Page 1 of 5
Data/Observations
1. What is the mass of copper oxide that reacted?
2. What is the mass of copper produced?
Analysis and Conclusions
4. Compare your result from question 2 to the mass percent of copper in copper(II) oxide by finding the difference between the two. Note: You calculated the
mass percent of copper in copper(II) oxide in the prelab
assignment.
1. How could you tell when the reaction was completed?
5. What is the formula of the copper oxide you reacted?
Justify your answer.
2. Calculate the percent copper by mass in your original
copper oxide.
6. Which would have a larger mass, an iron bar before or
after it rusts? Explain your answer.
3. Compare your result from question 2 to the mass percent of copper in copper(I) oxide by finding the difference between the two. Note: You calculated the mass
percent of copper in copper(I) oxide in the prelab
assignment.
Chapter Review
Key Terms
atomic mass unit (6.2)
average atomic mass (6.2)
mole (6.3)
Avogadro’s number (6.3)
Something Extra
Why was the gas sent through the test tube? Could this
lab be performed with a test tube open to the air? Explain your answer.
ger multiple of the empirical formula. The following
diagram summarizes these different ways of expressing
the same information.
actual masses
0.0806 g C
0.01353 g H
0.1074 g O
molar mass (6.4)
mass percent (6.5)
empirical formula (6.6)
molecular formula (6.6)
Summary
1. We can count individual units by weighing if we know
the average mass of the units. Thus, when we know
the average mass of the atoms of an element as that
element occurs in nature, we can calculate the number
of atoms in any given sample of that element by weighing the sample.
2. A mole is a unit of measure equal to 6.022 1023,
which is called Avogadro’s number. One mole of any
substance contains 6.022 1023 units.
3. One mole of an element has a mass equal to the element’s atomic mass expressed in grams. The molar
mass of any compound is the mass (in grams) of 1 mol
of the compound and is the sum of the masses of the
component atoms.
4. Percent composition consists of the mass percent of
each element in a compound:
Mass percent mass of a given element in 1 mol of compound
100%
mass of 1 mol of compound
5. The empirical formula of a compound is the simplest
whole-number ratio of the atoms present in the compound; it can be derived from the percent composition
of the compound. The molecular formula is the exact
formula of the molecules present; it is always an inte-
empirical
formula
CH2O
% composition
39.99% C
6.71% H
53.29% O
molar mass
molecular
formula
(CH2O)n
Questions and Problems*
All exercises with blue numbers have answers in the back
of this book.
6.1 Counting by Weighing
Problems
1. Merchants usually sell small nuts, washers, and bolts
by weight (like jelly beans!) rather than by individually counting the items. Suppose a particular type of
washer weighs 0.110 g on the average. What would
100 such washers weigh? How many washers would
there be in 100. g of washers?
*The element symbols and formulas are given in some
problems but not in others to help you learn this necessary
“vocabulary.”
Chapter Review
187
Page 2 of 5
2. A particular small laboratory cork weighs 1.63 g,
whereas a rubber lab stopper of the same size weighs
4.31 g. How many corks would there be in 500. g of
such corks? How many rubber stoppers would there
be in 500. g of similar stoppers? How many grams of
rubber stoppers would be needed to contain the same
number of stoppers as there are corks in 1.00 kg of
corks?
6.2 Atomic Masses: Counting Atoms
by Weighing
Questions
3. Define the amu. What is one amu equivalent to in
grams?
4. Why do we use the average atomic mass of the elements when performing calculations?
Problems
5. Using average atomic masses for each of the following elements (see the table inside the front cover of
this book), calculate the mass, in amu, of each of the
following samples.
a. 635 atoms of hydrogen
b. 1.261 104 atoms of tungsten
c. 42 atoms of potassium
d. 7.213 1023 atoms of nitrogen
e. 891 atoms of iron
6. Using average atomic masses for each of the following elements (see the table inside the front cover of
this book), calculate the number of atoms present in
each of the following samples.
a. 10.81 amu of boron
b. 320.7 amu of sulfur
c. 19,697 amu of gold
d. 19,695 amu of xenon
e. 3588.3 amu of aluminum
7. If an average atom of sulfur weighs 32.07 amu, how
many sulfur atoms are contained in a sample with
mass 8274 amu? What is the mass of 5.213 1024 sulfur atoms?
12. Calculate the average mass in grams of 1 atom of
oxygen.
13. Which weighs more, 0.50 mol of oxygen atoms or 4
mol of hydrogen atoms?
14. Using the average atomic masses given inside the front
cover of the text, calculate the number of moles of
each element in samples with the following masses.
a. 26.2 g of gold
b. 41.5 g of calcium
c. 335 mg of barium
d. 1.42 103 g of palladium
e. 3.05 105 g of nickel
f. 1.00 lb of iron
g. 12.01 g of carbon
15. Using the average atomic masses given inside the front
cover of the text, calculate the mass in grams of each
of the following samples.
a. 2.00 mol of iron
b. 0.521 mol of nickel
c. 1.23 103 mol of platinum
d. 72.5 mol of lead
e. 0.00102 mol of magnesium
f. 4.87 103 mol of aluminum
g. 211.5 mol of lithium
h. 1.72 106 mol of sodium
16. Using the average atomic masses given inside the front
cover of the text, calculate the indicated quantities.
a. the number of cobalt atoms in 0.00103 g of cobalt
b. the number of cobalt atoms in 0.00103 mol of
cobalt
c. the number of moles of cobalt in 2.75 g of cobalt
d. the number of moles of cobalt represented by
5.99 1021 cobalt atoms
e. the mass of 4.23 mol of cobalt
f. the number of cobalt atoms in 4.23 mol of cobalt
g. the number of cobalt atoms in 4.23 g of cobalt
6.4 Molar Mass
Questions
6.3 The Mole
17. The
of a substance is the mass (in grams) of
1 mol of the substance.
Questions
8. In 24.02 g of carbon, there are
carbon atoms.
9. A sample equal to the atomic mass of an element in
grams contains
atoms.
Problems
10. What mass of calcium metal contains the same number of atoms as 12.16 g of magnesium? What mass of
calcium metal contains the same number of atoms as
24.31 g of magnesium?
11. What mass of cobalt contains the same number of
atoms as 57.0 g of fluorine?
188
Chapter 6 Chemical Composition
18. The molar mass of a substance can be obtained by
the atomic weights of the component atoms.
Problems
19. Calculate the molar mass for each of the following
substances.
a. sodium nitride, Na3N
b. carbon disulfide, CS2
c. ammonium bromide, NH4Br
d. ethyl alcohol, C2H5OH
e. sulfurous acid, H2SO3
f. sulfuric acid, H2SO4
Page 3 of 5
20. Calculate the molar mass for each of the following
substances.
a. barium perchlorate
d. copper(II) nitrate
b. magnesium sulfate
e. tin(IV) chloride
c. lead(II) chloride
f. phenol, C6H6O
21. Calculate the number of moles of the indicated substance in each of the following samples.
a. 49.2 mg of sulfur trioxide
b. 7.44 104 kg of lead(IV) oxide
c. 59.1 g of chloroform, CHCl3
d. 3.27 g of trichloroethane, C2H3Cl3
e. 4.01 g of lithium hydroxide
22. Calculate the number of moles of the indicated substance in each of the following samples.
a. 4.26 103 g of sodium dihydrogen phosphate
b. 521 g of copper(I) chloride
c. 151 kg of iron
d. 8.76 g of strontium fluoride
e. 1.26 104 g of aluminum
23. Calculate the mass in grams of each of the following
samples.
a. 1.50 mol of aluminum iodide
b. 1.91 103 mol of benzene, C6H6
c. 4.00 mol of glucose, C6H12O6
d. 4.56 105 mol of ethanol, C2H5OH
e. 2.27 mol of calcium nitrate
24. Calculate the mass in grams of each of the following
samples.
a. 1.27 103 mol of carbon dioxide
b. 4.12 103 mol of nitrogen trichloride
c. 0.00451 mol of ammonium nitrate
d. 18.0 mol of water
e. 62.7 mol of copper(II) sulfate
25. Calculate the number of molecules present in each of
the following samples.
a. 6.37 mol of carbon monoxide
b. 6.37 g of carbon monoxide
c. 2.62 106 g of water
d. 2.62 106 mol of water
e. 5.23 g of benzene, C6H6
26. Calculate the number of moles of sulfur atoms present in each of the following samples.
a. 2.01 g of sodium sulfate
b. 2.01 g of sodium sulfite
c. 2.01 g of sodium sulfide
d. 2.01 g of sodium thiosulfate, Na2S2O3
6.5 Percent Composition of Compounds
Problems
27. Calculate the percent by mass of each element in the
following compounds.
a. sodium sulfate
b. sodium sulfite
c. sodium sulfide
d.
e.
f.
g.
h.
sodium thiosulfate, Na2S2O3
potassium phosphate
potassium hydrogen phosphate
potassium dihydrogen phosphate
potassium phosphide
28. Calculate the percent by mass of the element listed
first in the formulas for each of the following
compounds.
a. copper(II) bromide, CuBr2
b. copper(I) bromide, CuBr
c. iron(II) chloride, FeCl2
d. iron(III) chloride, FeCl3
e. cobalt(II) iodide, CoI2
f. cobalt(III) iodide, CoI3
g. tin(II) oxide, SnO
h. tin(IV) oxide, SnO2
29. Calculate the percent by mass of the element listed
first in the formulas for each of the following
compounds.
a. adipic acid, C6H10O4
b. ammonium nitrate, NH4NO3
c. caffeine, C8H10N4O2
d. chlorine dioxide, ClO2
e. cyclohexanol, C6H11OH
f. dextrose, C6H12O6
g. eicosane, C20H42
h. ethanol, C2H5OH
30. For each of the following ionic substances, calculate
the percentage of the overall molar mass of the compound that is represented by the positive ions the compound contains.
a. ammonium chloride
c. gold(III) chloride
b. copper(II) sulfate
d. silver nitrate
6.6 Formulas of Compounds
Questions
31. What experimental evidence about a new compound
must be known before its formula can be determined?
32. What does the empirical formula of a compound represent? How does the molecular formula differ from
the empirical formula?
33. Give the empirical formula that corresponds to each
of the following molecular formulas.
a. sodium peroxide, Na2O2
b. terephthalic acid, C8H6O4
c. phenobarbital, C12H12N2O3
d. 1,4-dichloro-2-butene, C4H6Cl2
34. Which of the following pairs of compounds have the
same empirical formula?
a. acetylene, C2H2, and benzene, C6H6
b. ethane, C2H6, and butane, C4H10
c. nitrogen dioxide, NO2, and dinitrogen tetroxide,
N2O4
d. diphenyl ether, C12H10O, and phenol, C6H5OH
Chapter Review
189
Page 4 of 5
6.7 Calculation of Empirical Formulas
Problems
Problems
46. A compound with the empirical formula CH2O was
found to have a molar mass between 89 and 91 g.
What is the molecular formula of the compound?
35. A new compound has been prepared. A 0.4791-g
sample was analyzed and was found to contain the
following masses of elements: carbon, 0.1929 g;
hydrogen, 0.01079 g; oxygen, 0.08566 g; chlorine,
0.1898 g. Determine the empirical formula of the new
compound.
36. In an experiment, a 2.514-g sample of calcium was
heated in a stream of pure oxygen, and was found to
increase in mass by 1.004 g. Calculate the empirical
formula of calcium oxide.
37. A compound has the following percentages by mass:
barium, 58.84%; sulfur, 13.74%; oxygen, 27.43%. Determine the empirical formula of the compound.
38. If a 1.271-g sample of aluminum metal is heated in a
chlorine gas atmosphere, the mass of aluminum chloride produced is 6.280 g. Calculate the empirical formula of aluminum chloride.
39. If cobalt metal is mixed with excess sulfur and heated
strongly, a sulfide is produced that contains 55.06%
cobalt by mass. Calculate the empirical formula of the
sulfide.
40. If 2.461 g of metallic calcium is heated in a stream of
chlorine gas, 4.353 g of Cl2 is absorbed in forming the
metal chloride. Calculate the empirical formula of calcium chloride.
41. If 10.00 g of copper metal is heated strongly in the
air, the sample gains 2.52 g of oxygen in forming
an oxide. Determine the empirical formula of this
oxide.
47. A compound with the empirical formula CH2 was
found to have a molar mass of approximately 84 g.
What is the molecular formula of the compound?
48. A compound with the empirical formula CH4O was
found in a subsequent experiment to have a molar
mass of approximately 192 g. What is the molecular
formula of the compound?
49. A compound having an approximate molar mass of
165–170 g has the following percentage composition
by mass: carbon, 42.87%; hydrogen, 3.598%; oxygen,
28.55%; nitrogen, 25.00%. Determine the empirical
and molecular formulas of the compound.
50. A compound consists of 65.45% C, 5.492% H, and
29.06% O on a mass basis and has a molar mass of
approximately 110. Determine the molecular formula
of the compound.
Critical Thinking
51. Use the periodic table inside the back cover of this
text to determine the atomic mass (per mole) or molar mass of each of the substances in column 1, and
find that mass in column 2.
Column 1
Column 2
(1) molybdenum
(a)
33.99 g
42. A compound has the following percentage composition by mass: copper, 33.88%; nitrogen, 14.94%; oxygen, 51.18%. Determine the empirical formula of the
compound.
(2) lanthanum
(b)
79.9 g
(3) carbon tetrabromide
(c)
95.94 g
(4) mercury(II) oxide
(d) 125.84 g
43. Sodium and nitrogen form two binary compounds.
The percentages of the elements in these compounds
are:
(5) titanium(IV) oxide
(e) 138.9 g
(6) manganese(II) chloride
(f) 143.1 g
(7) phosphine, PH3
(g) 156.7 g
Compound 1:
83.12% Na; 16.88% N
(8) tin(II) fluoride
(h) 216.6 g
Compound 2:
35.36% Na; 64.64% N
(9) lead(II) sulfide
(i) 239.3 g
(10) copper(I) oxide
(j) 331.6 g
Calculate the empirical formula of each of the compounds.
6.8 Calculation of Molecular Formulas
Questions
44. How does the molecular formula of a compound differ from the empirical formula? Can a compound’s
empirical and molecular formulas be the same?
Explain.
45. What information do we need to determine the molecular formula of a compound if we know only the
empirical formula?
190
Chapter 6 Chemical Composition
52. Complete the following table.
Mass of
Sample
Moles of
Sample
Atoms in
Sample
5.00 g A1
0.00250 mol Fe
2.6 1024 atoms Cu
0.00250 g Mg
2.7 103 mol Na
1.00 104 atoms U
Page 5 of 5
53. Complete the following table.
Mass of
Sample
Moles of
Sample
Molecules in
Sample
Column 2
Atoms in
Sample
(a) 6.022 1023
(b) atomic mass
(c) mass of 1000 hydrogen atoms
4.24 g C6H6
(d) benzene, C6H6, and acetylene, C2H2
0.224 mol H2O
2.71 1022
molecules CO2
(e) carbon dioxide
(f) empirical formula
(g) 1.66 1024 g
1.26 mol HCl
4.21 1024
molecules H2O
0.297 g
CH3OH
54. A binary compound of magnesium and nitrogen is analyzed, and 1.2791 g of the compound is found to
contain 0.9240 g of magnesium. When a second sample of this compound is treated with water and heated,
the nitrogen is driven off as ammonia, leaving a compound that contains 60.31% magnesium and 39.69%
oxygen by mass. Calculate the empirical formulas of
the two magnesium compounds.
55. When a 2.118-g sample of copper is heated in an
atmosphere in which the amount of oxygen present
is restricted, the sample gains 0.2666 g of oxygen
in forming a reddish-brown oxide. However, when
2.118 g of copper is heated in a stream of pure oxygen, the sample gains 0.5332 g of oxygen. Calculate
the empirical formulas of the two oxides of copper.
56. Find the item in column 2 that best explains or completes the statement or question in column 1.
Column 1
(1) 1 amu
(2) 1008 amu
(3) mass of the “average” atom of an element
(4) number of carbon atoms in 12.01 g of carbon
(5) 6.022 1023 molecules
(6) total mass of all atoms in 1 mol of a compound
(7) smallest whole-number ratio of atoms present in
a molecule
(8) formula showing the actual number of atoms present in a molecule
(9) product formed when any carbon-containing
compound is burned in O2
(10) have the same empirical formulas, but different
molecular formulas
(h) molecular formula
(i) molar mass
(j) 1 mol
57. Calculate the number of grams of iron that contain
the same number of atoms as 2.24 g of cobalt.
58. Calculate the number of grams of cobalt that contain
the same number of atoms as 2.24 g of iron.
59. A strikingly beautiful copper compound with the
common name “blue vitriol” has the following elemental composition: 25.45% Cu, 12.84% S, 4.036%
H, 57.67% O. Determine the empirical formula of the
compound.
60. A 0.7221-g sample of a new compound has been analyzed and found to contain the following masses of
elements: carbon, 0.2990 g; hydrogen, 0.05849 g; nitrogen, 0.2318 g; oxygen, 0.1328 g. Calculate the empirical formula of the compound.
61. When 4.01 g of mercury is strongly heated in air, the
resulting oxide weighs 4.33 g. Calculate the empirical
formula of the oxide.
62. When barium metal is heated in chlorine gas, a binary compound forms that consists of 65.95% Ba and
34.05% Cl by mass. Calculate the empirical formula
of the compound.
63. A particular compound in the chemistry laboratory is
found to contain 7.2 1024 atoms of oxygen, 56.0 g
of nitrogen, and 4.0 mol of hydrogen. What is its empirical formula?
64. The compound A2O is 63.7% A (a mystery element)
and 36.3% oxygen. What is the identity of element
A?
65. A molecule of an organic compound has twice as
many hydrogen atoms as carbon atoms, the same
number of oxygen atoms as carbon atoms, and oneeighth as many sulfur atoms as hydrogen atoms. The
molar mass of the compound is 152 g/mol. What are
the empirical and molecular formulas for this compound?
Chapter Review
191