1. DO NOW How many electrons are in Beryllium? Magnesium? Calcium? Strontium? What is similar and what is different about each of these elements? Is there a pattern?

2. Chapter 5 Electrons in atoms

3. Objectives 5.1a Describe what Bohr proposed in his model of the atom. 5.1b Describe what the quantum mechanical model determines about the electrons in an atom. 5.1c Explain how sublevels of principal energy levels differ. Electrons!

4. Back to Rutherford Nucleus containing protons and neutrons Electrons orbit around the outside of the nucleus

5. A new view: Niels Bohr Neils Bohr (1885-1962) was a Danish Physicist who studied under Rutherford developed the Bohr Model ▫Introduction of the quantum in atomic theory ▫http://science.sbcc.edu/ physics/flash/siliconsol arcell/bohratom.swfhttp://science.sbcc.edu/ physics/flash/siliconsol arcell/bohratom.swf

6. Let’s explain it musically- donna lee vs. reveille https://www.youtube.com/watch?v=QdD6XE9DtAE https://www.youtube.com/watch?v=83ozOX9l7M8

7. The Bohr Model Bohr proposed that electrons are found in circular paths (orbits) around a nucleus These orbits are called energy levels To move from one energy level to another, you need a quantum of energy

9. The quantum mechanical model The quantum mechanical model of the atom was developed from the equations of Erwin Schrödinger There is an “electron cloud” where atoms are constantly in motion Area of high electron probability

10. Schrödinger continued Energy levels have sublevels. ▫n=1 has 1 sublevel (1s) ▫n=2 has 2 sublevels (2s, 2p) Atomic orbitals are regions of space where there is high probability of finding an electron ▫Each orbital can hold 2 electrons

11. Atomic Orbitals (p. 131-132)

12. Animation http://www.shsu.edu/~chm_tgc/BbAIF/PDBs/ applet/PDBorbitals.htmlhttp://www.shsu.edu/~chm_tgc/BbAIF/PDBs/ applet/PDBorbitals.html

13. Summary of Principal Energy Levels and Sublevels Principal Energy Level Number of Sublevels Type of SublevelMaximum number of electrons n=1 n=2 n=3 n=4

14. Recap Energy Levels (n=1, 2, 3… etc) Sublevels (s, p, d, f) Orbitals (px, py, pz) electron

15. Atomic Orbitals (p. 131-132) How many orbitals are in energy level 1?  (1) ▫Level 2?  (4) ▫Level 3?  (9) ▫Level 4?  (16) Do you see a pattern? ▫Level 7?  (49)

16. Homework Read 5.1 Lesson Check, p. 132 questions 1-7

17. Do-now How many electrons does Silicon have? In which principal energy levels (n=?) do the electrons go?

18. Objectives 5.2a List the three rules for writing the electron configuration of elements.

19. Before we begin, What did we say an orbital is? ▫A region of space where there is a high probability of finding an electron. OK, we’re good then. Electrons found here Not here

20. Electron Configuration In order to determine how an atom will interact with other atoms, we need to know where the electrons are going to be. The electron configurations of an atom are the ways in which the electrons are arranged in various orbitals around the nuclei of the atoms

21. The three rules 1.Aufbau Principle 2.Pauli Exclusion Principle 3.Hund’s Rule

22. The Aufbau Principle-p. 135 The Aufbau principle: electrons occupy the orbital of the lowest energy levels first

23. The d sublevel The 3d sublevel actually has a higher energy than the 4s sublevel

24. The Pauli Exclusion principle The Pauli exclusion principle states that an orbital may only describe two electrons Electrons must have opposite spins, which can be thought of as clockwise or counterclockwise e-e- e-e- e-e- e-e-

25. Hund’s rule Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. In other words, fill one electron in each orbital for the sublevel before you double up

26. Lithium Beryllium Fluorine Periodic Table Practice

28. As partners- p. 135 Complete electron configurations for elements 12-19 (Magnesium to Potassium) We will review everybody’s list as a class. Leave plenty of space in between each configuration!

29. Shorthand Instead of drawing out each electron configuration, there is a shorthand used by most chemists Oxygen becomes: ▫1s22s22p4▫1s22s22p4 What about Bromine? ▫1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5 ▫(The 3d sublevel is written before the 4s sublevel, even though it is at a higher energy level)

30. Now complete each configuration in shorthand!

31. Recap 1. What are the three rules that govern the filling of atomic orbitals by electrons? ▫Aufbau Principle, Pauli Exclusion Principle, and Hund’s rule

32. Homework Worksheet- Electron Configurations Be ready for a quiz on Electron Configurations tomorrow Be ready for a lab tomorrow

33. Do-now Provide the electron configuration for Neon What would it look like if an electron jumped from the 2p orbital to the 3p orbital?

34. Objectives 5.3a Explain what causes atomic emission spectra. 5.3b Explain how the frequencies of emitted light are related to changes in electron energies.

35. Color Friday you saw that different chemicals have different colors when heated. ▫Why did different chemicals give off different colors?

36. What is light? Light consists of waves Draw another wave that has twice the amplitude of this wave.

37. Waves The amount of wave cycles in a given amount of time is called frequency (ν). ▫The amount of cycles in one second is called hertz (1/s) Draw another wave that has twice the frequency of the wave you have already drawn

38. The speed of light The wavelength times frequency gives us the speed of a wave ▫The speed of light is usually referred to as c ▫c= 2.998 x 10^8 m/s c=λν

39. Problem What is the frequency of radiation which has a wavelength of 4.00 x 10^-5 m?  7.49 x 10^12 hz

40. The electromagnetic spectrum includes the following types of radiation:

41. demo

42. Do-now What is the wavelength of electromagnetic radiation with a frequency of 5.31 x 10 7 /s? 5.65 m

43. Atomic Spectra When electrons in an atom are absorbed into a high energy level, they release the energy as light. Each atom has a unique color that is made of several different wavelengths. This is known as the atomic emission spectrum for that element

44. Neon

45. How “neon” signs are made

46. Atomic Spectra explained (p. 144-145) The ground state is the lowest possible energy level for an atom. Electrons absorb energy, promoting them to a higher energy level Electrons return to ground state, releasing a quanta of light. ▫This is called a photon.

47. Match the spectra Helium Mercury Neon Oxygen

48. Planck’s Constant To determine the amount of energy in a photon, you must use Planck’s Constant ▫h= 6.626x10^-34 Js E=hν

49. Problem What is the energy of a photon of light with a frequency of 4.30 x 10^4 hz? ▫2.85 x 10^-29 J

50. Problem A photon of light has a wavelength of 6.45 x 10^-5 cm. ▫What is the value of this wavelength in meters?  6.45 x 10^-7 m ▫Using Figure 5.8 on p. 139, guess which color of light this photon belongs to.  Red What is the frequency of this light? ▫4.65 x 10^14 hz What is the energy of this light? ▫3.08 x 10-19 J

51. De Broglie equation Louis de Broglie, a French Physicist, hypothesized that if light can behave both as a wave (has wavelength, frequency), and a particle (is quantized), then so can matter All objects can have a wavelength Mass must be very small

52. De Broglie Waves

53. The Heisenberg uncertainty principle Werner Heisenberg discovered that it is impossible to predict the position and velocity of a particle at the same time

54. Homework Read Section 5.3, 138-145 Complete Lesson Check, p 148 ▫Don’t do #20 or 22 P. 153, #63 Test Tuesday! Quiz Tomorrow

55. Do-now What is the energy of a photon that has a wavelength of 1.4 x 10 -8 m? 1.4 x 10 -17 J

57. Uncertainty

59. Homework Albert Einstein once famously said “God does not roll dice with the universe” when he heard about the uncertainty principle. ▫Write Einstein a persuasive essay as if you are either arguing for or against Heisenberg’s principle. ▫Due on day of Test- Wednesday!