1. Catalyst 1. 2. 3.

2. Ways to Organize Shapes Organize a sphere on xyz coordinate plane Organize dumb bells on xyz coordinate plane

3. Today’s Learning Targets 1.6 – I can characterize an electron based on its 4 quantum numbers (n, l, m l, and m s ). I can explain what each of these numbers indicate and discuss the importance of these numbers. 1.7 – I can describe the shape, number, and energy level of the s, p, d, and f orbitals. Furthermore, I can draw the s and p orbitals. 1.8 – I can write the electron configuration and orbital diagram for any element on the Periodic Table using the Pauli Exclusion Principle and Hund’s Rule.

4. Orbitals The solution to the Schrödinger equation gives us orbitals for a particular element. Every orbital has a special shape and energy level. We keep the idea of energy level (n = 1,2,3, etc.) from the Bohr model, but remove idea of fixed orbits. There are 4 characteristic shapes that an orbital can take.

5. s Orbitals Spherical shape Only one possible orientation around the nucleus. Seen whenever n (energy level) is 1 or greater All s orbitals have only one possible orientation.

6. p Orbitals All p-orbitals are dumb-bell shaped. Because of the unique shape there are 3 possible orientations around the nucleus (one along each of the 3 axis). Seen whenever n is 2 or greater

7. d Orbitals All d orbitals have a four-leaf clover shape. 4 possible orientations within the plane and one odd shape (d z 2 ) Appear when n is 3 or higher.

8. f Orbitals These orbitals appear whenever n is 4 or higher Only seen in elements that have many electrons.

9. Quick Talk 1 partner explain s and d orbitals. 1 partner explain p and f orbitals. Be able to draw s and p orbital

10. Quantum Numbers We can describe any orbital that an electron exists in using 4 quantum numbers. No 2 orbitals can have the same 4 quantum numbers. Quantum Numbers: 1. Principal Quantum Number (n) 2. Angular Momentum Quantum Number (l) 3. Magnetic Quantum Number (m l ) 4. Spin Magnetic Quantum Number (m s )

11. Principal Quantum Number (n) Describes the energy level of the orbital Can have any integer value of 1 or greater (1, 2, 3, etc.) The bigger n becomes, the higher the energy level

12. Angular Momentum ( ) Describes the shape of our orbital (s, p, d, or f) Can have any value from 0 to (n – 1) Each orbital has an assigned angular momentum: Value of0123 Letter usedspdf

13. Quick Write If I have an element at the n = 3 energy level, then what type(s) of orbitals do I have?

14. Magnetic Quantum Number (m ) m l is any number – to + Every orbital can take on a certain number of “allowed” orientations. Tells you the number of total electrons an orbital can hold The number of allowed orientations is the sum of – to +

15. Quick Write How many allowed orientations are there for the s orbital ( =0)? How many electrons can it hold? How many allowed orientations are there for the d orbital ( =2)? How many electrons can it hold?

16. Spin Magnetic Quantum Number (m s ) Any orbital can contain, at most, 2 electrons If electrons are in the same orbital, then they must have opposite spins or m s values An electron can either have a spin of +½ or -½

17. Summarize

18. Around the World Around the room there are 10 problems. Cycle through the problems to practice writing electron configurations. Complete all problems

19. 5 Minute Break

20. Who would perform at the ultimate concert?

21. Concert of a Lifetime

22. Stage

23. Orbitals and their Energy Orbitals, based on many different reasons, have varying energies. n does not determine energy levels when comparing orbitals

24. Electron Configurations Electron configurations describe the distribution of each electron among the various orbitals in the atom

25. Pauli Exclusion Principle Within an atom, no two electrons can have the same set of 4 quantum numbers.

26. Orbital energy diagrams Aufbau principle: Build up each atom from the preceding atom by “filling” electrons in from the bottom.

27. Hund’s Rule When placing electrons into a group of similar orbitals, electrons enter empty orbitals first before they form pairs

28. Class Example Draw the electron configuration for fluorine.

29. Table Talk Draw the electron configuration for magnesium

30. 3 Essential Principles 1. Pauli Exclusion Principle – Every electron gets its own unique quantum number per element. 2. Hund’s Rule – Electrons spread out within the same energy level. 3. Aufbau Principle – Fill the lowest energy orbital first

31. How to Remember Order of Orbital Filling 1s 2s 3s 4s 5s 6s 2p 3p 4p 5p 6p 3d 4d 5d 6d 4f 5f 6f

32. Periodic Table Quantum Shortcut! The position of the element determines the last orbital and energy level filled.

33. Condensed Electron Configurations Abbreviate configurations using Noble Gases. Choose the Noble Gas in the row just above the element of interest. Substitute this Noble Gas for all electrons prior to the row the element is in. 1s 2 2s 2 2p 6 3s 2 3p 4 [Ne]3s 2 3p 4 10 electrons = Neon!

34. Class Example Write the electron configuration for zinc using condensed notation

35. Table Talk Write the electron configuration for Tin.

36. JTPS: Analyze the Configuration Which configuration for chromium is more stable. WHY?: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 or

37. Strange Electron Configurations Due to the closeness in energy of the 3d and 4s orbitals, electrons will spread out in order to satisfy Hund’s Rule. Seen most often in chromium and copper.

38. Summarize

39. Coach and Correct

40. Question 1 In 3 words or less for each: o Describe the Pauli Exclusion Principle o Describe Hund’s Rule o Describe the Aufbau Principle

41. Question 2 What is the principal quantum number, number of electrons, and angular momentum of 5p 6

42. Question 3 Write the long-form electron configuration for sulfur

43. Question 4 Write the condensed electron configuration for titanium

44. Question 5 Why do electrons fill the lowest energy level first?

45. Question 6 Looking at the following electron configurations, which represents a chemically unreactive element? Justify your answer! a) 1s____2s ↑ _ b) 1s ↑↓ 2s ↑↓_ c) 1s ↑↓ 2s ↑↓ 2p ↑ ↑ ____ d) 1s ↑↓ 2s ↑↓ 2p ↑↓_ ↑↓ _↑↓_ e) [Ar]4s ↑↓ 3d ↑↓ ↑ ↑ ↑ ↑_

46. Question 7 The contour representation of one of the orbitals for the n = 3 shell of hydrogen atom is shown below a) What is the quantum number l for this orbital? b) What is the notation for this orbital

47. Question 8 An element has a valence shell configuration of ns 2 np 5, what element(s) is it?

48. Question 9 You friend writes the electron configuration for phosphorus as: [He]3s 2 3p 3. What is wrong with the way that they wrote the configuration?

49. Question 10 Write the condensed electron configuration for uranium

51. Closing Time You should have all Chapter 6 reading and homework problems done by Monday/Tuesday to be on track. Also, read corresponding chapter in Cracking the AP Chemistry Exam Pre-lab for Lab 2 due at start of class Monday/Tuesday. No pre-lab = you cannot start lab on time = staying after school to finish the lab.