1. How Atoms Differ

2. a. Properties of Subatomic Particles ParticleSymbolLocationRelative Charge Relative mass Actual mass (g) Electron Proton Neutron

3. b. Atomic Number the number of protons in an atom Identifies element c. Mass Number represents the total number of protons and neutrons in the nucleus A Z atomic number X Mass number

4. d. Isotopes Atoms that have the same number of protons but have a different masses Ex: 3 isotopes of carbon: 12 6 C 13 6 C 14 6 C

5. e. Average Atomic Mass the weighted average of the isotopes of that element. Formula: Atomic mass of an element = ( % abundance of Isotope #1 x mass of Isotope #1 ) + ( % abundance of Isotope #2 x mass of Isotope #2 ) + …

6. Example 1 Silver has two naturally occurring isotopes. Ag-107 has an abundance of 51.82% and mass of 106.9 amu. Ag-109 has a relative abundance of 48.18% and a mass of 108.9 amu. Calculate the average atomic mass of silver.

7. Example 2 Rubidium is a soft, silvery-white metal that has two common isotopes, Rb and Rb. If the abundance of 85 Rb is 72.2% and the abundance of 87 Rb is 27.8%, what is the average atomic mass of rubidium? 85 37 87 37

8. Example 3 Boron has two naturally occurring isotopes. If the abundance of 11 B is 80.10% with an amu of 11.0093, find the abundance of 10 B.