1. John C. Kotz State University of New York, College at Oneonta John C. Kotz Paul M. Treichel John Townsend http://academic.cengage.com/kotz Chapter 7 Atomic Electron Configurations and Chemical Periodicity

2. 2 © 2009 Brooks/Cole - Cengage Important – Read Before Using Slides in Class Instructor: This PowerPoint presentation contains photos and figures from the text, as well as selected animations and videos. For animations and videos to run properly, we recommend that you run this PowerPoint presentation from the PowerLecture disc inserted in your computer. Also, for the mathematical symbols to display properly, you must install the supplied font called “Symb_chm,” supplied as a cross-platform TrueType font in the “Font_for_Lectures” folder in the "Media" folder on this disc. If you prefer to customize the presentation or run it without the PowerLecture disc inserted, the animations and videos will only run properly if you also copy the associated animation and video files for each chapter onto your computer. Follow these steps: 1.Go to the disc drive directory containing the PowerLecture disc, and then to the “Media” folder, and then to the “PowerPoint_Lectures” folder. 2.In the “PowerPoint_Lectures” folder, copy the entire chapter folder to your computer. Chapter folders are named “chapter1”, “chapter2”, etc. Each chapter folder contains the PowerPoint Lecture file as well as the animation and video files. For assistance with installing the fonts or copying the animations and video files, please visit our Technical Support at http://academic.cengage.com/support or call (800) 423-0563. Thank you. http://academic.cengage.com/support

3. 3 © 2009 Brooks/Cole - Cengage ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

4. 4 © 2009 Brooks/Cole - Cengage Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (s) ORBITALS (m s )

5. 5 © 2009 Brooks/Cole - Cengage Each orbital can be assigned no more than 2 electrons! This is tied to the existence of a 4th quantum number, the electron spin quantum number, m s. Arrangement of Electrons in Atoms

6. 6 © 2009 Brooks/Cole - Cengage Electron Spin Quantum Number, m s PLAY MOVIE Can be proved experimentally that electron has an intrinsic property referred to as “spin.” Two spin directions are given by m s where m s = +1/2 and -1/2. Can be proved experimentally that electron has an intrinsic property referred to as “spin.” Two spin directions are given by m s where m s = +1/2 and -1/2.

7. 7 © 2009 Brooks/Cole - Cengage Electron Spin and Magnetism Diamagnetic : NOT attracted to a magnetic fieldDiamagnetic : NOT attracted to a magnetic field Paramagnetic : substance is attracted to a magnetic field.Paramagnetic : substance is attracted to a magnetic field. Substances with unpaired electrons are paramagnetic.Substances with unpaired electrons are paramagnetic. PLAY MOVIE

8. 8 © 2009 Brooks/Cole - Cengage Measuring Paramagnetism Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field See Active Figure 6.18

9. 9 © 2009 Brooks/Cole - Cengage n shell1, 2, 3, 4,... n f shell1, 2, 3, 4,... subshell0, 1, 2,... n - 1 s f subshell0, 1, 2,... n - 1 m orbital -... 0... + m s f orbital - s... 0... + s m s electron spin+1/2 and -1/2 m s f electron spin+1/2 and -1/2 n shell1, 2, 3, 4,... n f shell1, 2, 3, 4,... subshell0, 1, 2,... n - 1 s f subshell0, 1, 2,... n - 1 m orbital -... 0... + m s f orbital - s... 0... + s m s electron spin+1/2 and -1/2 m s f electron spin+1/2 and -1/2 QUANTUM NUMBERS Now there are four!

10. 10 © 2009 Brooks/Cole - Cengage Pauli Exclusion Principle No two electrons in the same atom can have the same set of 4 quantum numbers. That is, each electron has a unique address.

11. 11 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 1, then s = 0 this shell has a single orbital (1s) to which 2e- can be assigned. this shell has a single orbital (1s) to which 2e- can be assigned. When n = 2, then s = 0, 1 2s orbital 2e- 2s orbital 2e- three 2p orbitals6e- three 2p orbitals6e- TOTAL = 8e-

12. 12 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 3, then s = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e- When n = 3, then s = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e-

13. 13 © 2009 Brooks/Cole - Cengage Electrons in Atoms When n = 4, then s = 0, 1, 2, 3 4s orbital 2e- 4s orbital 2e- three 4p orbitals6e- three 4p orbitals6e- five 4d orbitals10e- seven 4f orbitals14e- seven 4f orbitals14e- TOTAL = 32e- And many more!

14. 14 © 2009 Brooks/Cole - Cengage

15. 15 © 2009 Brooks/Cole - Cengage Assigning Electrons to Atoms Electrons generally assigned to orbitals of successively higher energy.Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - C(1/n 2 ). E depends only on n.For H atoms, E = - C(1/n 2 ). E depends only on n. For many-electron atoms, energy depends on both n and s.For many-electron atoms, energy depends on both n and s. See Active Figure 7.1 and Figure 7.2 See Active Figure 7.1 and Figure 7.2 Electrons generally assigned to orbitals of successively higher energy.Electrons generally assigned to orbitals of successively higher energy. For H atoms, E = - C(1/n 2 ). E depends only on n.For H atoms, E = - C(1/n 2 ). E depends only on n. For many-electron atoms, energy depends on both n and s.For many-electron atoms, energy depends on both n and s. See Active Figure 7.1 and Figure 7.2 See Active Figure 7.1 and Figure 7.2

16. 16 © 2009 Brooks/Cole - Cengage Assigning Electrons to Subshells In H atom all subshells of same n have same energy.In H atom all subshells of same n have same energy. In many-electron atom:In many-electron atom: a) subshells increase in energy as value of n + s increases. b) for subshells of same n + s, subshell with lower n is lower in energy. PLAY MOVIE

17. 17 © 2009 Brooks/Cole - Cengage Electron Filling Order See Figure 7.2

18. 18 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3 Explains why E(2s) < E(2p)Explains why E(2s) < E(2p) Z* increases across a period owing to incomplete shielding by inner electrons.Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* = [ Z - (no. inner electrons) ]Estimate Z* = [ Z - (no. inner electrons) ] Charge felt by 2s e- in Li Z* = 3 - 2 = 1Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Be Z* = 4 - 2 = 2Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3and so on!B Z* = 5 - 2 = 3and so on!

19. 19 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge See Figure 7.3 Electron cloud for 1s electrons Z* is the nuclear charge experienced by the outermost electrons.

20. 20 © 2009 Brooks/Cole - Cengage Writing Atomic Electron Configurations 1 1 s value of n value of l no. of electrons spdf notation for H, atomic number = 1 Two ways of writing configs. One is called the spdf notation.

21. 21 © 2009 Brooks/Cole - Cengage Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. One electron has n = 1, s = 0, m s = 0, m s = + 1/2 Other electron has n = 1, s = 0, m s = 0, m s = - 1/2

22. 22 © 2009 Brooks/Cole - Cengage See “Toolbox” in ChemNow for Electron Configuration tool.

23. 23 © 2009 Brooks/Cole - Cengage Electron Configurations and the Periodic Table See Active Figure 7.4

24. 24 © 2009 Brooks/Cole - Cengage LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 f 3 total electrons

25. 25 © 2009 Brooks/Cole - Cengage BerylliumBeryllium Group 2A Atomic number = 4 1s 2 2s 2 f 4 total electrons

26. 26 © 2009 Brooks/Cole - Cengage BoronBoron Group 3A Atomic number = 5 1s 2 2s 2 2p 1 f 5 total electrons 5 total electrons

27. 27 © 2009 Brooks/Cole - Cengage CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 f 6 total electrons 6 total electrons Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

28. 28 © 2009 Brooks/Cole - Cengage NitrogenNitrogen Group 5A Atomic number = 7 1s 2 2s 2 2p 3 f 7 total electrons 7 total electrons

29. 29 © 2009 Brooks/Cole - Cengage OxygenOxygen Group 6A Atomic number = 8 1s 2 2s 2 2p 4 f 8 total electrons 8 total electrons

30. 30 © 2009 Brooks/Cole - Cengage FluorineFluorine Group 7A Atomic number = 9 1s 2 2s 2 2p 5 f 9 total electrons 9 total electrons

31. 31 © 2009 Brooks/Cole - Cengage NeonNeon Group 8A Atomic number = 10 1s 2 2s 2 2p 6 f 10 total electrons 10 total electrons Note that we have reached the end of the 2nd period, and the 2nd shell is full!

32. 32 © 2009 Brooks/Cole - Cengage Electron Configurations of p-Block Elements PLAY MOVIE

33. 33 © 2009 Brooks/Cole - Cengage SodiumSodium Group 1A Atomic number = 11 1s 2 2s 2 2p 6 3s 1 or “neon core” + 3s 1 “neon core” + 3s 1 [Ne] 3s 1 (uses rare gas notation) Note that we have begun a new period. All Group 1A elements have [core]ns 1 configurations.

34. 34 © 2009 Brooks/Cole - Cengage AluminumAluminum Group 3A Atomic number = 13 1s 2 2s 2 2p 6 3s 2 3p 1 [Ne] 3s 2 3p 1 All Group 3A elements have [core] ns 2 np 1 configurations where n is the period number.

35. 35 © 2009 Brooks/Cole - Cengage PhosphorusPhosphorus All Group 5A elements have [core ] ns 2 np 3 configurations where n is the period number. Group 5A Atomic number = 15 1s 2 2s 2 2p 6 3s 2 3p 3 [Ne] 3s 2 3p 3 Yellow P Red P

36. 36 © 2009 Brooks/Cole - Cengage CalciumCalcium Group 2A Atomic number = 20 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s 2 All Group 2A elements have [core]ns 2 configurations where n is the period number.

37. 37 © 2009 Brooks/Cole - Cengage Electron Configurations and the Periodic Table

38. 38 © 2009 Brooks/Cole - Cengage Transition Metals Table 7.4 All 4th period elements have the configuration [argon] ns x (n - 1)d y and so are d-block elements. CopperIronChromium

39. 39 © 2009 Brooks/Cole - Cengage Transition Element Configurations 3d orbitals used for Sc-Zn (Table 7.4)

40. 40 © 2009 Brooks/Cole - Cengage

41. 41 © 2009 Brooks/Cole - Cengage Lanthanides and Actinides All these elements have the configuration [core] ns x (n - 1)d y (n - 2)f z and so are f-block elements. Cerium [Xe] 6s 2 5d 1 4f 1 Uranium [Rn] 7s 2 6d 1 5f 3

42. 42 © 2009 Brooks/Cole - Cengage Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 7.2)

43. 43 © 2009 Brooks/Cole - Cengage

44. 44 © 2009 Brooks/Cole - Cengage Ion Configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]. P [Ne] 3s 2 3p 3 - 3e- f P 3+ [Ne] 3s 2 3p 0

45. 45 © 2009 Brooks/Cole - Cengage Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s 2 3d 6 loses 2 electrons f Fe 2+ [Ar] 4s 0 3d 6 To form cations, always remove electrons of highest n value first!

46. 46 © 2009 Brooks/Cole - Cengage Ion Configurations How do we know the configurations of ions? Determine the magnetic properties of ions. Sample of Fe 2 O 3 Sample of Fe 2 O 3 with strong magnet

47. 47 © 2009 Brooks/Cole - Cengage Ion Configurations How do we know the configurations of ions? Determine the magnetic properties of ions. Ions with UNPAIRED ELECTRONS are PARAMAGNETIC. Without unpaired electrons DIAMAGNETIC. Fe 3+ ions in Fe 2 O 3 have 5 unpaired electrons and make the sample paramagnetic.

48. 48 © 2009 Brooks/Cole - Cengage PERIODIC TRENDS PLAY MOVIE

49. 49 © 2009 Brooks/Cole - Cengage General Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy Electron affinityElectron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.

50. 50 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3Z* is the nuclear charge experienced by the outermost electrons. See Figure 7.3 Explains why E(2s) < E(2p)Explains why E(2s) < E(2p) Z* increases across a period owing to incomplete shielding by inner electrons.Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* = [ Z - (no. inner electrons) ]Estimate Z* = [ Z - (no. inner electrons) ] Charge felt by 2s e- in Li Z* = 3 - 2 = 1Charge felt by 2s e- in Li Z* = 3 - 2 = 1 Be Z* = 4 - 2 = 2Be Z* = 4 - 2 = 2 B Z* = 5 - 2 = 3and so on!B Z* = 5 - 2 = 3and so on!

51. 51 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge See Figure 7.3 Electron cloud for 1s electrons Z* is the nuclear charge experienced by the outermost electrons.

52. 52 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge Z* The 2s electron PENETRATES the region occupied by the 1s electron. 2s electron experiences a higher positive charge than expected. PLAY MOVIE

53. 53 © 2009 Brooks/Cole - Cengage Effective Nuclear Charge, Z* AtomZ* Experienced by Electrons in Valence Orbitals Li+1.28 Be------- B+2.58 C+3.22 N+3.85 O+4.49 F+5.13 Increase in Z* across a period [Values calculated using Slater’s Rules]

54. 54 © 2009 Brooks/Cole - Cengage Orbital Energies ChemNow Screens 8.9 - 8.13, Simulations Orbital energies “drop” as Z* increases

55. 55 © 2009 Brooks/Cole - Cengage General Periodic Trends Atomic and ionic sizeAtomic and ionic size Ionization energyIonization energy Electron affinityElectron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly.

56. 56 © 2009 Brooks/Cole - Cengage Atomic Radii See Active Figure 7.8

57. 57 © 2009 Brooks/Cole - Cengage Atomic Size Size goes UP on going down a group. See Figure 7.8.Size goes UP on going down a group. See Figure 7.8. Because electrons are added further from the nucleus, there is less attraction.Because electrons are added further from the nucleus, there is less attraction. Size goes DOWN on going across a period.Size goes DOWN on going across a period. Size goes UP on going down a group. See Figure 7.8.Size goes UP on going down a group. See Figure 7.8. Because electrons are added further from the nucleus, there is less attraction.Because electrons are added further from the nucleus, there is less attraction. Size goes DOWN on going across a period.Size goes DOWN on going across a period.

58. 58 © 2009 Brooks/Cole - Cengage Atomic Size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge. Large Small Increase in Z*

59. 59 © 2009 Brooks/Cole - Cengage Trends in Atomic Size See Active Figure 7.8

60. 60 © 2009 Brooks/Cole - Cengage Sizes of Transition Elements See Figure 7.9

61. 61 © 2009 Brooks/Cole - Cengage Sizes of Transition Elements See Figure 7.9 3d subshell is inside the 4s subshell.3d subshell is inside the 4s subshell. 4s electrons feel a more or less constant Z*.4s electrons feel a more or less constant Z*. Sizes stay about the same and chemistries are similar!Sizes stay about the same and chemistries are similar!

62. 62 © 2009 Brooks/Cole - Cengage Density of Transition Metals 6th period 5th period 4th period

63. 63 © 2009 Brooks/Cole - Cengage Ion Sizes Does the size go up or down when losing an electron to form a cation? Does the size go up or down when losing an electron to form a cation?

64. 64 © 2009 Brooks/Cole - Cengage Ion Sizes CATIONS are SMALLER than the atoms from which they come.CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES.The electron/proton attraction has gone UP and so size DECREASES. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p + Forming a cation.

65. 65 © 2009 Brooks/Cole - Cengage Ion Sizes Does the size go up or down when gaining an electron to form an anion?

66. 66 © 2009 Brooks/Cole - Cengage Ion Sizes ANIONS are LARGER than the atoms from which they come.ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES.The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes.Trends in ion sizes are the same as atom sizes. Forming an anion. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

67. 67 © 2009 Brooks/Cole - Cengage Trends in Ion Sizes See Active Figure 7.12

68. 68 © 2009 Brooks/Cole - Cengage Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons? Why do metals lose electrons in their reactions? Why does Mg form Mg 2+ ions and not Mg 3+ ? Why do nonmetals take on electrons?

69. 69 © 2009 Brooks/Cole - Cengage Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ f Mg + (g) + e- PLAY MOVIE

70. 70 © 2009 Brooks/Cole - Cengage Mg (g) + 738 kJ f Mg + (g) + e- Mg + (g) + 1451 kJ f Mg 2+ (g) + e- Mg + has 12 protons and only 11 electrons. Therefore, IE for Mg + > Mg. IE = energy required to remove an electron from an atom in the gas phase. Ionization Energy PLAY MOVIE

71. 71 © 2009 Brooks/Cole - Cengage Mg (g) + 735 kJ f Mg + (g) + e- Mg + (g) + 1451 kJ f Mg 2+ (g) + e- Mg 2+ (g) + 7733 kJ f Mg 3+ (g) + e- Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no. Ionization Energy PLAY MOVIE

72. 72 © 2009 Brooks/Cole - Cengage 2nd IE / 1st IE Li Na K B Al

73. 73 © 2009 Brooks/Cole - Cengage Trends in Ionization Energy See Active Figure 7.10

74. 74 © 2009 Brooks/Cole - Cengage Trends in Ionization Energy

75. 75 © 2009 Brooks/Cole - Cengage Orbital Energies CD-ROM Screens 8.9 - 8.13, Simulations As Z* increases, orbital energies “drop” and IE increases.

76. 76 © 2009 Brooks/Cole - Cengage Trends in Ionization Energy IE increases across a period because Z* increases.IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals.Metals lose electrons more easily than nonmetals. Metals are good reducing agents.Metals are good reducing agents. Nonmetals lose electrons with difficulty.Nonmetals lose electrons with difficulty.

77. 77 © 2009 Brooks/Cole - Cengage Trends in Ionization Energy IE decreases down a groupIE decreases down a group Because size increases.Because size increases. Reducing ability generally increases down the periodic table.Reducing ability generally increases down the periodic table. See reactions of Li, Na, KSee reactions of Li, Na, K

78. 78 © 2009 Brooks/Cole - Cengage Periodic Trend in the Reactivity of Alkali Metals with Water Lithium SodiumPotassium PLAY MOVIE

79. 79 © 2009 Brooks/Cole - Cengage Electron Affinity A few elements GAIN electrons to form anions. Electron affinity is the energy involved when an atom gains an electron to form an anion. A(g) + e- f A - (g) E.A. = ∆U A(g) + e- f A - (g) E.A. = ∆U

80. 80 © 2009 Brooks/Cole - Cengage Electron Affinity of Oxygen ∆U is EXOthermic because O has an affinity for an e-. [He]      O atom EA = - 141 kJ + electron O [He]       - ion

81. 81 © 2009 Brooks/Cole - Cengage Electron Affinity of Nitrogen ∆U is zero for N - due to electron- electron repulsions. EA = 0 kJ [He]     Natom  [He]    N - ion  + electron

82. 82 © 2009 Brooks/Cole - Cengage Trends in Electron Affinity See Active Figure 7.11

83. 83 © 2009 Brooks/Cole - Cengage See Figure 7.11 and Appendix FSee Figure 7.11 and Appendix F Affinity for electron increases across a period (EA becomes more positive).Affinity for electron increases across a period (EA becomes more positive). Affinity decreases down a group (EA becomes less positive).Affinity decreases down a group (EA becomes less positive). Atom EA F+328 kJ Cl+349 kJ Br+325 kJ I+295 kJ Atom EA F+328 kJ Cl+349 kJ Br+325 kJ I+295 kJ Trends in Electron Affinity Note effect of atom size on F vs. Cl