Relative atomic mass — AQA GCSE Chemistry
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Relative atomic mass explained
Relative atomic mass covers 2 specification points, including The relative atomic mass of an element is an… and Students should be able to calculate the….
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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. 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.
Students should be able to calculate the relative atomic mass of an element given the percentage abundance of its isotopes.
To calculate relative atomic mass from percentage abundances, multiply each isotope's mass number by its percentage abundance, add the products, then divide by 100. For example, if an element is 20% isotope-10 and 80% isotope-11, the calculation is ((10 × 20) + (11 × 80)) ÷ 100 = (200 + 880) ÷ 100 = 10.8. The answer is a weighted average and may be quoted to a suitable number of decimal places. If abundances are given as decimals or fractions, multiply each mass by its decimal abundance and add, without dividing by 100. Always check that the abundances total 100% before starting.
Your focus
- Define relative atomic mass as a weighted average of isotope masses.
- Explain why the relative atomic mass of an element may not be a whole number.
- Link isotope abundance to the value of relative atomic mass shown on the periodic table.
Show all 6 objectives
- Calculate relative atomic mass from percentage abundances of isotopes.
- Adapt the calculation when abundances are given as decimals or fractions.
- Check that a calculated relative atomic mass lies between the isotope mass numbers.
Relative atomic mass flashcards
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Isotopes
Relative atomic mass exam tips
Marking 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.
- The periodic table value for an element is its relative atomic mass.
- Multiply each isotope's mass number by its percentage abundance.
- Add the products for all isotopes.
- Divide the total by 100 when abundances are percentages.
- If abundances are given as decimals, multiply each mass by its decimal abundance and add without dividing by 100.
- Check that percentage abundances sum to 100 before calculating.
- Quote the final answer as a weighted average, often to one decimal place.
Examiner 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.
- 💡Lay out the calculation as [ (mass × abundance) + (mass × abundance) ] ÷ 100 to keep it mathematically correct.
- 💡Check that the abundances add to 100% before you begin.
- 💡Sanity-check the answer: it must lie between the lowest and highest isotope mass numbers.
Common Mistakes
- Treating relative atomic mass as the mass of the most common isotope; instead treat it as a weighted average of all isotopes.
- Averaging isotope masses without including their abundances; instead weight each mass by its percentage abundance.
- Confusing relative atomic mass with mass number; instead note that mass number is a whole number for one isotope, while Aᵣ is an average for the element.
- Forgetting to divide by 100 when using percentage abundances; instead divide the sum of the products by 100.
- Using the number of isotopes rather than their abundances as weights; instead weight each mass by its percentage or decimal abundance.
- Adding the isotope mass numbers and dividing by the number of isotopes; instead use the abundance-weighted method.