Average Atomic Mass
Relative Atomic Masses
 Masses of atoms (in grams) are very small, so
for convenience we use relative masses.
 Carbon-12 is our standard
 Atomic mass (u or amu): 1/12 the mass of a
carbon-12 atom or 1.6605402 x 10-24 g.
**Mass number and atomic mass are close to
each other, but not identical because the
proton and neutron masses deviate slightly
from 1 amu and atomic mass includes
electrons**
Average Atomic Mass
 Average Atomic Mass: the weighted
average of the atomic masses of the
naturally occurring isotopes of an element
 This number appears on the periodic table
below the symbol
Steps
1. Convert % abundance to decimal form

(80 %  0.80)
2. Multiply decimal form of % abundance by
the actual mass of isotope
3. Add the products of each isotope to get the
average atomic mass
Example 1
Copper exists as two isotopes, copper-63
(actual mass 62.939598) with 69.17%
abundance and copper-65 (actual mass
64.927793) with 30.83% abundance. Find the
average atomic mass.
Cu-63
=
(0.6917) (62.939598) = 43.53531994
Cu-65
=
(0.3083) (64.927793) = 20.01723858
63.55 amu
+
Example 2
Calculate the estimated average atomic mass of neon if
neon exists naturally as 90.22% neon-20, 0.57% neon-21,
and 8.82% neon-22.
Note: If the actual mass of each isotope is not available,
the average atomic mass may be estimated by using
the mass number of each isotope.
Ne-20
=
(0.9092) (20) =
18.184
Ne-21
=
(0.0057) (21) =
0.1197
Ne-22
=
(0.0882) (22) =
1.764
+
20.07 amu