2 KClO3(s)  2 KCl(s) + 3 O2(g)
K2CO3(s)  K2O(s) + CO2(g)
KCl is not affected by the heat. What is the percent composition of the original mixture?
ANSWERS:
1. 5 mol O2
2. 2 mol O2
3. 0.3 mol Na2S2O3
4. 189 g Al2O3
5. 3214 g H2O and 13.2 kg slaked lime [Ca(OH) 2]
6. 998 g water
7. 34,783 g O2
8. 0.344 mol NaBr
9. 9.8 x 105 molecules CaO
10. 8.9 x 1023 atoms Al
11. 0.25 g CH4
12. 9.375 g C6H12O6
13. 1915 kg Al
14. silver nitrate
15. 2.5 x 1024 atoms Ag
16. 87.3 % yield
17. 365 L N2
18. 0.012 mol F2
19. 142 L NO2
20. 39.8 mol O2
21. 56.25 g H2O
22. 942,000 J
23. 7.1 x 1023 atoms Ag
24. 1.9 x 10-16mol (NOT 2.5 x 10-16mol: 75% Yield)
25a.
g Li
b.
L O2
26.
27a.
b.
c.
d.
28a.
b. 6 c. 336 d. 15 e. 9 x 1024 f. 9.32
29. 641 L O2
30. 15.2 g KCl
73
Unit 3/4 – Moles / Stoichiometry
Name:
________________________
Hour:
____
Date:
___________
Chemistry: Percent Yield
Directions: Solve each of the following problems. Show your work, including proper units, to
earn full credit.
1. “Slaked lime,” Ca(OH)2, is produced when water reacts with “quick lime,” CaO. If you start
with 2 400 g of quick lime, add excess water, and produce 2 060 g of slaked lime, what is
the percent yield of the reaction?
2. Some underwater welding is done via the thermite reaction, in which rust (Fe 2O3) reacts with
aluminum to produce iron and aluminum oxide (Al2O3). In one such reaction, 258 g of
aluminum and excess rust produced 464 g of iron. What was the percent yield of the
reaction?
3. Use the balanced equation to find out how many liters of sulfur dioxide are actually produced
at STP if
1.5 x 1027 molecules of zinc sulfide are reacted with excess oxygen and the
percent yield is 75%.
2 ZnS(s) + 3 O2(g)  2 ZnO(s) + 2 SO2(g)
4. The Haber process is the conversion of nitrogen and hydrogen at high pressure into ammonia,
as follows:
N2(g) + 3 H2(g)  2 NH3(g)
74
Unit 3/4 – Moles / Stoichiometry
If you must produce 700 g of ammonia, what mass of nitrogen should you use in the
reaction, assuming that the percent yield of this reaction is 70%?
Answers:
1. 65%
4. 824 g N2
2. 87%
3. 4.19 x 104 L SO2
75
Unit 3/4 – Moles / Stoichiometry
76
Unit 3/4 – Moles / Stoichiometry
77
Unit 3/4 – Moles / Stoichiometry
78
Unit 3/4 – Moles / Stoichiometry
79
Unit 3/4 – Moles / Stoichiometry
80
Unit 3/4 – Moles / Stoichiometry
81
Unit 3/4 – Moles / Stoichiometry
82
Unit 3/4 – Moles / Stoichiometry
83
Unit 3/4 – Moles / Stoichiometry
84
Unit 3/4 – Moles / Stoichiometry
85
Unit 3/4 – Moles / Stoichiometry
86
Unit 3/4 – Moles / Stoichiometry
CHEMICAL REACTIONS OF COPPER AND PERCENT YIELD
Objective
To gain familiarity with basic laboratory procedures, some chemistry of a typical transition
element, and the concept of percent yield.
Apparatus and Chemicals
0.5 g piece of no. 16 or no. 18 copper wire
evaporating dish
250 mL beaker (2)
weighing paper
concentrated HNO3 (4 – 6 mL)
6.0 M H2SO4 (15 mL)
graduated cylinder
granular zinc
3.0 M NaOH (30 mL)
methanol
carborundum boiling chips
acetone
stirring rod
towel
iron ring and ring stand
balance
wire gauze
aluminum foil cut in 1-inch squares
Bunsen burner
concentrated HCl (drops)
Discussion
Most chemical synthesis involves separation and purification of the desired product from
unwanted side products. Some methods of separation, such as filtration, sedimentation,
decantation, extraction, and sublimation were discussed earlier. This experiment is designed as
a quantitative evaluation of your individual laboratory skills in carrying out some of these
operations. At the same time you will become more acquainted with two fundamental types of
chemical reactions -- redox reactions and metathesis (double-displacement) reactions. By means
of these reactions, you will finally recover the copper sample with maximum efficiency. The
chemical reactions involved are the following.
Cu(s) + 4 HNO3(aq) ----->
Redox
[1]
Cu(NO3)2(aq) + 2 NaOH(aq)
Metathesis
[2]
Cu(OH)2(s) ----->
Dehydration [3]
CuO(s) +
Metathesis
CuO(s)
H2SO4(aq)
[4]
Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l)
----->
+
----->
Cu(OH)2(s)
+
2 NaNO3(aq)
H2O(g)
CuSO4(aq)
+
H2O(l)
87
Unit 3/4 – Moles / Stoichiometry
CuSO4(aq)
Redox
+ Zn(s)
[5]
----->
ZnSO4(aq)
+
Cu(s)
Each of these reactions proceeds to completion. Metathesis reactions proceed to completion
whenever one of the components is removed from the solution, such as in the formation of a gas
or an insoluble precipitate (driving forces). This is the case for reaction [1], [2], and [3], where in
reactions [1] and [3] a gas and in reaction [2] an insoluble precipitate are formed. Reaction [5]
proceeds to completion because zinc has a lower ionization energy or oxidation potential that
copper.
The objective in this experiment is to recover all of the copper you begin with, in analytically
pure form.
This is the test of your laboratory skills.
The percent yield of the copper can be expressed as the ratio of the recovered weight to initial
weight, multiplied by 100:
recovered wt of Cu
% yield =
x 100
initial wt of Co
Procedure

Weight approximately 0.500 g of no. 16 or no. 18 copper wire (1) to the nearest 0.0001 g
and place it in a 250 mL beaker. Add 4-5 mL of concentrated HNO3 to the beaker, IN
THE HOOD. After the reaction is complete, add 100 mL distilled H2O. Describe the
reaction (6) as to color change, evolution of gas, and change in temperature (exothermic
or endothermic) in the report sheet.

Add 30 mL of 3.0 M NaOH to the solution in your beaker and describe the reaction (7).
Add two or three boiling chips and carefully heat the solution -- while stirring with a glass
stirring rod -- just to the boiling point. Describe the reaction on your report sheet (8).
Remove the boiling chips.

Allow the black CuO to settle; then decant the supernantant liquid. Add about 200 mL of
very hot distilled water and allow the CuO to settle. Decant once more. What are you
removing by washing and decanting (9)?

Add 15 mL of 6.0 M H2SO4. What copper compound is present in the beaker now (10)?
Your teacher will tell you whether you should use Zn or Al for the reduction of Cu (II) in the
following step.
A. Zinc
In the hood, add 2.0 g of 30-mesh zinc metal all at once and stir until the supernatant liquid is
colorless. Describe the reaction on your report sheet (11). What is present in solution (12)?
When gas evolution has become very slow, heat the solution gently (but do not boil) and allow it
to cool. What gas is formed in this reaction (13)? How do you know (14)?
B. Aluminum
88
Unit 3/4 – Moles / Stoichiometry
In the hood, add several 1-inch squares of aluminum foil and a few drops of concentrated HCl.
Continue to add pieces of aluminum until the supernatant liquid is colorless. Describe the
reaction on your report sheet (11). What is present in solution (12)? What gas is formed in this
reaction (13)? How do you know (14)?
When gas evolution has ceased, decant the solution and transfer the precipitate to a
preweighed porcelain evaporating dish (3). Wash the precipitated copper with about 5 mL of
distilled water, allow it to settle, decant the solution, and repeat the process. What are you
removing by washing (15)? Wash the precipitate with about 5 mL of methanol (KEEP THE
METHANOL AWAY FROM FLAMES _ IT IS FLAMMABLE!) Allow the precipitate to settle, and
decant the methanol. (METHANOL IS ALSO EXTREMELY TOXIC: AVOID BREATHING THE
VAPORS AS MUCH AS POSSIBLE.) Finally, wash the precipitate with about 5 mL of acetone
(KEEP THE ACETONE AWAY FROM FLAMES - IT IS EXTREMELY FLAMMABLE!), allow the
precipitate to settle, and decant the acetone from the precipitate. Prepare a steam bath as
illustrated and dry the product on your steam bath for at least 5 minutes.
Wipe the bottom of the evaporating dish with a towel, remove the boiling chips and weigh the
evaporating dish plus copper (2). Calculate the final a\weight of copper (4). Compare the weight
with your initial weight and calculate the percent yield (5). What color is your copper sample
(16)? Is it uniform in appearance (17)? Suggest possible sources of error in this experiment (18).
Chemical Reactions of Copper and Percent Yield
Pre-lab (Review Questions)
1. Give an example, other than the ones listed in this experiment, of redox and metathesis
reactions.
2. When will reactions proceed to completion?
3. Define percent yield in general terms.
4. Name six methods of separating materials.
89
Unit 3/4 – Moles / Stoichiometry
5. Give criteria in terms of temperature changes for exothermic and endothermic reactions.
6. If 1.65 g of Cu(NO3)2 are obtained from allowing 0.93 g of Cu to react with excess HNO 3, what
is the percent
yield of the reaction?
Prelab – Chemical Reactions of Copper and Percent Yield
7. What is the maximum percent yield in any reaction?
8. What is meant by the terms decantation and filtration?
9. When Cu(OH)2(s) is heated, Copper (II) oxide and water are formed.
Write a balanced equation for the reaction.
90
Unit 3/4 – Moles / Stoichiometry
10. When sulfuric acid and copper (II) oxide are allowed to react, copper (II) sulfate and water
are formed.
Write a balanced equation for this reaction.
11. When copper (II) sulfate and aluminum are allowed to react, aluminum sulfate and copper
are formed.
What kind of reaction is this? Write a balanced equation for this reaction.
REPORT SHEET
Chemical Reactions of Copper and Percent Yield
1. Weight copper initial
_______________
2. Weight of copper and evaporating dish
_______________
3. Weight of evaporating dish
_______________
4. Weight of copper final
_______________
5. % Yield (show calculations)
_______________
6. Describe the reaction Cu(s) + HNO3(aq) -->
7. Describe the reaction Cu(NO3)2(aq) + NaOH(aq) -->
8. Describe the reaction Cu(OH)2(s) -->
9. What are you removing by this washing (be specific)?
91
Unit 3/4 – Moles / Stoichiometry
10. What copper compound is present in the beaker?
11. Describe the reaction CuSO4(aq) + Zn(s), or CuSO4(aq) + Al(s)
12. What is present in solution (aqueous)?
REPORT SHEET – page 2
Chemical Reactions of Copper and Percent Yield
13. What is the gas? Hint: Where did the HCl(aq) go? Why did you have to do this reaction in the
fume hood?
14. How do you know?
15. What are you removing by washing?
16. What color is your copper sample?
17. Is it uniform in appearance?
18. Suggest possible sources of error in this experiment.
POST LAB QUESTIONS
1. If your percent yield of copper was greater than 100%, what are two plausible errors you may
have made?
92
Unit 3/4 – Moles / Stoichiometry
2. Consider the combustion of methane, CH4:
CH4(g) + 2 O2(g) ----->
CO2(g) + 2 H2O(g)
Suppose 2 mole of methane is allowed to react with 3 mol of oxygen.
a) What is the limiting reagent? (show work)
b) How many moles of CO2 can be made from this mixture? How many grams of CO2?
3. Suppose 8.00 g of CH4 is allowed to burn in the presence of 6.00 g of oxygen.
How much (in grams) CH4, O2, CO2, and H2O remain after the reaction is complete?
4. How many milliliters of 6.0 M H2SO4 are required to react with 0.80 g of CuO according to
Equation [4]?
5. If 2.00 g of Zn is allowed to react with 1.75 g of CuSO 4 according to Equation [5], how many
grams of Zn will
remain after the reaction is complete?
93
Unit 3/4 – Moles / Stoichiometry
6. What is meant by the term limiting reagent? Explain
94
Unit 3/4 – Moles / Stoichiometry
95
Unit 3/4 – Moles / Stoichiometry
96
Unit 3/4 – Moles / Stoichiometry
97
Unit 3/4 – Moles / Stoichiometry
98
Unit 3/4 – Moles / Stoichiometry
99
Unit 3/4 – Moles / Stoichiometry
Show your work for all calculations.
1.
What volume of carbon dioxide can be produced by reacting 10.5 g baking soda (NaHCO 3) with
excess hydrochloric acid at STP according to the reaction below?
NaHCO3(s) + HCl(aq) → NaCl(aq) + CO2(g) + H2O(l)
2.
Potassium chlorate decomposes to produce oxygen gas by the reaction shown below. What
volume of oxygen gas can be produced by heating 5.89 g KClO 3 at a pressure of 0.50 atmospheres
and a temperature of 79 °C?
2 KClO3(s) → 2KCl(s) + 3O2(g)
3.
Cetyl alcohol (C16H34O) is a flammable solid at room temperature. If 15.8 g of the alcohol are
combusted, what volume of water vapor will be produced at a pressure of 1.26 atmospheres and a
temperature of 110 °C? (Watch out — this equation isn't balanced!)
C16H34O(s) + O2(g) → CO2(g) + H2O(g)
Conclusion
Now, address the Focus Question: How much metal does it take to float a blimp? Given the information
below, you should be able to calculate an answer.
100
Unit 3/4 – Moles / Stoichiometry
Zinc reacts with sulfuric acid by the reaction shown below. If 759 kg zinc were to react with excess sulfuric
acid, how large would a blimp be if it were filled with the hydrogen gas produced from this reaction at a
pressure of 1.10 atm and a temperature of 22 °C?
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
Compound
Table 2
Formula Mass (amu)
MgO
K2S
C3H8O3
Ca(NO3)2
Mg3(PO4)2
C6H2CH3(NO2)3
Pt(NH3)2Cl2
Fe(ClO4)2
Na2CO3 • 10H2O
MgSO4 • 7H2O
Conclusion
Now address the Focus Question: Which has more mass—salt (NaCl) or sugar (C12H22O11)? Show
calculations to support your answer.
A comparison of two domestic fuels - pollution
This sheet compares the amount of pollution produced by two domestic fuels - anthracite,
the
purest form of coal, and propane, sold in cylinders as liquefied petroleum gas, LPG.
Both anthracite and propane produce carbon dioxide when they burn.
1. What environmental problem does carbon dioxide contribute to?
..........................................................................................................................................
2. Give two practical consequences of this problem.
..........................................................................................................................................
101
Unit 3/4 – Moles / Stoichiometry
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
Anthracite contains between 89% and 98% pure carbon. When carbon burns,
carbon dioxide is
formed and heat energy is produced.
C(s) + O2(g) ’ CO2(g) DH = -394 kJ mol-1
3. Calculate the mass of 1 mole of carbon dioxide. (RAMs: C = 12; O = 16)
..........................................................................................................................................
4. Calculate the mass of carbon dioxide formed per kJ of heat energy produced.
..........................................................................................................................................
..........................................................................................................................................
Propane burns according to the equation:
C3H8(g) + 5O2(g) ’ 3CO2(g) + 4H2O(l) DH = -2220 kJ mol-1
5. Calculate the mass of carbon dioxide formed per kJ of heat energy produced
when propane
burns.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
6. Which is the more environmentally friendly fuel in terms of carbon dioxide
emission?
..........................................................................................................................................
Sulphur dioxide production
Sulphur dioxide is formed when sulphur and sulphur-containing compounds burn.
7. What environmental problem does sulphur dioxide contribute to?
102
Unit 3/4 – Moles / Stoichiometry
..........................................................................................................................................
8. Give two practical consequences of this problem.
..........................................................................................................................................
..........................................................................................................................................
..........................................................................................................................................
Coal often contains sulphur compounds, and these are converted into sulphur
dioxide when the
coal burns. At the moment there is no way of removing sulphur dioxide from the
gases
produced by homes. Propane is freed of sulphur-containing compounds at the
refinery. The sulphur compounds are converted into sulphur dioxide which can
then be used to manufacture a useful chemical.
9. Which chemical?
..........................................................................................................................................
10. Which is the more environmentally friendly fuel in terms of sulphur dioxide
emission?
..........................................................................................................................................
103
Unit 3/4 – Moles / Stoichiometry
1. Copper reacts with nitric acid according to the following reaction:
3 Cu (s) + 8 HNO3 (aq)  3 Cu(NO3)2 (aq) + 2 NO (g) + 4 H2O (l)
If an pre 1982 copper penny contains 3.10 grams of copper, what volume of 8.00 M
nitric acid is required to exactly consume it? What volume of nitrogen monoxide gas
measured at STP would be produced? (Remember that 1 mole of any gas measured at
STP = 22.4 dm3)
2. Hydrogen peroxide decomposes slowly to produce oxygen gas and water. Write
a balanced equation for the decomposition of hydrogen peroxide. What mass
of hydrogen peroxide (H2O2) must decompose to produce 0.77g of water?
3. Carbon monoxide is an air pollutant found in automobile exhausts. It slowly
reacts with oxygen in the atmosphere to produce carbon dioxide. Write a
balanced equation for this process. What mass of carbon dioxide would be
produced from 10.00 grams of carbon monoxide? What mass of oxygen
is required to convert the carbon monoxide to carbon dioxide?
4. Suppose that 50.0 cm3 of 2.00 M hydrochloric acid are added to 4.00 grams of
zinc metal.
What it the limiting reagent? How many grams of zinc are actually consumed? What
is the concentration of the hydrochloric acid solution after the reaction has occurred?
What mass of hydrogen will be produced?
Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)
104
Unit 3/4 – Moles / Stoichiometry
5. Consider the reaction: 2 VO + 3 Fe2O3  6 FeO + V2O5
If 609.5 grams of vanadium (II) oxide, VO, and 832 grams of iron(III) oxide,
Fe2O3, are put into a container and allowed to react according to the equation
above, which substance(s) and how many grams of each would be present in the
container after the reaction is complete?
6. Methyl alcohol (wood alcohol), CH3OH, is produced via the reaction
CO(g) + 2 H2(g)  CH3OH(l)
A mixture of 1.20 g H2(g) and 7.45 g CO(g) are allowed to react.
(a) Which reagent is the limiting reagent?
(b) What is the yield of CH3OH? [Assume theoretical yield in g is what is
wanted here.]
(c) How much of the reagent present in excess is left over?
(d) Suppose the actual yield is 7.52 g of CH3OH. What is the % yield?
7. A 0.32570 gram sample of NaClO3 is decomposed to form oxygen gas according to the following
reaction:
2 NaClO3 (s)  2 NaCl (s) + 3 O2 (g)
The gas is collected over water at a total pressure of 742 torr at 23.0 oC. What volume of oxygen will
be collected? (The vapor pressure of water at 23.0 oC is 20.65 torr)
8. Hydrogen is produced when zinc reacts with sulfuric acid according to the following reaction
Zn (s) + H2SO4 (aq)  ZnSO4 (aq) + H2 (g)
If 159 cm3 of wet hydrogen is collected over water at 23.0o C and a total pressure of 738 torr, how many
grams of zinc are consumed? (The vapor pressure of water at 23.0 oC is 20.65 torr)
105
Unit 3/4 – Moles / Stoichiometry
Stoichiometry Challenge
1. (pg 129, #118) Nitric acid is produced commercially by the Ostwald process, represented by the following
equations:
a. 4HN3(g) +5O2(g)  4NO(g) + 6H2O(g)
b. 2NO(g) + O2(g)  2NO2(g)
c. 3NO2(g) + H2O(l)  2HNO3(aq) + NO(g)
What mass of NH3 must be used to produce 1.o x 106 kg HNO3 by the Ostwald process? Assume 100% yield in
each reaction and assume that the NO produced in the third step is not recycled.
2. (pg 129, #120) The aspirin substitute, acetaminophen (C8H9O2N), is produced by the following three-step
synthesis:
a. C6H5O3N(s) + 3H2(g) + HCl(aq)  C6H8ONCl(s) + 2H2O(l)
b. C6H8ONCl(s) + NaOH(aq)  C6H7ON(s) + H2O(l) + NaCl(aq)
c. C6H7ON(s) + C4H6O3(l)  C8H9O2N(s) + HC2H3O2(l)
The first two reactions have percent yields of 87% and 98% by mass, respectively. The overall reaction yields 3
mol of acetaminophen product for every 4 mol CH5O3N reacted.
a. What is the percent yield by mass for the overall process?
b. What is the percent yield by mass of step III?
- 106 -