*Calculating Atomic Masses (430 only)
a. The periodic table lists atomic masses for each element.
b. The atomic mass is not a whole number because it is the weighted average of the masses of the different isotopes of that element.
c. Atomic masses do not always increase directly with relation to atomic number. See the exceptions for Co and Ni; Cu and Te; and Te, I and Xe.
Example 1 An element has two isotopes: one of mass 63 u; the other with a mass of 65 u.
If the relative abundance of the isotopes is 69.1% and 30.9 %, respectively, find the atomic mass of the element.
First convert percents to decimals out of 1. Then, multiply each mass by their relative abundance and add them up:
the weighted average = atomic mass = 0.691(63) + 0.309(65) = 63.6 units
Example 2 If the relative abundance of a neon isotope was 5.7%, and the rest was only Ne-20, what was the mass number of the minor isotope?
The rest = 100% - 5.7% = 94.3% = mass of 20
Unknown mass = 5.7%.
Periodic table reveals the weighted average to be 20.18 units, so
0.943(20) + 0.057x = 20.18
x = 23.1, so the mass number of the other isotope is about 23.