Final Chemistry Notebook

Learning Goal for this section:
The student will learn why the mole is important and how themolecular formula of a compound can be determined
experimentally.
The Mole: Its History and Use
Simply put, the mole represents a number. Just as the term dozen refers to the number twelve, the
mole represents the number 6.02 x 10 23 . (If you're confused by the form of this number refer to our
The Metric System module).
Now that's a big number! While a dozen eggs will make a nice omelet, a mole of eggs will fill all of the
oceans on earth more than 30 million times over. Think about it: It would take 10 billion chickens
laying 10 eggs per day more than 10 billion years to lay a mole of eggs. So why would we ever use
such a big number? Certainly the local donut store is not going to "supersize" your dozen by giving
you a mole of jelly-filled treats.
The mole is used when we're talking about numbers of atoms and molecules. Atoms and molecules
are very tiny things. A drop of water the size of the period at the end of this sentence would contain
10 trillion water molecules. Instead of talking about trillions and quadrillions of molecules (and more),
it's much simpler to use the mole.
History of the mole
The number of objects in one mole, that is, 6.02 x 10 23 , is commonly referred to as Avogadro's
number. Amedeo Avogadro was an Italian physics professor who proposed in 1811 that equal
volumes of different gases at the same temperature contain equal numbers of molecules. About fifty
years later, an Italian scientist named Stanislao Cannizzaro used Avogadro's hypothesis to develop a
set of atomic weights for the known elements by comparing the masses of equal volumes of gas.
Building on this work, an Austrian high school teacher named Johann Josef Loschmidt calculated the
size of a molecule of air in 1865, and thus developed an estimate for the number of molecules in a
given volume of air. While these early estimates have since been refined, they led to the concept of
the mole - that is, the theory that in a defined mass of an element (its atomic weight) there is a
precise number of atoms: Avogadro's number.
Molar mass
A sample of any element with a mass equal to that element's atomic weight (in grams) will contain
precisely one mole of atoms (6.02 x 10 23 atoms). For example, helium has an atomic weight of 4.00
amu. Therefore, 4.00 grams of helium will contain one mole of helium atoms. You can also work with
fractions (or multiples) of moles:
Mole/weight relationship
examples using helium
moles
helium
# helium
atoms
grams
helium
¼ 1.505 X 10 23 1g
½ 3.01 X 10 23 2g
1 6.02 X 10 23 4g
2 1.204 X 10 24 8g
10 6.02 X 10 24 40g
92