Atomic Mass

1. Atomic Mass

2. Isotopes Reminder • Yesterday we learned that the mass of all atoms of a certain element are not always the same • Some atoms may have more neutrons than protons • We call these atoms isotopes

3. Atomic Mass • On your Periodic Tables, each element has its own atomic mass • This mass is the weightedaverage mass of the atoms that make up that element • The masses of the potential isotopes of that element are taken into consideration when the atomic mass is calculated

4. Percent Natural Abundance • For almost all elements, the relative amounts of each isotope in a naturally occurring sample are the same For example… • In any natural sample of Neon atoms: 90.48% are Ne-20 0.27% are Ne-21 9.25% are Ne-22 These percentages are referred to as the percent natural abundance of the isotopes

5. Mass Number vs. Atomic Mass • The mass number can be calculated by adding the number of protons and neutrons The mass of a proton is ____________amu The mass of a neutron is ___________amu 1 1 With that said, an atom’s mass number can be used to determine the atomic mass of that atom **If you are not given the mass of an isotope, you can always use the mass number to figure out the atom’s mass

6. Mass Number vs. Atomic Mass An atom’s mass number can be used to determine the atomic mass of that atom For example If an isotope of carbon has 6protons and 8 neutrons the mass number of the isotope would be 14 The 6 protons would have a total mass of 6 amu: 6 protons x 1amu = 6 amu 1 proton The 8 neutrons would have a total mass of 8 amu: 8 neutronsx 1 amu = 8amu 1 neutron 6 amu + 8 amu = 14 amu

7. Calculating Atomic Mass • These percent natural abundances are used when the atomic mass is calculated Atomic mass = (Fraction of isotope 1 x Mass of isotope 1) + (Fraction of isotope 2 x Mass of isotope 2) + (Fraction of isotope 3 x Mass of isotope 3) + . . . ***As a way to check your calculations, the weighted average atomic mass should be in the range of the masses of the isotopes or the mass numbers used in the calculation

8. Calculating Atomic Mass • These percent natural abundances are used when the atomic mass is calculated • *The percentages need to be converted to their decimal value before they can be included into the calculation 54.3% = 54.3/100 = 0.543 Atomic mass = (Fraction of isotope* 1 x Mass of isotope 1) + (Fraction of isotope* 2 x Mass of isotope 2) + (Fraction of isotope* 3 x Mass of isotope 3) + . . .

9. Calculating Atomic Mass Example 1:(from p. 109 in your book!!) Naturally occurring chlorine consists of 75.77% chlorine-35 (mass 34.97 amu) and 24.23% chlorine-37 (mass 36.97 amu). It’s atomic mass is: Atomic mass = (Fraction of isotope 1 x Mass of isotope 1) + (Fraction of isotope 2 x Mass of isotope 2) Atomic mass = (0.7577 x 34.97 amu) + (0.2423 x 36.97 amu) 35.45 amu

10. Calculating Atomic Mass Example 2: A sample of Cesium is 75%133Cs, 20%132Cs, and 5%134Cs. What is its average atomic mass? Hint: For this problem, you need to use the mass number to find the mass of each isotope Atomic mass = (Fraction of isotope 1 x Mass of isotope 1) + (Fraction of isotope 2 x Mass of isotope 2) + (Fraction of isotope 3 x Mass of isotope 3) Atomic mass = (0.75 x 133 amu) + (0.20 x 132 amu) + (0.05 x 134 amu) 132.85 amu