Practice Atomic Mass in Chemistry
Use these practice problems to test your method after reviewing the concept explanation and worked examples.
Quick Recap
The weighted average mass of all naturally occurring isotopes of an element, in atomic mass units.
The number under each element on the periodic table—a weighted average of all its isotopes.
Example 1
easyDefine atomic mass and explain why the atomic mass of carbon is listed as 12.01\,\text{amu} rather than exactly 12\,\text{amu} on the periodic table.
Example 2
mediumSilicon has three stable isotopes: {}^{28}\text{Si} (92.23%, mass 27.977), {}^{29}\text{Si} (4.67%, mass 28.976), and {}^{30}\text{Si} (3.10%, mass 29.974\,\text{amu}). Calculate the atomic mass of silicon.
Example 3
mediumChlorine has two stable isotopes: {}^{35}\text{Cl} (mass 34.97\,\text{amu}, 75.77%) and {}^{37}\text{Cl} (mass 36.97\,\text{amu}, 24.23%). Calculate its atomic mass.
Example 4
hardAn element has two isotopes. Isotope A has mass 10.01\,\text{amu} and isotope B has mass 11.01\,\text{amu}. The atomic mass listed on the periodic table is 10.81\,\text{amu}. Calculate the percent abundance of each isotope.