1. Review: Atomic #  The number of protons in a nucleus Mass #  The number of protons and neutrons in the nucleus Atomic Mass  The average mass of all isotopes of an element (this has units of g/mol or a.m.u.)

2. Notes on Electrons Electrons exist in energy levels surrounding the nucleus. –These are like the rungs of a ladder. Neils Bohr was the first to describe this.

3. Each energy level can only hold a specific number of electrons. Electrons fill the inner energy levels first. Electrons have more energy the further from the nucleus.

4. Electron Capacity = 2n 2

5. The electrons in the energy level furthest from the nucleus are called VALENCE ELECTRONS. Electrons from this level can be gained or lost forming IONS.

6. Group 1—Alkali Metals –1 valence e - Group 2—Alkaline-Earth Metals –2 valence e - Groups 3-12—Transition Metals –Variable valence e - Group 17—Halogens –7 valence e - Group 18—Noble Gases –8 valence e -

7. IONS: –If an atom loses an electron it forms an ion with a positive charge - cation. Why? Because now there are more protons (+) than electrons (-). –If an atom gains an electron it forms an ion with a negative charge - anion. Why? Because now there are more electrons (-) than protons (+)

8. Ion Notation: Li +1 O 2- Cl - How many electrons does each ion have?

9. Summary Atomic # = protons Mass # = protons + neutrons Neutrons = Mass # - atomic # Electrons = protons (if neutral) Electrons = protons - charge