Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver - PowerPoint PPT Presentation

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Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver PowerPoint Presentation
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Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver
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Chemistry and Chemical Reactivity 6th Edition John C. Kotz Paul M. Treichel Gabriela C. Weaver

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    2. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY

    3. Arrangement of Electrons in Atoms Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml)

    4. 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. Arrangement of Electrons in Atoms

    5. Electron Spin Quantum Number, ms

    6. Electron Spin and Magnetism

    7. Measuring Paramagnetism

    8. n ---> shell 1, 2, 3, 4, ... l ---> subshell 0, 1, 2, ... n - 1 ml ---> orbital -l ... 0 ... +l ms ---> electron spin +1/2 and -1/2

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

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

    12. Electrons in Atoms When n = 4, then l = 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 l. See Active Figure 8.4, Figure 8.5, and Screen 8. 7.

    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 ? increases. S<p<d<f b) for shells increase in energy: 1<2<3<4<

    16. Electron Filling Order Figure 8.5

    17. Effective Nuclear Charge, Z* 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) Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> [ 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. Effective Nuclear Charge

    19. 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.

    22. Electron Configurations and the Periodic Table

    23. The electron configuration for chlorine is: 1s2 2s22p6 3s5 1s2 2s22p6 3s23p5 1s2 2s22p5 1s2 2s22p6 3s23p6

    24. Lithium Group 1A Atomic number = 3 1s22s1 ---> 3 total electrons

    25. Beryllium Group 2A Atomic number = 4 1s22s2 ---> 4 total electrons

    26. Boron Group 3A Atomic number = 5 1s2 2s2 2p1 ---> 5 total electrons

    27. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 ---> 6 total electrons

    28. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3 ---> 7 total electrons

    29. Oxygen Group 6A Atomic number = 8 1s2 2s2 2p4 ---> 8 total electrons

    30. Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5 ---> 9 total electrons

    31. Neon Group 8A Atomic number = 10 1s2 2s2 2p6 ---> 10 total electrons

    32. What element has the electron configuration [Kr] 4d5 5s1? W Mo Ru Pm

    33. Electron Configurations of p-Block Elements

    34. 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.

    35. Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 [Ne] 3s2 3p1

    36. Phosphorus Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3 [Ne] 3s2 3p3

    37. 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.

    38. Electron Configurations and the Periodic Table

    39. What element has the electron configuration [Xe] 4f14 5d10 6s2 6p2? Hf Lu Pb Sn

    40. Transition Metals Table 8.4 All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are d-block elements.

    41. Transition Element Configurations

    43. What element has the electron configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2? Zn Ca Ge Ni

    44. The electron configuration for tin, Sn, element 50, is: [Ne] 4s2 3d10 4p2 [Ar] 4s2 3d10 4p2 [Kr] 5s2 4d10 5p2 [Xe] 5s2 4d10 5p2

    45. Lanthanides and Actinides All these elements have the configuration [core] nsx (n - 1)dy (n - 2)fz and so are f-block elements.

    46. Lanthanide Element Configurations

    48. 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- ---> P3+ [Ne] 3s2 3p0

    49. Ion Configurations For transition metals, remove ns electrons and then (n - 1) electrons. Fe [Ar] 4s2 3d6 loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6

    50. Ion Configurations How do we know the configurations of ions? Determine the magnetic properties of ions.

    51. 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.

    52. PERIODIC TRENDS

    53. General Periodic Trends Atomic and ionic size Ionization energy Electron affinity

    54. Effective Nuclear Charge, Z* 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) Z* increases across a period owing to incomplete shielding by inner electrons. Estimate Z* by --> [ 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!

    55. Effective Nuclear Charge

    56. Effective Nuclear Charge Z* The 2s electron PENETRATES the region occupied by the 1s electron. 2s electron experiences a higher positive charge than expected.

    57. Effective Nuclear Charge, Z* Atom Z* 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

    58. Orbital Energies

    59. General Periodic Trends Atomic and ionic size Ionization energy Electron affinity

    60. Atomic Radii

    61. Atomic Size Size goes UP on going down a group. See Figure 8.9. Because electrons are added further from the nucleus, there is less attraction. Size goes DOWN on going across a period.

    62. Atomic Size Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge.

    63. Trends in Atomic Size See Active Figure 8.11

    64. Sizes of Transition Elements See Figure 8.12

    65. Sizes of Transition Elements See Figure 8.12 3d subshell is inside the 4s subshell. 4s electrons feel a more or less constant Z*. Sizes stay about the same and chemistries are similar!

    66. Density of Transition Metals

    67. Ion Sizes

    68. Ion Sizes CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES.

    69. Ion Sizes Does the size go up or down when gaining an electron to form an anion?

    70. Ion Sizes ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Trends in ion sizes are the same as atom sizes.

    71. Trends in Ion Sizes

    72. Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest atomic radius? Na B Al C

    73. Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest (most negative) electron affinity? Na B Al C

    74. Compare the elements Na, B, Al, and C with regard to the following properties: Which has the largest (most positive) ionization energy? Na B Al C

    75. Redox Reactions Why do metals lose electrons in their reactions? Why does Mg form Mg2+ ions and not Mg3+? Why do nonmetals take on electrons?

    76. Ionization Energy See CD Screen 8.12 IE = energy required to remove an electron from an atom in the gas phase.

    77. Mg (g) + 738 kJ ---> Mg+ (g) + e- Ionization Energy See Screen 8.12

    78. Mg (g) + 735 kJ ---> Mg+ (g) + e- Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e- Ionization Energy See Screen 8.12

    79. 2nd IE / 1st IE B

    80. Trends in Ionization Energy

    81. Trends in Ionization Energy

    82. Orbital Energies

    83. Trends in Ionization Energy IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals. Metals are good reducing agents. Nonmetals lose electrons with difficulty.

    84. Trends in Ionization Energy IE decreases down a group Because size increases. Reducing ability generally increases down the periodic table. See reactions of Li, Na, K

    85. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? Mg < C < N < F N < Mg < C < F Mg < N < C < F F < C < Mg < N

    86. Which of the following elements would have the greatest difference between the first and the second ionization energy? lithium carbon fluorine nitrogen

    87. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? B < O < Al < F B < O < F < Al Al < B < O < F F < O < B < Al

    88. Which of the following groups of elements is arranged correctly in order of increasing affinity for electrons (that is, electron affinity becomes more negative)? Mg < S < Al < Cl Mg < Al < S < Cl Al < Mg < S < Cl Cl < S < Mg < Al

    89. Which of the following groups of elements is arranged correctly in order of decreasing atomic radius? Mg > S > Al > Cl Mg > Al > S > Cl Al > Mg > S > Cl Cl > S > Mg > Al

    90. Periodic Trend in the Reactivity of Alkali Metals with Water

    91. 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

    92. Electron Affinity of Oxygen ?E is EXOthermic because O has an affinity for an e-.

    93. Electron Affinity of Nitrogen ?E is zero for N- due to electron-electron repulsions.

    94. Trends in Electron Affinity

    95. See Figure 8.14 and Appendix F Affinity for electron increases across a period (EA becomes more positive). Affinity decreases down a group (EA becomes less positive). Trends in Electron Affinity

    96. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? Mg < C < N < F N < Mg < C < F Mg < N < C < F F < C < Mg < N

    97. Which of the following elements would have the greatest difference between the first and the second ionization energy? lithium carbon fluorine nitrogen

    98. Which of the following groups of elements is arranged correctly in order of increasing first ionization energy? B < O < Al < F B < O < F < Al Al < B < O < F F < O < B < Al

    99. Which of the following groups of elements is arranged correctly in order of increasing affinity for electrons (that is, electron affinity becomes more negative)? Mg < S < Al < Cl Mg < Al < S < Cl Al < Mg < S < Cl Cl < S < Mg < Al

    100. Which of the following groups of elements is arranged correctly in order of decreasing atomic radius? Mg > S > Al > Cl Mg > Al > S > Cl Al > Mg > S > Cl Cl > S > Mg > Al