1. Chemistry 11 Resource: Chang’s Chemistry, Chapter 7

2. Objectives 1. Explain how the lines in the emission spectrum of hydrogen are related to the electron energy levels. 2. State the relative energies of s, p, d, and f orbitals in a single energy level. 3. State the maximum number of orbitals in a given energy level. 4. Draw the shape of s and p orbitals. 5. Apply the Aufbau principle, Hund’s rule, and the Pauli exclusion principle to write electronic configurations for atoms and ions up to Z = 20.

3. Activities  Exercises from the text  Quizzes  3-d models of atomic orbitals

4. Bohr’s model  Ever since the 17 th century, the phenomenon of emission spectra has fascinated physicists.  The emission spectrum of a substance can be seen by energizing a sample of material.

5. Bohr’s model Emission by a heated object p 258 p 267

6. Bohr’s model  The emission spectra of gases are quite different.  Gases were found to emit light only at certain wavelengths.

7. Bohr’s theory Emission spectra of gases p 268

8. Bohr’s theory  What was the model of the atom before Bohr?  Could that model possibly explain the emission spectra phenomenon?

9. Bohr’s theory  Before Bohr, physicists knew that the atom consisted of protons and electrons.  They believed that the electrons moved around the nucleus in circular orbits (Rutherford’s model).  Why was this model acceptable to scientists?

10. Bohr’s model  In the early 20 th century, Bohr added to the contemporary model of the atom: The single electron in the hydrogen atom can only be located in certain orbits. Each orbit has a particular energy associated with it.

11. Bohr’s model Bohr’s model of the atom p 269

12. Bohr’s model  Only certain orbits are permitted.  Each orbit has an associated energy value.  Therefore, the energy associated with e- motion is quantized, or fixed in value.

13. Bohr’s model  Bohr attributed the emission spectrum of hydrogen to the following process: The electron absorbs energy and jumps to a higher orbit. When the electron returns to its ground (normal) state, it emits energy through a photon (light particle).  Since only certain orbits (energy levels) are permitted, light at a certain wavelength is emitted.

14. Bohr’s model Emission of light by a hydrogen atom p 269

15. Electron cloud model  Bohr’s model could not account for the emission spectra of atoms with more than one electron.  It became even more insufficient when physicists discovered that electrons are wavelike.  How can you pinpoint the location of an electron if it is a wave?

16. Electron cloud model  Heisenberg’s uncertainty principle: It is impossible to know [the momentum p and] the position of a particle with certainty.  How does this principle defy Bohr’s model of the atom?

17. Electron cloud model  In the 1920s, Schrödinger applied this to the model of the atom: The exact location of an electron cannot be pinpointed.  Therefore, the representation of the electron was modified from lines to a cloud where an electron is more likely to be found.

18. Electron cloud model The electron cloud model p 278

19. Electron cloud model  Schrödinger ushered in a new age of physics called quantum mechanics.  We now refer to the “location” of electrons as atomic orbitals.  Each atomic orbital has a certain associated energy and a distribution of electron density.

20. Quantum numbers  As a result of the discoveries in the 1920s, electrons were assigned quantum numbers to describe their distribution or “location”.  Three quantum numbers are required to describe the distribution of electrons. the principal quantum number n the angular momentum quantum number l the magnetic quantum number m l

21. Quantum numbers  The principal quantum number n is designated an integer value greater than 0, i.e. 1, 2, 3, 4, …  It relates to the average distance of the e- from the nucleus.  The larger n is, the farther away it is from the nucleus.  If n is larger, is the orbital bigger or smaller?

22. Quantum numbers  The angular momentum quantum number l tells us the “shape” of the orbital.  l is related to n The values of l can vary from 0 to (n -1).  If n = 1, what are the possible values of l?  What if n = 3?

23. Quantum numbers  The value of l is generally designated by the letters s, p, d, … as follows: l012345 Name of orbitalspdfgh

24. Quantum numbers If an e- has a principal quantum number of 1 (n = 1), how many orbitals are possible?

25. Quantum numbers Since n = 1, the only possible value of l is 0. remember: l varies from 0 to n – 1 since n – 1 = 0, 0 is the only possible l value therefore: there is only 1 orbital when n = 1. This is called the 1s orbital. l012345 Name of orbitalspdfgh

26. Quantum numbers If an e- has a principal quantum number of 2, how many orbitals are possible?

27. Quantum numbers If n = 2, l can be 0 and 1 therefore: TWO orbitals are possible. These orbitals are called 2s and 2p. l012345 Name of orbitalspdfgh

28. Quantum numbers  A group of orbitals that have the same value for n (e.g. 2s and 2p) are frequently called a shell.

29. Quantum numbers  The magnetic quantum number m l describes the orbital’s orientation in space.  The value of m l depends on l and varies as follows: -l, (-l +1), … 0, … (l - 1), l

30. Quantum numbers If n = 2 and l = 1, how many orbitals are possible?

31. Quantum numbers Three orbitals in that subshell are possible: since l = 1, m l = -1, 0, 1 Therefore: 3 orbitals are possible. These orbitals are called 2p x, 2p y, and 2p z. This will all make a little more sense later on

32. Quantum numbers Relation between quantum numbers and atomic orbitals nlmlml Number of orbitals Atomic orbital designations 10011s 2 3

33. Quantum numbers Relation between quantum numbers and atomic orbitals nlmlml Number of orbitals Atomic orbital designations 10011s 2 0101 0 -1, 0, -1 1313 2s 2p x, 2p y, 2p z 3 012012 0 -1, 0, -1 -2, -1, 0, 1, 2 135135 3s 3p x, 3p y,, 3p z 3d xy, 3d yz, 3d xz, 3d x2- y2, 3d z2

34. Quantum numbers  A fourth quantum number m s is used to denote the spin of the electron.  Electrons are known to spin two ways: up or down. This electron spin quantum number will be discussed later on.

35. Atomic orbitals  Both Bohr and Schrödinger made significant contributions to our understanding of the atom.  We will use their ideas to get a better picture of atomic structure.

36. Atomic orbitals  In principle, an electron can be found anywhere in the atom.  In a typical hydrogen atom, where would the single electron most likely be?

37. Atomic orbitals  Common sense dictates that the single electron will probably be close to the nucleus.  Thus we can represent the 1s orbital by drawing a boundary that encloses about 90% of the total electron density:  p 282

38. Atomic orbitals  Recall that each value of n has an s orbital (1s, 2s, 3s, …) The shape of the s orbital p 282 How does the value of n affect the shape/size of the orbital?

39. Atomic orbitals At what value for n do we see s orbitals?

40. Atomic orbitals There is an s orbital at every value of n. Think of it as the “basic” orbital.

41. Atomic orbitals If n = 1, does a p orbital (l = 1) exist?

42. Atomic orbitals No. p orbitals exist when n = 2 or higher: when n = 1, l = 0; therefore only the 1s is possible. p orbitals are associated with l = 1.

43. Atomic orbitals  p orbitals appear when n is 2 or higher: since n = 2, l = 0, 1 These correspond to the 2s and the 2p orbitals.

44. Atomic orbitals  Furthermore, when n = 2 and l = 1 (2p) m l = -1, 0, 1  Therefore there are THREE possible 2p orbitals.

45. Atomic orbitals Shape of the 2p orbitals p 283

46. Atomic orbitals The shape of the 3d orbitals p 283

47. Atomic orbitals  Remember that each orbital has a shape (cloud) and a certain energy associated with it.  Which orbital has the lowest energy associated with it?

48. Atomic orbitals The 1s orbital has the lowest energy.

49. Atomic orbitals Orbital energy levels p 285

50. Atomic orbitals The order in which atomic subshells are filled p 285

51. Electron configuration  An electron can be identified by its four quantum numbers.  You may think of the quantum numbers as the “address” of the e- because they describe its location.

52. Electron configuration Summary of quantum numbers Quantum number Information about e- Part of the “address” ndistance from nucleusprovince / state lshapecity mlml orientation in spacestreet msms spinnumber

53. Electron configuration What are the four quantum numbers of hydrogen’s single electron? (n, l, m l, m s )

54. Electron configuration (n, l, m l, m s ) (1, 0, 0, +½) or (1, 0, 0, -½)

55. Electron configuration Write the four quantum numbers of an electron in the 3p orbital.

56. Electron configuration  Homework p 299: 7.55-7.61, odd; 7.64  Review for a quiz next class

57. Electron configuration  The electron configuration of an atom is how the electrons are distributed among the various atomic orbitals.  This is the electron configuration of hydrogen which has 1 e-. 1s11s1 Denotes the principal quantum number n Denotes the angular momentum quantum number l Denotes the number of electrons in the orbital or subshell

58. Electron configuration  Electron configuration can also be represented by an orbital diagram that shows the spin of the electron: 1s11s1

59. Pauli exclusion principle  No two electrons in an atom can have the same four quantum numbers.  If they are in the same orbital (i.e. same values for n, l, and m l ) then they must have different values for m s. 1s21s2 1s21s2 1s21s2

60. Hund’s rule  The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (same spins).

61. Hund’s rule  Carbon (Z = 6) is 1s 2 2s 2 2p 2 Which configuration satisfies Hund’s rule? 1s21s2 2p 2 2s22s2 1s21s2 2s22s2 1s21s2 2s22s2

62. Aufbau principle  “Aufbau” is the German word for “building up”  Just as protons are added one-by-one to build up the elements, so are electrons into the atomic orbitals.  This introduces a different way of showing of electronic configuration.

63. Aufbau principle  The configuration shows the noble gas (in brackets) that most nearly precedes the element being considered. So instead of: Na 1s 2 2s 2 2p 6 3s 1 You may represent Na as: Na [Ne]3s 1  Very convenient.

64. The transition metals  The electronic configurations of elements from Z = 1 to Z = 20 are relatively straightforward.  The electronic configurations of the transition metals have “strange” electronic configurations that do not necessarily follow convention (p292)  Why do you think this is?

65. Electronic configuration  Homework pp 299 – 302 ○ # 7.71, 72, 79. 85. 87. 124  Quiz next class  Bring old newspapers

66. Atomic orbital models  You will divide yourselves into THREE groups of fairly equal numbers.  Each group has a different assignment.  This will count as a project.  It is due on 27 January 2009.

67. Atomic orbital models  Group 1 Build models for the 1s, 2s, and 3s orbitals. Create 3 posters: ○ Orbital energy levels ○ Order in which atomic subshells are filled ○ Pauli exclusion principle and Hund’s rule  Group 2 Build models for the 2p and 3p orbitals.  Group 3 Build models for the 3d orbitals.

68. Atomic orbital models  Color codes: Blue: n = 1 Yellow: n = 2 Red: n = 3

69. Atomic orbital models Grading the orbital models (for each model) Correct shape5 pts Correct size (relative to similar orbitals with different values for n) 5 Correct color2 Stability3 Total15

70. Atomic orbital models Grading the posters (for each poster) Accurate information5 pts Pertinent details present5 Organization of information3 Creativity / aesthetic value2 Total15