AQA · GCSE · Chemistry
Course version: 8462 · Version from Sept 2018
The relative atomic mass of an element is an average value that takes…
Full specification wording
Relative atomic mass: The relative atomic mass of an element is an average value that takes account of the abundance of the isotopes of the element.
Most elements exist as a mixture of isotopes: atoms with the same number of protons but different numbers of neutrons.
Checked against the AQA 8462 specification (opens in a new tab) on .
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- 3 key terms
Study notes
From the checked teaching notes for this specification point.
The idea
Most elements exist as a mixture of isotopes: atoms with the same number of protons but different numbers of neutrons. Because the isotopes of an element have different masses and occur in different abundances, the relative atomic mass (Aᵣ) is a weighted average, not simply the mass of one atom. For example, chlorine consists of about 75% chlorine-35 and 25% chlorine-37, giving a relative atomic mass of about 35.5. The value is relative to 1/12 of the mass of a carbon-12 atom, and it is not a whole number when isotopes are present in significant amounts. Aᵣ is the value shown for each element on the periodic table.
What earns the marks
- Isotopes are atoms of the same element with different numbers of neutrons and therefore different mass numbers.
- Relative atomic mass is a weighted average that reflects the abundance of each isotope.
- The average is relative to 1/12 of the mass of a carbon-12 atom.
- Aᵣ is often not a whole number because it is an average of isotope masses.
- The periodic table value for an element is its relative atomic mass.
Common mistakes and how to put them right
Treating relative atomic mass as the mass of the most common isotope.
Put it right: Treat it as a weighted average of all isotopes.
Averaging isotope masses without including their abundances.
Put it right: Weight each mass by its percentage abundance.
Confusing relative atomic mass with mass number.
Put it right: Note that mass number is a whole number for one isotope, while Aᵣ is an average for the element.
Exam tips
- Quote the definition of relative atomic mass precisely, including the word average.
- Link the idea of abundance to the reason Aᵣ is not a whole number.
- Use a named element such as chlorine or copper to illustrate the idea in an explanation.
Try the lesson
The start of the Relative atomic mass: The relative atomic mass of an element is an average value that takes account of the abundance of the isotopes of the element. lesson, made from checked teaching notes for this specification point.
Have a go first
Why is the relative atomic mass of an element often not a whole number?
Listen
Most elements exist as a mixture of isotopes: atoms with the same number of protons but different numbers of neutrons. Because the isotopes of an element have different masses and occur in different abundances, the relative atomic mass (Aᵣ) is a weighted average, not simply the mass of one atom. For example, chlorine consists of about 75% chlorine-35 and 25% chlorine-37, giving a relative atomic mass of about 35.5. The value is relative to 1/12 of the mass of a carbon-12 atom, and it is not a whole number when isotopes are present in significant amounts. Aᵣ is the value shown for each element on the periodic table. A strong answer earns marks for these points. Isotopes are atoms of the same element with different numbers of neutrons and therefore different mass numbers. Relative atomic mass is a weighted average that reflects the abundance of each isotope. The average is relative to 1/12 of the mass of a carbon-12 atom. Aᵣ is often not a whole number because it is an average of isotope masses. A common slip: Treating relative atomic mass as the mass of the most common isotope. Instead, treat it as a weighted average of all isotopes. Exam tip: Quote the definition of relative atomic mass precisely, including the word average.
See it
Key words
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Isotopes
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Relative atomic mass
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Abundance
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See every definition
- Isotopes
- Atoms of the same element with different numbers of neutrons and therefore different mass numbers.
- Relative atomic mass
- A weighted average that reflects the abundance of each isotope, relative to 1/12 of the mass of a carbon-12 atom.
- Abundance
- How much of each isotope of an element is present, which the weighted average reflects.
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