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Chapter 7 Atomic Electron Configurations and Chemical Periodicity

Chapter 7 Atomic Electron Configurations and Chemical Periodicity. Important – Read Before Using Slides in Class

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Chapter 7 Atomic Electron Configurations and Chemical Periodicity

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  1. Chapter 7Atomic Electron Configurations and Chemical Periodicity

  2. 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.

  3. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

  4. Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (s) ORBITALS (ms)

  5. Arrangement of Electrons in Atoms 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, ms.

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

  7. Electron Spin and Magnetism • Diamagnetic: NOT attracted to a magnetic field • Paramagnetic: substance is attracted to a magnetic field. • Substances with unpaired electrons are paramagnetic. PLAY MOVIE

  8. 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. QUANTUM NUMBERS Now there are four! n fshell 1, 2, 3, 4, ... sf subshell 0, 1, 2, ... n - 1 msf orbital - s ... 0 ... + s msfelectron spin +1/2 and -1/2

  10. 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. Electrons in Atoms When n = 1, then s = 0 this shell has a single orbital (1s) to which 2e- can be assigned. When n = 2, then s = 0, 1 2s orbital 2e- three 2p orbitals 6e- TOTAL = 8e-

  12. Electrons in Atoms When n = 3, then s = 0, 1, 2 3s orbital 2e- three 3p orbitals 6e- five 3d orbitals 10e- TOTAL = 18e-

  13. And many more! Electrons in Atoms When n = 4, then s = 0, 1, 2, 3 4s orbital 2e- three 4p orbitals 6e- five 4d orbitals 10e- seven 4f orbitals 14e- TOTAL = 32e-

  14. Assigning Electrons to Atoms • Electrons generally assigned to orbitals of successively higher energy. • For H atoms, E = - C(1/n2). E depends only on n. • For many-electron atoms, energy depends on both n and s. • See Active Figure 7.1 and Figure 7.2

  15. Assigning Electrons to Subshells • In H atom all subshells of same n have same energy. • 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

  16. Electron Filling OrderSee Figure 7.2

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

  18. Electron cloud for 1s electrons Effective Nuclear Charge See Figure 7.3 Z* is the nuclear charge experienced by the outermost electrons.

  19. spdf notation for H, atomic number = 1 1 no. of s 1 electrons value of l value of n Writing Atomic Electron Configurations Two ways of writing configs. One is called the spdf notation.

  20. Writing Atomic Electron Configurations Two ways of writing configs. Other is called the orbital box notation. One electron has n = 1, s = 0, ms = 0, ms = + 1/2 Other electron has n = 1, s = 0, ms = 0, ms = - 1/2

  21. See “Toolbox” in ChemNow for Electron Configuration tool.

  22. Electron Configurations and the Periodic Table See Active Figure 7.4

  23. Lithium Group 1A Atomic number = 3 1s22s1f 3 total electrons

  24. Beryllium Group 2A Atomic number = 4 1s22s2f 4 total electrons

  25. Boron Group 3A Atomic number = 5 1s2 2s2 2p1f 5 total electrons

  26. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2f 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. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3f 7 total electrons

  28. Oxygen Group 6A Atomic number = 8 1s2 2s2 2p4f 8 total electrons

  29. Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5f 9 total electrons

  30. Neon Group 8A Atomic number = 10 1s2 2s2 2p6f 10 total electrons Note that we have reached the end of the 2nd period, and the 2nd shell is full!

  31. Electron Configurations of p-Block Elements PLAY MOVIE

  32. Sodium Group 1A Atomic number = 11 1s2 2s2 2p6 3s1 or “neon core” + 3s1 [Ne] 3s1 (uses rare gas notation) Note that we have begun a new period. All Group 1A elements have [core]ns1 configurations.

  33. Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 [Ne] 3s2 3p1 All Group 3A elements have [core] ns2 np1 configurations where n is the period number.

  34. Yellow P Red P Phosphorus Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3 All Group 5A elements have [core ] ns2 np3 configurations where n is the period number.

  35. Calcium Group 2A Atomic number = 20 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2 All Group 2A elements have [core]ns2 configurations where n is the period number.

  36. Electron Configurations and the Periodic Table

  37. Transition MetalsTable 7.4 All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are d-block elements. Chromium Iron Copper

  38. Transition Element Configurations 3d orbitals used for Sc-Zn (Table 7.4)

  39. Cerium [Xe] 6s2 5d1 4f1 Uranium [Rn] 7s2 6d1 5f3 Lanthanides and Actinides All these elements have the configuration [core] nsx (n - 1)dy (n - 2)fz and so are f-block elements.

  40. Lanthanide Element Configurations 4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 7.2)

  41. Ion Configurations To form cations from elements remove 1 or more e- from subshell of highest n [or highest (n + l)]. P [Ne] 3s2 3p3 - 3e- f P3+ [Ne] 3s2 3p0

  42. Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s2 3d6 loses 2 electrons f Fe2+ [Ar] 4s0 3d6 To form cations, always remove electrons of highest n value first!

  43. Sample of Fe2O3 with strong magnet Sample of Fe2O3 Ion Configurations How do we know the configurations of ions? Determine the magnetic properties of ions.

  44. Fe3+ ions in Fe2O3 have 5 unpaired electrons and make the sample paramagnetic. 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.

  45. PLAY MOVIE PERIODIC TRENDS PLAY MOVIE PLAY MOVIE

  46. Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. General Periodic Trends • Atomic and ionic size • Ionization energy • Electron affinity

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

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