The Mole
The identity of a substance is defined not only by the types of atoms or ions it contains, but by the
quantity of each type of atom or ion. For example, water, H2O, and hydrogen peroxide, H2O2, are alike
in that their respective molecules are composed of hydrogen and oxygen atoms. However, because a
hydrogen peroxide molecule contains two oxygen atoms, as opposed to the water molecule, which
has only one, the two substances exhibit very different properties. Today, we possess sophisticated
instruments that allow the direct measurement of these defining microscopic traits; however, the
same traits were originally derived from the measurement of macroscopic properties (the masses
and volumes of bulk quantities of matter) using relatively simple tools (balances and volumetric
glassware). This experimental approach required the introduction of a new unit for amount of
substances, the mole, which remains indispensable in modern chemical science. The mole is an
amount unit similar to familiar units like pair, dozen, gross, etc. It provides a specific measure of the
number of atoms or molecules in a bulk sample of matter. A mole is defined as the amount of
substance containing the same number of discrete entities (such as atoms, molecules, and ions) as
the number of atoms in a sample of pure 12C weighing exactly 12 g. One Latin connotation for the
word “mole” is “large mass” or “bulk,” which is consistent with its use as the name for this unit. The
mole provides a link between an easily measured macroscopic property, bulk mass, and an extremely
important fundamental property, number of atoms, molecules, and so forth. The number of entities
composing a mole has been experimentally determined to be 6.02214179 × 1023, a fundamental
constant named Avogadro’s number (NA) or the Avogadro constant in honor of Italian scientist
Amedeo Avogadro. This constant is properly reported with an explicit unit of “per mole,” a
conveniently rounded version being 6.022 × 1023/mol.
Consistent with its definition as an amount unit, 1 mole of any element contains the same number of
atoms as 1 mole of any other element. The masses of 1 mole of different elements, however, are
different, since the masses of the individual atoms are drastically different. The molar mass of an
element (or compound) is the mass in grams of 1 mole of that substance, a property expressed in
units of grams per mole (g/mol). Because the definitions of both the mole and the atomic mass unit
are based on the same reference substance, 12C, the molar mass of any substance is numerically
equivalent to its atomic or formula weight in amu. Per the amu definition, a single 12C atom weighs
12 amu (its atomic mass is 12 amu). According to the definition of the mole, 12 g of 12C contains 1
mole of 12C atoms (its molar mass is 12 g/mol). This relationship holds for all elements, since their
atomic masses are measured relative to that of the amu-reference substance, 12C. Extending this
principle, the molar mass of a compound in grams is likewise numerically equivalent to its formula
mass in amu.
QUESTIONS:
1. Why do you think chemists prefer using the mole? Why don't they just count each particle?
2. Can you give an example of a way that you use moles in your everyday life? If you can't think
of one, can you create an instance where using the mole might be a benefit for you?
3. What's another name for the mole?
A. The Chemist's Constant
B. Sextillion Power
C. Power Point
D. Avogadro's Number
4. The mole is a really helpful way to count really small things. Are there other methods we use
to help us count? Do we have a number that helps us group really large things?
