Light and Quantized Energy

1. Light and Quantized Energy Section 5.3

2. Ground State Electron Configurations Review from Section #2: According to the quantum mechanical model of the atom… Principle Energy Level (n) Energy Sublevel Orbitals s, p, d, f, … 2 e-s per orbital

3. Ground State Electron Configurations • The arrangement of e-s in an atom is called the atom’s electron configuration. • The ground state is the “normal” electron configuration of the atom…but why? • The atom wants to be in a state that is as low in energy as possible. • I guess that means that atom’s are lazy??? • Atoms are the most stable when they are in their lowest energy state.

4. Ground State Electron Configurations • The Aufbau Principle • Aufbau (in German) = “building up” or “construction • The aufbau principle states that each e occupies the lowest energy orbital available.

5. Ground State Electron Configurations • The Aufbau Principle – Statements: • All orbitals within a specific sublevel are equal in energy. • The energy sublevels contained within a single principle energy level have different energies. • Sublevels in order of increasing energy: s, p, d, f • For related sublevels, the one with the higher principle quantum # is higher in energy.

6. Ground State Electron Configurations • The Aufbau Principle – Statements: • All orbitals within a specific sublevel are equal in energy. • Example: All 2p orbitals are equal in energy • The energy sublevels contained within a single principle energy level have different energies. • For n=2, the p sublevel is higher in energy than the s sublevel • Sublevels in order of increasing energy: s, p, d, f • For related sublevels, the one with the higher principle quantum # is higher in energy. • The 3p sublevel is higher in energy than the 2p sublevel

7. Ground State Electron Configurations • The Pauli Exclusion Principle: • Electrons have a property known as spin…an electron can have an up spin or a down spin. • The Pauli exclusion principle states that only 2 e-s can occupy an orbital at one time and that the e-s must have opposite spin to share the same orbital!

8. Ground State Electron Configurations • Hund’s Rule: • Hund’s “Hungry” Rule • Within an energy sublevel, each orbital must receive one electron before any orbital receives a 2nd electron. Yum, Yum!!!

9. Orbital Diagrams • There are two major ways to represent an atom’s electron configuration: orbital diagrams and electron configuration notation. • In the orbital diagram method, an empty box represents an orbital.

10. Orbital Diagrams Li Be B

11. Orbital Diagrams C N F

12. Orbital Diagrams Si

13. Electron Configuration Notation • The electron configuration method uses the principle energy levels and the sublevel names in writing the notation. The number of electrons in each sublevel is indicated with a superscript above the sublevel symbol. Orbital Diagrams e- configuration notation orbitals sublevels

14. Electron Configuration Notation • How can you decide on the order to fill all of the orbitals? The diagram below can help you! (Hint: Follow the arrows)

15. Electron Configuration Notation Homework: Practice Problems #18, 19, 20 (p.139) Section Assessment #24, 26 (p.141)

16. Electron Configuration Notation • Why are the electron sublevels arranged in a non-numerical order? In particular, look at the last two arrows in the diagram below:

17. Electron Configuration Notation Remember…electrons will stay in the lowest energy orbital that is available to them!!!

18. Electron Configuration Notation

19. Electron Configuration Notation

20. Shorthand Electron Configuration Notation • What is the election configuration notation for… • Argon (Ar) • Potassium (K)

21. Shorthand Electron Configuration Notation • Putting the noble gas in brackets [ ] indicates that the atom of interest contains the same e- configuration up until that point!

22. Electron Configuration Notation • One last note about writing electron configurations… you will not always be able to use the “cheat sheet” that you have.

23. Electron Configuration Notation • But never fear, there is a way for you to get help! You can use…the periodic table!!!

24. Electron Configuration Notation

25. Electron Configuration Notation

26. Valence Electrons • The chemistry of elements only depends on the e-s in the outermost principle energy level. • This level or “shell” of e-s are known as valence electrons. • Using the shorthand e- configuration notation, it is incredibly easy to identify the valence electrons for an atom.

27. Valence Electrons

28. Valence Electrons

29. Valence Electrons • First, let’s get some more practice writing out electron configurations…write out the following elements using both the long and shorthand method: • Copper – Cu • Promethium – Pu • Uranium – U

30. Valence Electrons • An atom can have anywhere between 1 and 8 valence e-s. • Why? Only the outermost s and p orbitals matter when counting valence electrons!!! • How many valence electrons are in the following atoms, according to their electron configurations? • Fr • Br • Sb

31. Electron Dot Structures • As stated earlier, valence e-s are involved in the formation of chemical bonds. Therefore, it is useful to have a visual, shorthand way of representing the valence e-s. • An electron-dot structure is made up of the element’s chemical symbol surrounded by dots that represent the valence electrons. See page 140 for lots of examples!!!

32. Electron Dot Structures • Rule: You must give each side of the chemical symbol one electron before any side gets two. • Electron Dot Structure Examples: Carbon Arsenic Iodine Krypton

33. Electron Configuration Notation

34. Homework (11/4): Practice Problems p. 139: #18d,e (can use shorthand!) 21,22 Practice Problems p. 141: #23a,b,c Section Assessment p. 141: #24,26,28

35. Homework (11/5): Practice Problems p. 141: #23a,b,c,d,f Section Assessment p. 141: #25 (e- configuration only), 28