Section 5.3 Quantum numbers and Atomic Orbitals

1. Section 5.3Quantum 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. s orbital p orbitals Shapes of 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: Hydrogen Example: Lithium • Example: Helium Example: Oxygen

16. Electron configurations(electron arrangements) • Hund’s rule • Orbitals of equal energy are each occupied by 1 electron before a 2nd 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”