1. Section 5.3 Quantum numbers and Atomic Orbitals Quantum numbers are numbers that specify the properties of atomic orbitals and of the electrons in that orbital It’s the electrons “address”

2. Four Quantum Numbers Principal quantum number Orbital quantum number Magnetic quantum number Spin quantum number

3. Principal quantum number Symbol, n Indicates the main energy levels To this point, only 1-7 Where do we see 7 main energy levels in this room?

4. Orbital quantum number Shape of an orbital Four shapes s, p, d, and f Within each main energy level there are different shapes of orbitals

5. Shapes of orbitals s orbitalp orbitals

6. Shapes of d orbitals

7. Examples of f-shaped orbitals

8. Magnetic quantum number Indicates the orientation (or position) of an orbital around the nucleus –s orbital has 1 orientation –p orbitals have 3 orientations –d orbitals have 5 orientations –f orbitals have 7 orientations Each orbital can contain only 0, 1, or 2 electrons.

9. Spin quantum number Indicates the spin of the electron –+1/2 –-1/2 –So if there are two electrons in one orbital, they spin in opposite directions *** no two electrons can have the same 4 quantum numbers***

10. Electron configurations (electron arrangements) Pauli Exclusion Principle –No two electrons in the same atom will have the same set of 4 quantum numbers

11. How to “read” orbitals How we determine which orbital gets filled with electrons first? Must follow the ________________: –Orbital of Lowest energy gets filled before going to the next lowest energy orbital –In other words we fill from lowest energy to highest energy –“building up” principle: electrons occupy the lowest-energy orbital that is available. –For example, Hydrogen’s electron goes into the __ orbital, because it is the lowest energy orbital

12. Electron configurations (electron arrangements) How do we know which orbitals are higher or lower in energy? –Read Periodic Table from Left to Right, Top to Bottom

13. Periodic Table Sections

14. 3 types of notation Orbital Notation Electron-Configuration Notation Electron Dot Notation

15. Orbital Notation Unoccupied orbital __ Orbital with1 e-↑ or ↓ Orbital with 2 e-↓↑ Example: HydrogenExample: Lithium Example: HeliumExample: Oxygen

16. Electron configurations (electron arrangements) Hund’s rule –Orbitals of equal energy are each occupied by 1 electron before a 2 nd electron is added. –All electrons in singly occupied orbitals must have the same spin –For example, there are 3 p orbitals. If you have 3 electrons, there will be one in each orbital and all will have spin quantum number of +1/2 or -1/2 –Example N:

17. Electron-Configuration notation Similar to orbital notation, but uses superscripts instead of lines Example: Hydrogen Example: Helium Example: Lithium

18. Electron-Dot Notation Uses only the Valence electrons Valence electrons = the electrons in the highest (outermost) main energy level H He K

19. Practice Problems (orbital and dot notation) Carbon Sodium Sulfur

20. Shorthand Notation Use the last noble gas before your element as a “building block” Example: Phosphorous

21. Practice Problems (d and f orbitals) Fe Au

22. Trick to Electron Dot Notation Use the group number that the element is in Hydrogen is in group 1, 1 valence electron Oxygen is in group 6, 6 valence electrons These 8 groups are sometimes called the 8 “main groups”