Relative mass

• The strong nuclear force holding together protons and neutrons comes at the expense of the loss of a fraction of their mass. This small amount of mass lost is called the mass defect. This is why the masses of all elements are compared to carbon-12 as a standard for the measurement of atomic masses
• • The mass of a carbon-12 isotope is defined as exactly 12 atomic mass units (12 u)
  • The standard mass for atomic mass is 1u, the mass of 1/12th of an atom of carbon-12
  • On this scale, 1u is approximately the mass of a proton or a neutron
• Relative isotopic mass is the mass of an isotope relative to 1/12th of the mass of an atom of carbon-12.
• Relative atomic mass $A_r$ is the weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12
• The weighted mean mass takes account of:

  • the percentage abundance of each isotope
  • the relative isotopic mass of each isotope
• The percentage abundance of the isotopes in a sample of an element are found experimentally using a mass spectrometer
• When given isotopes of an element and their percentage abundances, to find the relative atomic mass of the element:

  The sum of the isotopic masses multiplied my their percentage abundances, all over 100 or the sum of the percentages
• Relative isotopic mass is the exact mass of an atom of a particular isotope and not a mean value like relative atomic mass