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Chapter 8: The Mole Activity 2
Name_____________________
Period ______Date__________
Today’s activity will introduce you to a unit of measure without which chemistry would not exist. It is a
unit much like a dozen, which helps us count things. “Why would a chemist need to count things?” you
might ask. Examine the following chemical equation:
2H2 + O2 ------> 2H2O
This can be interpreted as two molecules of hydrogen and one molecule of oxygen combining to form two
water molecules. But how often do we as chemists limit our reactions to one or two molecules? Usually a
reaction is done on a billion, trillion scale where an unimaginable number of molecules are involved. We
need to count our molecules, how can we do this? This brings us to the Mole.
We will start by investigating two egg farmers (a quail farmer and a chicken farmer). They produce such
large numbers of eggs each year that they can’t count them individually, so they use dozens, in some
cases they use mass (weighing is often easier than counting). Although eggs vary in size, shape, and mass
we can use average values and get a general idea. Fill in the following table using the information below.
1 quail egg = 2.34 grams
1 dozen quail eggs = 12 eggs
Quail
# eggs in
sample
1
10
438
1 dozen
12 dozen
1 million
Quail
Mass of
Sample
1 chicken egg = 37.44 grams
1 dozen chicken eggs = 12 eggs
Chicken
# eggs in
sample
1
10
438
1 dozen
12 dozen
1 million
Chicken
Mass of
Sample
Ratio of #
of eggs
Ratio of
Masses
1) How many times larger in mass is a chicken egg than a quail egg?
2) What conclusion can you draw about the masses of two samples when the number of eggs in those
samples is equal?
Let’s use the egg model and apply it to atoms. Like eggs, atoms of the same kind may have different
masses (isotopes). But, the Periodic Table lists an average atomic mass for each type of atom. Keep in
mind that the masses on the periodic table are in “atomic mass units” where 1 a.m.u. is equal to the mass
of a proton (or neutron). To find the mass of an entire atom we add up the number of protons and
neutrons in that atom. To find the mass of a molecule we add up the atoms in that molecule. Remember in
this class we will be rounding the atomic mass units (amu’s) to the nearest whole number. Here are some
examples on the next page:
Atomic Mass
1 amu
1 H atom
16 amu
1 O atom
12 amu
1 C atom
23 amu
1 Na atom
64 amu
1 Cu atom
Table A
1
1
1
1
1
Molecular Mass
18 amu
H2O molecule
44 amu
CO2 molecule
32 amu
O2 molecule
2 amu
H2 molecules
16 amu
CH4 molecules
5) Take a moment to check the values in the table. How were the values for atomic mass calculated?
How were the values for molecular mass calculated? Fill in the table below.
Oxygen
# of atoms
in sample
1
10
438
1 dozen
12 dozen
1 million
1 mole
Oxygen
Mass of
Sample
16 amu
16 grams
Sulfur
# of atoms
in sample
1
10
438
1 dozen
12 dozen
1 million
1 mole
Sulfur
Mass of
Sample
32 amu
Ratio of # of
atoms
Ratio of
Masses
32 grams
6) Even though you may not be sure “how many” a mole is, did you have problems filling in the above table?
Why or why not?
7) If we were to change the mass unit we used to grams, what could you say about the relative number of atoms of
oxygen in a 16 gram sample, and the number of atoms of sulfur in a 32 gram sample?
8) A bottle contains a 32 gram sample of sulfur. How many atoms do you think are in this bottle?
A long time ago chemists discovered what you just discovered by answering question 8. If they were talking about
the mass of one atom of an element they talked about its mass in amu’s. This was not very helpful as most often
they were dealing with many more atoms than just one or two. What they found was that if they grouped the
atoms or molecules together into a large sample, 6.02 x 1023 atoms or molecules, this number of atoms or molecules
had a mass in grams equal to the mass of 1 atom in amu’s. Below are the “molar masses” of the same substances
we looked at on the previous page.
Table B
Gram Atomic Mass
1 gram
1 mole of H atoms
16
gram
1 mole of O atoms
12 gram
1 mole of C atoms
1 mole of Na atoms 23 gram
1 mole of Cu atoms 64 gram
Gram Molecular Mass
1 mole of H2O molecules 18 gram
44 gram
1 mole of CO2 molecules
32 gram
1 mole of O2 molecules
2 gram
1 mole of H2 molecules
1 mole of CH4 molecules 16 gram
9) Compare the two tables on the last page, A and B. What changed?
10) What can you conclude about the numbers of atoms in a 64 gram sample of copper and a 23 gram
sample of sodium?
So how many “things” are in a mole? By estimating the size of atoms, and taking volume measurements of
1 mole samples it has been estimated that:
mole = 6.02 X 1023 (particles) OR 602,200,000,000,000,000,000,000 particles
11) Checking for understanding:
a. 1 mole of chlorine = ___________atoms of chlorine = ___________ g of chlorine
b. 1 mole of ammonia (NH3) = _______________molecules of ammonia = ___________g of ammonia
c. 36 grams of carbon atoms = ___________moles of carbon = _________________ atoms of carbon
d. 1 mole of chlorine gas (Cl2) = __________________ atoms of chlorine
e. 6.25 moles of cars = _______________________ individual cars
f. Theoretically we could count cars in moles, but why don’t we?
g. NaCl is the formula unit for table salt. What is the gram formula mass of 1.0 mole of NaCl?
h. Why do we use gram formula mass with NaCl instead of gram molecular mass (like H2O and NH3)?