General Chemistry

1. Chemistry Unit 9

2.  Chemists use the mole to count atoms,
molecules, ions and formula units.
 A mole always contains the same number of
particles, however, moles of different
substances have different masses.
 The molar mass of a compound can be
calculated from its chemical formula and can be
used to convert from mass to moles of that
compound.
 A molecular formula of a compound is a whole-
number multiple of its empirical formula

4.  Explain how a mole is used to indirectly
count the number of particles of matter.
 Relate the mole to a common everyday
counting unit.
 Convert between moles and number of
representative particles.

5. The mole is the SI base unit for measure of
amount of a substance: 6.0221367 x 1023
 The number of carbon atoms in exactly 12 g
of pure carbon-12.
 Called Avogadro’s number – Italian physicist
who in 1811, determined the volume of 1
mol of gas.
 By mass, we can determine the number of
particles (atoms, molecules) in a sample.
 We typically round to 3 sig figs – 6.02 x 1023

6. The mole is a number. What other unit is used
in a similar manner?
 A dozen flowers, doughnuts or eggs.
 A baker’s dozen of cookies or bagels.
 A pair of socks or friends
If you have a dozen flowers and a dozen eggs,
do they weigh the same?

7. In order to convert between moles and number
of particles we need to use this ratio of
equivalent values (conversion factor) to
express the same quantity in different units.

8. How many particles are in 3.5 mols?
How many moles of atoms are in 9.63 x 1026
atoms?

9. What does the mole measure?
A. mass of a substance
B. amount of a substance
C. volume of a gas
D. density of a gas

10. What is the conversion factor for determining
the number of moles of a substance from a
known number of particles?
A.
B.
C. 1 particle  6.02  1023
D. 1 mol  6.02  1023 particles

11.  Page 322 #1-4; page 324 #5-14

13.  Relate the mass of an atom to the mass of a
mole of atoms.
 Convert between number of moles and the
mass of an element.
 Convert between number of moles and
number of atoms of an element.

14. Molar mass is the mass in grams of one mole of
any pure substance.
 Units are given in g/mol
 Mass of the periodic table is given in amu,
but also g/mol

15. If I need 3 mols of Cu, how do I measure the
amount?
I measured 5.0g of Iron, how many atoms do I
have?

16. 

moles
x
grams
mole
= grams
grams
x
moles
gram
= moles

17. How many atoms of gold are in a U.S. Eagle
bullion coin with a mass of 31.1g?
How much does 5.8 x 1015 atoms of lead weigh?

18. The mass in grams of 1 mol of any pure
substance is:
A. molar mass
B. Avogadro’s number
C. atomic mass
D. 1 g/mol

19. Molar mass is used to convert what?
A. mass to moles
B. moles to mass
C. atomic weight
D. particles

20.  Page 328 #15-16; page 329 #17-18
 Page 331 #19-21; page 332 #22-27

22.  Recognize the mole relationships shown by a
chemical formula
 Calculate the molar mass of a compound.
 Convert between number of moles and mass
of a compound.
 Apply conversion factors to determine the
number of atoms or ions in a known mass of
a compound

23. Steps to calculate molar mass:
1. Count the number of atoms in each
molecule.
2. Find the molar mass of each atom.
3. Multiply the molar mass of each atom to
the number of atoms in a compound.
4. Add the total molar masses together.

24. Find the molar mass of the following
compounds/molecules.
 H2O
 NaCl
 H2SO4

25.  Al2O3
 Fe2(SO4)3
 CCl2F2

26. To determine the number of atoms or ions in a
known mass of a compound
1. Find the molar mass of the compound.
2. Use molar mass and the mole as conversion
factors to get the units needed.

27. What is the mass of 2.5 mols of (C3H5)2 S?

28. Calculate the number of moles of Ca(OH)2 in
325g of the compound?

29. How many atoms are in 212g of water?

30. How many moles of OH— ions are in 2.50 moles
of Ca(OH)2?
A. 2.00
B. 2.50
C. 4.00
D. 5.00

31. How many particles of Mg are in 10 moles of
MgBr2?
A. 6.02 x 1023
B. 6.02 x 1024
C. 1.20 x 1024
D. 1.20 x 1025

32.  Page 335 #29-36; page 336 #37-41; page 339
#42-46

34.  Explain what is meant by the percent
composition of a compound.
 Determine the empirical and molecular
formulas for a compound from percent and
actual mass data.
 Explain what a hydrate is and relate the
name of the hydrate to its composition.
 Determine the formula of a hydrate from
laboratory data.

35. The percent composition is a percent by mass
of each element in a compound.
Steps to determine percent composition of a
compound:
1. Assume 1 mole of a compound.
2. Calculate molar mass of each element in
the compound.
3. Use each element’s molar mass to calculate
percent by mass.

36. Percent by mass is a description of the amount
of an element in a compound.
 Percent by mass =
%mass =
mass of 1 mole of element
molar mass of compound
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37. What is the percent by mass of each element
in NaHCO3?

38. The empirical formula is the smallest whole
number ratio of elements in a compound
 This ratio provides the subscripts for the
elements.
 May or may not be the same as the actual
molecular formula.
 If they are different the molecular formula will
be a simple multiple of the empirical formula.
 Hydrogen peroxide: HO- empirical formula
H2O2 – actual formula (molecular formula)

39. Steps to figure empirical formula from percent
composition:
1. Assume an overall 100g sample of the compound.
2. Each element’s percentage can be used as mass
in calculations.
3. Use this ‘mass’ to convert to moles. This
provides a ‘mole ratio’ for the compound.

40.  Since these mole ratios are not whole
numbers, we convert them to whole numbers
what can be used as subscripts by dividing
them all by the smallest ratio. (We assume
the smallest mole ratio is a 1 in the
compound)

41. A compound has the following mass percentages:
C – 48.64%, H – 8.16%, O – 43.20%
What is the empirical formula for this molecule?

42. The molecular formula specifies the actual
number of atoms of each element in one
molecule/formula unit of the substance.

43. Steps to determine the molecular formula:
1. Determine the molar mass of the empirical
formula.
2. Determine the molar mass of the actual
compound. (might be given to you)
3. Divide the molar mass of the actual
compound by the molar mass of the
empirical formula
4. Multiply all subscripts of the empirical
formula by this molar mass ratio.

45. The mass of benzene has been experimentally
determined to be 78.12 g. We know that
benzene is 92 % C by mass and 8 % H by mass.
What is the molecular formula of benzene?

46.  Page 344 #54-57
 Page 348 #58-61
 Page 350 #62-66

47. Hydrates are solid ionic compounds in which
water molecules are trapped.
 Hydrates are formed when water molecules
adhere to the ions as the solid forms.
 Water molecules become a part of the
crystal solid structure.
 The number of water molecules associated
with each molecule is written following a dot
after the molecular formula:
 Na2CO310H2O

48.  Names of these compounds are named with a
prefix representing the number of water
molecules and the word hydrate.
 Na2CO3  10H2O – sodium carbonate decahydrate
 Prefixes are the same as the ones used in naming
covalent molecules

50. An anhydrous is a compound without water.
 When a hydrate is heated, water molecules
are driven off leaving the compound.

51. Steps to determining the formula of a hydrate:
1. Determine the initial mass of the compound
prior to heating.
2. Determine the final mass of the compound
after heating.
3. The final mass is used to determine the
number of moles of the anhydrous
compound. Grams to moles

52. 4. Calculate the difference of the initial mass
and the final mass and use this mass to
determine number of moles of water, grams
to moles
5. Number of hydrates per compound
molecule (molar ratio – hydrate: anhydrous)
= moles of H2O/moles of compound
#H2O=
moles of H2O
moles of compound

53. A mass of 2.50 g of blue, hydrated copper
sulfate (CuSO4) ?H2O) is place in a crucible
and heated. After heating, 1.59g of white
anhydrous copper sulfate (CuSO4) remains.
What is the formula for the hydrate? Name
the hydrate.

54. What is the empirical formula for the
compound C6H12O6?
A. CHO
B. C2H3O2
C. CH2O
D. CH3O

55. Which is the empirical formula for hydrogen
peroxide?
A. H2O2
B. H2O
C. HO
D. none of the above

56. Heating a hydrate causes what to happen?
A. Water is driven from the hydrate.
B. The hydrate melts.
C. The hydrate conducts
electricity.
D. There is no change in the
hydrate.

57. A hydrate that has been heated and the water
driven off is called:
A. dehydrated compound
B. antihydrated compound
C. anhydrous compound
D. hydrous compound

58.  Page 353 #74-75; page 354 #76-82

60.  How does the mole apply to balanced
equations?

61. How many grams of each reactant are needed
to run the following reaction? How many
grams of each product are produced? (Hint:
complete, balance, convert)
 CuSO45H2O(aq) + CaCl2(aq) 

62. From the reaction above, how much is needed of
each reactant in the net ionic equation to
produce the balanced amount of the precipitate?
How much precipitate is produced?

64.  The mole is a unit used to count particles
of matter indirectly. One mole of a pure
substance contains Avogadro’s number of
particles.
 Representative particles include atoms,
ions, molecules, formula units, electrons,
and other similar particles.

65.  One mole of carbon-12 atoms has a mass
of exactly
12 g.
 Conversion factors written from
Avogadro’s relationship can be used to
convert between moles and number of
representative particles.
 The mass in grams of 1 mol of any pure
substance is called its molar mass.

66.  The molar mass of an element is
numerically equal to its atomic mass.
 The molar mass of any substance is the
mass in grams of Avogadro’s number of
representative particles of the substance.
 Molar mass is used to convert from moles
to mass. The inverse of molar mass is
used to convert from mass to moles.

67.  Subscripts in a chemical formula indicate
how many moles of each element are
present in 1 mol of the compound.
 The molar mass of a compound is
calculated from the molar masses of all
of the elements in the compound.
 Conversion factors based on a
compound’s molar mass are used to
convert between moles and mass of a
compound.

68.  The percent by mass of an element in a
compound gives the percentage of the
compound’s total mass due to that element.
 The subscripts in an empirical formula give
the smallest whole-number ratio of moles of
elements in the compound.
 The molecular formula gives the actual
number of atoms of each element in a
molecule or formula unit of a substance.
 The molecular formula is a whole-number
multiple of the empirical formula.

69.  The formula of a hydrate consists of the
formula of the ionic compound and the
number of water molecules associated
with one formula unit.
 The name of a hydrate consists of the
compound name and the word hydrate
with a prefix indicating the number of
water molecules in 1 mol of the
compound.
 Anhydrous compounds are formed when
hydrates are heated.

70. What does Avogadro’s number represent?
A. the number of atoms in 1 mol of
an element
B. the number of molecules in
1 mol of a compound
C. the number of Na+ ions in
1 mol of NaCl (aq)
D. all of the above

71. The molar mass of an element is numerically
equivalent to what?
A. 1 amu
B. 1 mole
C. its atomic mass
D. its atomic number

72. How many moles of hydrogen atoms are in one
mole of H2O2?
A. 1
B. 2
C. 3
D. 0.5

73. What is the empirical formula of Al2Br3?
A. AlBr
B. AlBr3
C. Al2Br
D. Al2Br3

74. What is an ionic solid with trapped water
molecules called?
A. aqueous solution
B. anhydrous compound
C. hydrate
D. solute

75. How many water molecules are associated with
3.0 mol of CoCl2 • 6H2O?
A. 18
B. 1.1  1025
C. 3.6  1024
D. 1.8  1024

76. How many moles of Al are in 2.0 mol of Al2Br3?
A. 2
B. 4
C. 6
D. 1

77. How many atoms of hydrogen are in
3.5 mol of H2S?
A. 7.0  1023
B. 2.1  1023
C. 6.0  1023
D. 4.2  1024

78. Which is not the correct formula for an ionic
compound?
A. CO2
B. NaCl
C. Na2SO4
D. LiBr2