The atomic
weight
of an element is the weighted average of the masses of the isotopes of
that element. The weighted average is determined using the abundance and
mass of each isotope. Most elements have more than one naturally occurring
isotopes.
For example,
there are two naturally occurring isotopes of copper, copper-63
and copper-65. The natural abundances of the
isotopes is 69.2% and 30.8%
respectively.
To determine
the atomic weight:
Step
1: Multiply the mass number and the relative abundance (as a
decimal). The mass of the electron is insignificant in this calculation
and is not used.
isotope |
mass number
(amu) |
x |
abundance
(as a decimal) |
= |
result |
copper-63 |
63 amu |
x |
0.692 |
= |
43.596 |
copper-65 |
65 amu |
x |
0.308 |
= |
20.020 |
Step
2: Add
up your results.
|
Atomic Weight |
|
63.616 |
On your periodic table you can see that the atomic weight is 63.546.
The reason for the difference between the actual atomic weight (as seen
on the perioidc table) and the calculated atomic weight is due to the
fact that the masses of the proton and neutron is not exactly 1 amu. The
mass of a proton is approx 1.008 amu. The mass of a neutron is approximately
1.009 amu.
|