Worksheet: Relative Atomic Masses and Isotopic Abundance
In this worksheet, we will practice calculating the relative atomic masses of elements based on the relative abundance of their isotopes.
A sample of lithium contains a mixture of and isotopes with atomic masses 6.01512 u and 7.01600 u respectively. The sample is by moles.
Calculate, to 3 significant figures, the average atomic mass of the sample.
- A6.52 u
- B6.96 u
- C6.06 u
- D6.04 u
- E6.97 u
A second sample of lithium also contains a mixture of and but displays a higher abundance of . How does the average atomic mass of this sample compare with that of the first?
Protons and neutrons have masses of approximately 1 unified atomic mass unit (u) and cannot easily be split into smaller particles. Magnesium consists of protons, neutrons, and electrons and has an atomic mass of 24.3 u. Why is the atomic mass of magnesium significantly greater than 24 u?
- AProton and neutron have masses slightly larger than 1 u.
- BThe atomic mass is calculated based on the relative abundances of natural isotopes.
- CProtons and neutrons increase in mass when they interact to form an atomic nucleus.
- DThe atomic mass is an average of the relative atomic mass, which varies depending on the measurement conditions.
- EThe electrons in magnesium have a total mass of 0.3 u.