1. 1 © 2006 Brooks/Cole - Thomson Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver CHAPTER 8 Atomic Electron Configurations and Chemical Periodicity © 2006 Brooks/Cole Thomson Lectures written by John Kotz

2. 2 © 2006 Brooks/Cole - Thomson ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

3. 3 © 2006 Brooks/Cole - Thomson Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (m l )

4. 4 © 2006 Brooks/Cole - Thomson 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

5. 5 © 2006 Brooks/Cole - Thomson Electron Spin Quantum Number, m s Can be proved experimentally that electron has a spin. Two spin directions are given by m s where m s = +1/2 and -1/2. Can be proved experimentally that electron has a spin. Two spin directions are given by m s where m s = +1/2 and -1/2.

6. 6 © 2006 Brooks/Cole - Thomson 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.

7. 7 © 2006 Brooks/Cole - Thomson Measuring Paramagnetism Paramagnetic : substance is attracted to a magnetic field. Substance has unpaired electrons. Diamagnetic : NOT attracted to a magnetic field Active Figure 8.2

8. 8 © 2006 Brooks/Cole - Thomson n ---> shell1, 2, 3, 4,... l ---> subshell0, 1, 2,... n - 1 m l ---> orbital -l... 0... +l m s ---> electron spin+1/2 and -1/2 n ---> shell1, 2, 3, 4,... l ---> subshell0, 1, 2,... n - 1 m l ---> orbital -l... 0... +l m s ---> electron spin+1/2 and -1/2 QUANTUM NUMBERS Now there are four!

9. 9 © 2006 Brooks/Cole - Thomson 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.

10. 10 © 2006 Brooks/Cole - Thomson Electrons in Atoms When n = 1, then l = 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 l = 0, 1 2s orbital 2e- 2s orbital 2e- three 2p orbitals6e- three 2p orbitals6e- TOTAL = 8e-

11. 11 © 2006 Brooks/Cole - Thomson Electrons in Atoms When n = 3, then l = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e- When n = 3, then l = 0, 1, 2 3s orbital 2e- 3s orbital 2e- three 3p orbitals6e- three 3p orbitals6e- five 3d orbitals10e- TOTAL = 18e-

12. 12 © 2006 Brooks/Cole - Thomson Electrons in Atoms When n = 4, then l = 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!

13. 13 © 2006 Brooks/Cole - Thomson

14. 14 © 2006 Brooks/Cole - Thomson 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 l.For many-electron atoms, energy depends on both n and l. See Active Figure 8.4, Figure 8.5, and Screen 8. 7. See Active Figure 8.4, Figure 8.5, and Screen 8. 7. 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 l.For many-electron atoms, energy depends on both n and l. See Active Figure 8.4, Figure 8.5, and Screen 8. 7. See Active Figure 8.4, Figure 8.5, and Screen 8. 7.

15. 15 © 2006 Brooks/Cole - Thomson 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 + l increases. b) for subshells of same n + l, subshell with lower n is lower in energy.

16. 16 © 2006 Brooks/Cole - Thomson Electron Filling Order Figure 8.5

17. 17 © 2006 Brooks/Cole - Thomson Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See Figure 8.6 and and Screen 8.6.Z* is the nuclear charge experienced by the outermost electrons. See Figure 8.6 and and Screen 8.6. 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* by --> [ Z - (no. inner electrons) ]Estimate Z* by --> [ 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!

18. 18 © 2006 Brooks/Cole - Thomson Effective Nuclear Charge Figure 8.6 Electron cloud for 1s electrons Z* is the nuclear charge experienced by the outermost electrons.

19. 19 © 2006 Brooks/Cole - Thomson 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.

20. 20 © 2006 Brooks/Cole - Thomson Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. One electron has n = 1, l = 0, m l = 0, m s = + 1/2 Other electron has n = 1, l = 0, m l = 0, m s = - 1/2

21. 21 © 2006 Brooks/Cole - Thomson See “Toolbox” on CD for Electron Configuration tool.

22. 22 © 2006 Brooks/Cole - Thomson Electron Configurations and the Periodic Table Active Figure 8.7

23. 23 © 2006 Brooks/Cole - Thomson LithiumLithium Group 1A Atomic number = 3 1s 2 2s 1 ---> 3 total electrons

24. 24 © 2006 Brooks/Cole - Thomson BerylliumBeryllium Group 2A Atomic number = 4 1s 2 2s 2 ---> 4 total electrons

25. 25 © 2006 Brooks/Cole - Thomson BoronBoron Group 3A Atomic number = 5 1s 2 2s 2 2p 1 ---> 5 total electrons 5 total electrons

26. 26 © 2006 Brooks/Cole - Thomson CarbonCarbon Group 4A Atomic number = 6 1s 2 2s 2 2p 2 ---> 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.

27. 27 © 2006 Brooks/Cole - Thomson NitrogenNitrogen Group 5A Atomic number = 7 1s 2 2s 2 2p 3 ---> 7 total electrons 7 total electrons

28. 28 © 2006 Brooks/Cole - Thomson OxygenOxygen Group 6A Atomic number = 8 1s 2 2s 2 2p 4 ---> 8 total electrons 8 total electrons

29. 29 © 2006 Brooks/Cole - Thomson FluorineFluorine Group 7A Atomic number = 9 1s 2 2s 2 2p 5 ---> 9 total electrons 9 total electrons

30. 30 © 2006 Brooks/Cole - Thomson NeonNeon Group 8A Atomic number = 10 1s 2 2s 2 2p 6 ---> 10 total electrons 10 total electrons Note that we have reached the end of the 2nd period, and the 2nd shell is full!

31. 31 © 2006 Brooks/Cole - Thomson Electron Configurations of p-Block Elements

32. 32 © 2006 Brooks/Cole - Thomson 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.

33. 33 © 2006 Brooks/Cole - Thomson 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.

34. 34 © 2006 Brooks/Cole - Thomson 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

35. 35 © 2006 Brooks/Cole - Thomson 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.

36. 36 © 2006 Brooks/Cole - Thomson Electron Configurations and the Periodic Table

37. 37 © 2006 Brooks/Cole - Thomson Transition Metals Table 8.4 All 4th period elements have the configuration [argon] ns x (n - 1)d y and so are d-block elements. CopperIronChromium

38. 38 © 2006 Brooks/Cole - Thomson Transition Element Configurations 3d orbitals used for Sc-Zn (Table 8.4)

39. 39 © 2006 Brooks/Cole - Thomson

40. 40 © 2006 Brooks/Cole - Thomson 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

41. 41 © 2006 Brooks/Cole - Thomson Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)

42. 42 © 2006 Brooks/Cole - Thomson

43. 43 © 2006 Brooks/Cole - Thomson 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- ---> P 3+ [Ne] 3s 2 3p 0

44. 44 © 2006 Brooks/Cole - Thomson Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s 2 3d 6 loses 2 electrons ---> Fe 2+ [Ar] 4s 0 3d 6 To form cations, always remove electrons of highest n value first!

45. 45 © 2006 Brooks/Cole - Thomson 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

46. 46 © 2006 Brooks/Cole - Thomson 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.

47. 47 © 2006 Brooks/Cole - Thomson PERIODIC TRENDS

48. 48 © 2006 Brooks/Cole - Thomson 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.

49. 49 © 2006 Brooks/Cole - Thomson Effective Nuclear Charge, Z* Z* is the nuclear charge experienced by the outermost electrons. See Figure 8.6 and and Screen 8.6.Z* is the nuclear charge experienced by the outermost electrons. See Figure 8.6 and and Screen 8.6. 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* by --> [ Z - (no. inner electrons) ]Estimate Z* by --> [ 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!

50. 50 © 2006 Brooks/Cole - Thomson Effective Nuclear Charge Figure 8.6 Electron cloud for 1s electrons Z* is the nuclear charge experienced by the outermost electrons.

51. 51 © 2006 Brooks/Cole - Thomson Effective Nuclear Charge Z* The 2s electron PENETRATES the region occupied by the 1s electron. 2s electron experiences a higher positive charge than expected.

52. 52 © 2006 Brooks/Cole - Thomson 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]

53. 53 © 2006 Brooks/Cole - Thomson Orbital Energies CD-ROM Screens 8.9 - 8.13, Simulations Orbital energies “drop” as Z* increases

54. 54 © 2006 Brooks/Cole - Thomson 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.

55. 55 © 2006 Brooks/Cole - Thomson Atomic Radii Active Figure 8.11

56. 56 © 2006 Brooks/Cole - Thomson Atomic Size Size goes UP on going down a group. See Figure 8.9.Size goes UP on going down a group. See Figure 8.9. 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 8.9.Size goes UP on going down a group. See Figure 8.9. 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.

57. 57 © 2006 Brooks/Cole - Thomson 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*

58. 58 © 2006 Brooks/Cole - Thomson Trends in Atomic Size See Active Figure 8.11

59. 59 © 2006 Brooks/Cole - Thomson Sizes of Transition Elements See Figure 8.12

60. 60 © 2006 Brooks/Cole - Thomson Sizes of Transition Elements See Figure 8.12 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!

61. 61 © 2006 Brooks/Cole - Thomson Density of Transition Metals 6th period 5th period 4th period

62. 62 © 2006 Brooks/Cole - Thomson 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?

63. 63 © 2006 Brooks/Cole - Thomson 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.

64. 64 © 2006 Brooks/Cole - Thomson Ion Sizes Does the size go up or down when gaining an electron to form an anion?

65. 65 © 2006 Brooks/Cole - Thomson 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 -

66. 66 © 2006 Brooks/Cole - Thomson Trends in Ion Sizes Active Figure 8.15

67. 67 © 2006 Brooks/Cole - Thomson 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?

68. 68 © 2006 Brooks/Cole - Thomson Ionization Energy See CD Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg + (g) + e-

69. 69 © 2006 Brooks/Cole - Thomson Mg (g) + 738 kJ ---> Mg + (g) + e- Mg + (g) + 1451 kJ ---> 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 See Screen 8.12

70. 70 © 2006 Brooks/Cole - Thomson Mg (g) + 735 kJ ---> Mg + (g) + e- Mg + (g) + 1451 kJ ---> Mg 2+ (g) + e- Mg 2+ (g) + 7733 kJ ---> 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 See Screen 8.12

71. 71 © 2006 Brooks/Cole - Thomson 2nd IE / 1st IE Li Na K B Al

72. 72 © 2006 Brooks/Cole - Thomson Trends in Ionization Energy Active Figure 8.13

73. 73 © 2006 Brooks/Cole - Thomson Trends in Ionization Energy

74. 74 © 2006 Brooks/Cole - Thomson Orbital Energies CD-ROM Screens 8.9 - 8.13, Simulations As Z* increases, orbital energies “drop” and IE increases.

75. 75 © 2006 Brooks/Cole - Thomson 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.

76. 76 © 2006 Brooks/Cole - Thomson 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

77. 77 © 2006 Brooks/Cole - Thomson Periodic Trend in the Reactivity of Alkali Metals with Water Lithium SodiumPotassium

78. 78 © 2006 Brooks/Cole - Thomson 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- ---> A - (g) E.A. = ∆E A(g) + e- ---> A - (g) E.A. = ∆E

79. 79 © 2006 Brooks/Cole - Thomson Electron Affinity of Oxygen ∆E is EXOthermic because O has an affinity for an e-. [He]      O atom EA = - 141 kJ + electron O [He]       - ion

80. 80 © 2006 Brooks/Cole - Thomson Electron Affinity of Nitrogen ∆E is zero for N - due to electron- electron repulsions. EA = 0 kJ [He]     Natom  [He]    N - ion  + electron

81. 81 © 2006 Brooks/Cole - Thomson Trends in Electron Affinity Active Figure 8.14

82. 82 © 2006 Brooks/Cole - Thomson See Figure 8.14 and Appendix FSee Figure 8.14 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