Atomic Mass Examples in Chemistry

Start with the recap, study the fully worked examples, then use the practice problems to check your understanding of Atomic Mass.

This page combines explanation, solved examples, and follow-up practice so you can move from recognition to confident problem-solving in Chemistry.

Concept Recap

The weighted average mass of all naturally occurring isotopes of an element, expressed in atomic mass units (amu), where each isotope's mass is weighted by.

The number under each element on the periodic table—a weighted average of all its isotopes.

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How to Use These Examples

  • Read the first worked example with the solution open so the structure is clear.
  • Try the practice problems before revealing each solution.
  • Use the related concepts and background knowledge badges if you feel stuck.

What to Focus On

Core idea: Atomic mass accounts for the natural mix of isotopes and their relative abundances.

Common stuck point: It's not a whole number because isotopes have different masses and abundances.

Sense of Study hint: When calculating atomic mass from isotope data, use a weighted average. First list each isotope's mass and its percent abundance. Then convert each percent to a decimal (divide by 100). Finally, multiply each isotope's mass by its decimal abundance and sum: \text{atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance}).

Worked Examples

Example 1

easy
Define 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.

Solution

  1. 1
    Atomic mass is the weighted average of the masses of all naturally occurring isotopes of an element, measured in atomic mass units (amu).
  2. 2
    Carbon has two stable isotopes: {}^{12}\text{C} (mass 12, abundance 98.9%) and {}^{13}\text{C} (mass 13, abundance 1.1%).
  3. 3
    Weighted average: 12 \times 0.989 + 13 \times 0.011 = 11.87 + 0.14 = 12.01\,\text{amu}.

Answer

\text{Atomic mass of C} = 12.01\,\text{amu (weighted average of isotopes)}
The atomic mass on the periodic table reflects the natural isotope distribution. Since {}^{12}\text{C} is by far the most abundant isotope, the average is very close to 12 but slightly above due to the small contribution of {}^{13}\text{C}.

Example 2

medium
Silicon 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.

Practice Problems

Try these problems on your own first, then open the solution to compare your method.

Example 1

medium
Chlorine 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 2

hard
An 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.
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Related Concepts

IsotopeMolar Mass

Background Knowledge

These ideas may be useful before you work through the harder examples.

isotope