1. Average Atomic Mass

2. 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**

3. 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

4. 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

5. 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= Cu-65= (0.6917) (0.3083) (62.939598) (64.927793) = = 43.53531994 20.01723858 + 63.55 amu

6. 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= Ne-21= (0.9092) (0.0057) (20) (21) = = 18.184 0.1197 + 20.07 amu Ne-22=(0.0882)=1.764(22)