1. Ch. 4.3 Electron Configuration

2. POINT > Define electron configuration POINT > Explain the Aufbau principle POINT > Explain the Pauli exclusion principle POINT > Explain Hund’s rule POINT > Demonstrate 3 methods to indicate electron configuration

3. Review: Recall that electrons in an atom can exist in different main energy levels (principle quantum number) Main energy levels may be divided into sublevels (angular momentum quantum number) A sublevel can contain multiple orbitals (magnetic quantum number) Each orbital can have a maximum of two electrons (spin quantum number)

4. The electron configuration shows where the electrons in an atom are located with respect to the quantum energy levels Three rules guide our understanding of electron configuration…..

5. The Aufbau principle states that electrons occupy the orbitals of lowest energy first Lowest energy levels are (usually) closer to the nucleus

7. Sublevels within an energy level can overlap sublevels of another principal energy level For example, the 4s orbital is a lower energy level than the 3d orbital

9. Which shows the order that electrons fill orbitals? a)1s  2s  3s  4s  2p  3p  4p  3d  b) 4d  4p  4s  3d  3p  3s  2p  2s  c)1s  2s  2p  3s  3p  4s  3d  4p  d)1s  2s  2p  3s  3p  3d  4s  4p 

10. The Pauli exclusion principle states that two electrons cannot have the same four quantum numbers So, two electrons in a single orbital must have opposite spins Spin is noted as arrows or when paired

11. Two electrons cannot have the same four quantum numbers. Example: Two electrons could have main energy level n = 2 and could both be in sublevel p and could both be in orbital p x but they would have opposite spin

12. Hund’s rule states that within a given sublevel, 1) A single electron goes to each orbital before any pairs are formed 2) The single electrons have the same spin

13. Orbital notation of a hydrogen atom Electron configuration notation: 1s 1 1s2s2p x 2p y 2p z 3s

14. Orbital notation of a lithium atom E - config: 1s 2 2s 1 1s2s2p x 2p y 2p z 3s

15.  Orbital notation of a carbon atom: E - config: 1s 2 2s 2 2p 2 1s2s2p x 2p y 2p z 3s

16.  Orbital notation of an oxygen atom: E - config: 1s 2 2s 2 2p 4 1s2s2p x 2p y 2p z 3s

17. Illustrate the orbital notation of a sodium atom E - config: 1s 2 2s 2 2p 6 3s 1 1s2s2p x 2p y 2p z 3s

18. What atom is this? 1s2s2p x 2p y 2p z 3s

19. Show the e - config notation of a sulfur atom: 1s 2 2s 2 2p 6 3s 2 3p 4 Note: the sum of the superscripts = the number of electrons

20. Show the e - config notation of a titanium atom: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2

21. By convention, after determining electron configuration, we group main energy levels together (Theory)1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 (Actual)1s 2 2s 2 2p 6 3s 2 3p 6 3d 2 4s 2

22. Show the e - config notation of a chlorine atom: 1s 2 2s 2 2p 6 3s 2 3p 5

23. What atom is this? 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 Cobalt

24. Noble gas notation is devised to save time: Aluminum:1s 2 2s 2 2p 6 3s 2 3p 1 Neon:1s 2 2s 2 2p 6 Aluminum: [Ne]3s 2 3p 1

25. Noble gas notation is devised to save time: Iron: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 Argon: 1s 2 2s 2 2p 6 3s 2 3p 6 Iron: [Ar]3d 6 4s 2

26. Show the noble gas notation of a titanium atom: [Ar] 4s 2 3d 2

27. What is this atom? [Kr] 5s 2 4d 9

28. Some configurations differ from those predicted by these rules Some atoms prefer filled and half-filled sublevels to other configurations (Copper & Chromium)

29. Ex. Chromium Theoretical: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 Actual: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1

30. Read pages 105-117 Practice #1-2 page 107 Practice #1A, 2-4 Page 115 F.A. #1-5 page 116