1. Isotopic Abundance SCH 3U

2. Atomic Mass The mass of an atom (protons, neutrons, electrons) Relative Atomic Mass: An element’s atomic mass relative to the mass of a carbon-12 atom Atomic Mass Units (amu or u) A unit of mass defined as 1/12 the mass of one C-12 atom OR 1 proton or neutron E.g. C-12 has a mass of 12 X 1 u = 12 u

3. E.g. The mass of a sodium atom is 1.92 times that of C-12. Calculate the atomic mass of sodium in atomic mass units. m c-12 = 12u m Na = 1.92 X 12 u = 23.04 u Compare this to the average atomic mass. 22.99 u What accounts for the difference? It’s an average of an element’s isotopes

4. Isotopic Abundance For most elements, the isotopic composition remains constant E.g. any sample of naturally occurring carbon has the same percentage of C-12 and C-13 Because this is constant, we can use an average value for the atomic mass of carbon The percentage of an isotope in a sample is called its isotopic abundance

5. E.g. Given Isotopic Abundance of C-12 = 98.89% Isotopic Abundance of C-13 = 1.110% Calculate the average atomic mass of carbon. m avg = [m C-12 X % C-12 ] + [m C-13 X % C-13 ] = [12 u X 0.9889] + [13 u X 0.0110] = 11.8668 u + 0.1443 u = 12.01 u

6. E.g. Chlorine has 2 isotopes, Isotopic Abundance of Cl-35 = 75.53% Isotopic Abundance of Cl-37 = 24.47% Calculate the average atomic mass of chlorine. m avg = [m Cl-35 X % Cl-35 ] + [m Cl-37 X % Cl-37 ] = [35 u X 0.7553] + [37 u X 0.2447] = 26.4355 u + 9.0539 u = 35.49 u