The Foundations of Chemistry

THE ATOMIC WEIGHT SCALE AND ATOMIC WEIGHTS

We said in Section 2-5 that the atomic weight scaleis based on the mass of the carbon-
12 isotope. As a result of action taken by the International Union of Pure and Applied
Chemistry in 1962,

one amuis exactly 1/12 of the mass of a carbon-12 atom.

This is approximately the mass of one atom of^1 H, the lightest isotope of the element
with lowest mass.
In Section 2-6 we said that one mole of atoms, contains 6.022 1023 atoms. The mass
of one mole of atoms of any element, in grams, is numerically equal to the atomic weight
of the element. Because the mass of one carbon-12 atom is exactly 12 amu, the mass of
one mole of carbon-12 atoms is exactly 12 grams.
To show the relationship between atomic mass units and grams, let us calculate the
mass, in amu, of 1.000 gram of^1206 C atoms.

5-9

Described another way, the mass of

one atom of^1206 C is exactly12 amu.

5-9 The Atomic Weight Scale and Atomic Weights 191

_?_amu1.000 g^1206 C atoms

6.022 1023 amu (in 1 gram)

12 amu

 12

06 C atom

6.022 10 23 12 06 C atoms



1 mol^1206 C atoms

1 mol^1206 C



12 g^1206 C atoms

Thus,

1 g6.022 1023 amu or 1 amu1.660 10 
^24 g

At this point, we emphasize the following:

1.The atomic number, Z,is an integer equal to the number of protons in the nucleus

of an atom of the element. It is also equal to the number of electrons in a neutral

atom. It is the same for all atoms of an element.

2.The mass number, A,is an integer equal to the sumof the number of protons and

the number of neutrons in the nucleus of an atom of a particular isotopeof an element.

It is different for different isotopes of the same element.

3.Many elements occur in nature as mixtures of isotopes. The atomic weightof such

an element is the weighted average of the masses of its isotopes. Atomic weights are

fractional numbers, not integers.

The atomic weight that we determine experimentally (for an element that consists of more
than one isotope) is such a weighted average. The following example shows how an atomic
weight can be calculated from measured isotopic abundances.

EXAMPLE 5-2 Calculation of Atomic Weight

Three isotopes of magnesium occur in nature. Their abundances and masses, determined by
mass spectrometry, are listed in the following table. Use this information to calculate the atomic
weight of magnesium.

We saw in Chapter 2 that Avogadro’s

number is the number of particles of a

substance in one mole of that

substance. We now see that Avogadro’s

number also represents the number of

amu in one gram. You may wish to

verify that the same result is obtained

regardless of the element or isotope

chosen.