1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f

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1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f
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d block p block s block Periodic Table: s , p , d , and f Orbitals f block
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Slide 15 Orbital Energy Diagrams and Electron Configurations Preliminary Considerations Pauli Exclusion Principle We discussed this in Chapter 7: No two electrons in an atom can have the same 4 quantum numbers. As discussed before, the consequence is that electrons pair up with two in each orbital. As an example, consider a specific 3p orbital (corresponding to m l = 0). The two sets of quantum numbers allowed are: n = 3, l = 1, m l = 0, m s = +½ n = 3, l = 1, m l = 0, m s = -½ We cannot put a third electron into this orbital without duplicating one of the above sets of quantum numbers.
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Slide 16 Partial Orbital Energy Diagrams and Electron Configurations For Period 3 Elements Condensed Electron Configurations
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Slide 17 Condensed Electron Configurations The valence (outer shell) electrons are the ones that are most important for chemical bonding. To emphasize the outer shell configuration, one often writes the condensed electron configuration. The core electrons are represented by the noble gas [in square brackets] of lower atomic number closest to the atom of interest. Li (Z=3): 1s 2 2s 1 or [He]2s 1 Na (Z=11): 1s 2 2s 2 2p 6 3s 1 or [Ne]3s 1 Cl (Z=17): 1s 2 2s 2 2p 6 3s 2 3p 5 or [Ne]3s 2 3p 5
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Slide 18 Electron Configurations of Main Group Elements Sr (Z=38) Te (Z=52) Bi (Z=83) [Kr]5s 2 4d 10 5p 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 [Kr]5s 2 [Xe]6s 2 4f 14 5d 10 6p 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 3 10 e 18 e 36 e 10 e 18 e 36 e 10 e 18 e 36 e 54 e
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Slide 19 Electron Configurations and the Periodic Table Li (Z=3) 1s 2 2s 1 [He] 2s 1 Core electrons Valence electrons F (Z=9) 1s 2 2s 2 2p 5 [He] 2s 2 2p 5 Cl (Z=17) 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne] 3s 2 3p 5 Br (Z=35) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 [Ar] 4s 2 3d 10 4p 5
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Slide 20 Electron Configurations and the Periodic Table Gp 1A: [core] ns 1 Gp 2A: [core] ns 2 Gp 3A: [core] ns 2 np 1 Gp 4A: [core] ns 2 np 2 Gp 5A: [core] ns 2 np 3 Gp 6A: [core] ns 2 np 4 Gp 7A: [core] ns 2 np 5 Gp 8A: [core] ns 2 np 6
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Slide 21 Co (Z=27) Pt (Z=78) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 [Ar]4s 2 3d 7 [Xe]6s 2 4f 14 5d 8 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 8 Electron Configurations of Transition Metal Elements 10 e 18 e 10 e 18 e 36 e 54 e Two Anomalous Transition Metals Cr (Z=24): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu (Z=29): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10
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Slide 22 Electron Configurations of Transition Metal Elements Two Anomalous Transition Metals Cr (Z=24 ): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu (Z=29): 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 The standard explanation for these anomalous transition metals is the increased stability of half-filled and completely-filled d subshells.
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