Valence Bond Theory PPTX


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Valence Bond Theory PPTX

  1. 1. Theories of Covalent Bonding<br />11.1 Valence Shell Electron Pair Repulsion Theory<br />11.2 Valence Bond (VB) Theory and Orbital Hybridization<br />11.3 Molecular Orbital (MO)Theory and Electron Delocalization<br />
  2. 2. Prentice Hall © 2003<br />Chapter 9<br />Covalent Bonding and Orbital Overlap<br /><ul><li>Lewis structures and VSEPR do not explain why a bond forms. How do we account for shape in terms of quantum mechanics?
  3. 3. What are the orbitals that are involved in bonding?
  4. 4. We use Valence Bond Theory:
  5. 5. Bonds form when orbitals on atoms overlap.
  6. 6. A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.
  7. 7. There are two electrons of opposite spin in the orbital overlap.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Covalent Bonding and Orbital Overlap<br />Hydrogen, H2<br />
  8. 8. Hydrogen fluoride, HF<br />Fluorine, F2<br />
  9. 9. Prentice Hall © 2003<br />Chapter 9<br />Covalent Bonding and Orbital Overlap<br /><ul><li>As two nuclei approach each other their atomic orbitals overlap.
  10. 10. As the amount of overlap increases, the energy of the interaction decreases.
  11. 11. At some distance the minimum energy is reached.
  12. 12. The minimum energy corresponds to the bonding distance (or bond length).
  13. 13. As the two atoms get closer, their nuclei begin to repel and the energy increases.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Covalent Bonding and Orbital Overlap<br /><ul><li>At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).</li></li></ul><li>
  14. 14. Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br /><ul><li>Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding.
  15. 15. Hybridization is determined by the electron domain geometry.</li></ul>sp Hybrid Orbitals<br /><ul><li>Consider the BeF2 molecule (experimentally known to exist):</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br />sp Hybrid Orbitals<br /><ul><li>Be has a 1s22s2 electron configuration.
  16. 16. There is no unpaired electron available for bonding.
  17. 17. We conclude that the atomic orbitals are not adequate to describe orbitals in molecules.
  18. 18. We know that the F-Be-F bond angle is 180 (VSEPR theory).
  19. 19. We also know that one electron from Be is shared with each one of the unpaired electrons from F.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br />sp Hybrid Orbitals<br /><ul><li>We assume that the Be orbitals in the Be-F bond are 180 apart.
  20. 20. We could promote and electron from the 2s orbital on Be to the 2p orbital to get two unpaired electrons for bonding.
  21. 21. BUT the geometry is still not explained.
  22. 22. We can solve the problem by allowing the 2s and one 2p orbital on Be to mix or form a hybrid orbital.
  23. 23. The hybrid orbital comes from an s and a p orbital and is called an sp hybrid orbital.</li></li></ul><li>
  24. 24. Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br />sp Hybrid Orbitals<br /><ul><li>The lobes of sp hybrid orbitals are 180º apart.
  25. 25. Since only one of the Be 2p orbitals has been used in hybridization, there are two unhybridized p orbitals remaining on Be.</li></li></ul><li>The sp hybrid orbitals in gaseous BeCl2.<br />Figure 11.2<br />atomic orbitals<br />hybrid orbitals<br />orbital box diagrams<br />
  26. 26. The sp hybrid orbitals in gaseous BeCl2(continued).<br />Figure 11.2<br />orbital box diagrams with orbital contours<br />
  27. 27.
  28. 28. The sp2 hybrid orbitals in BF3.<br />Figure 11.3<br />
  29. 29. sp2 and sp3<br />Hybrid Orbitals<br />
  30. 30. The sp3 hybrid orbitals in CH4.<br />Figure 11.4<br />
  31. 31. Figure 11.5<br />The sp3 hybrid orbitals in NH3.<br />
  32. 32. Figure 11.5 continued<br />The sp3 hybrid orbitals in H2O.<br />
  33. 33. Figure 11.6<br />The sp3d hybrid orbitals in PCl5.<br />
  34. 34. The sp3d2hybrid orbitals in SF6.<br />Figure 11.7<br />
  35. 35. Key Points<br />Types of Hybrid Orbitals<br />sp<br />sp2<br />sp3<br />sp3d<br />sp3d2<br />Hybrid Orbitals<br />The number of hybrid orbitals obtained equals the number of atomic orbitals mixed.<br />The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.<br />
  36. 36.
  37. 37. Step 1<br />Step 2<br />Step 3<br />Figure 11.8<br />The conceptual steps from molecular formula to the hybrid orbitals used in bonding.<br />Molecular shape and e- group arrangement<br />Molecular formula<br />Lewis structure<br />Hybrid orbitals<br />
  38. 38. PROBLEM:<br />Use partial orbital diagrams to describe mixing of the atomic orbitals of the central atom leads to hybrid orbitals in each of the following:<br />PLAN:<br />Use the Lewis structures to ascertain the arrangement of groups and shape of each molecule. Postulate the hybrid orbitals. Use partial orbital box diagrams to indicate the hybrid for the central atoms.<br />SAMPLE PROBLEM 11.1<br />Postulating Hybrid Orbitals in a Molecule<br />(a) Methanol, CH3OH<br />(b) Sulfur tetrafluoride, SF4<br />SOLUTION:<br />(a) CH3OH<br />The groups around C are arranged as a tetrahedron.<br />O also has a tetrahedral arrangement with 2 nonbonding e- pairs.<br />
  39. 39. hybridized C atom<br />hybridized O atom<br />single C atom<br />single O atom<br />hybridized S atom<br />S atom<br />SAMPLE PROBLEM 11.1<br />Postulating Hybrid Orbitals in a Molecule<br />continued<br />(b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs. <br />
  40. 40. Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br />Hybridization Involving d Orbitals<br /><ul><li>Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron domain geometries must involve d-orbitals.
  41. 41. Trigonal bipyramidal electron domain geometries require sp3d hybridization.
  42. 42. Octahedral electron domain geometries require sp3d2 hybridization.
  43. 43. Note the electron domain geometry from VSEPR theory determines the hybridization.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Hybrid Orbitals<br />Summary<br />Draw the Lewis structure.<br />Determine the electron domain geometry with VSEPR.<br />Specify the hybrid orbitals required for the electron pairs based on the electron domain geometry.<br />
  44. 44.
  45. 45.
  46. 46. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br /><ul><li>-Bonds: electron density lies on the axis between the nuclei.
  47. 47. All single bonds are -bonds.
  48. 48. -Bonds: electron density lies above and below the plane of the nuclei.
  49. 49. A double bond consists of one -bond and one -bond.
  50. 50. A triple bond has one -bond and two -bonds.
  51. 51. Often, the p-orbitals involved in -bonding come from unhybridized orbitals.</li></li></ul><li>Multiple Bonds<br />
  52. 52. both C are sp3 hybridized<br />s-sp3 overlaps to  bonds<br />sp3-sp3 overlap to form a  bond<br />relatively even distribution of electron density over all  bonds<br />The  bonds in ethane(C2H6).<br />Figure 11.9<br />
  53. 53. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />Ethylene, C2H4, has:<br /><ul><li>one - and one -bond;
  54. 54. both C atoms sp2 hybridized;
  55. 55. both C atoms with trigonal planar electron pair and molecular geometries.</li></li></ul><li>overlap in one position - <br />p overlap - <br />electron density<br />The  and  bonds in ethylene (C2H4).<br />Figure 11.10<br />
  56. 56. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />
  57. 57.
  58. 58. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br /><ul><li>When triple bonds form (e.g. N2) one -bond is always above and below and the other is in front and behind the plane of the nuclei.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />Consider acetylene, C2H2<br /><ul><li>the electron pair geometry of each C is linear;
  59. 59. therefore, the C atoms are sp hybridized;
  60. 60. the sp hybrid orbitals form the C-C and C-H -bonds;
  61. 61. there are two unhybridized p-orbitals;
  62. 62. both unhybridized p-orbitals form the two -bonds;
  63. 63. one -bond is above and below the plane of the nuclei;
  64. 64. one -bond is in front and behind the plane of the nuclei.</li></li></ul><li>overlap in one position - <br />p overlap - <br />The  and  bonds in acetylene (C2H2).<br />Figure 11.11<br />
  65. 65. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />
  66. 66. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />
  67. 67. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />Delocalized  Bonding<br /><ul><li>So far all the bonds we have encountered are localized between two nuclei.
  68. 68. In the case of benzene
  69. 69. there are 6 C-C  bonds, 6 C-H  bonds,
  70. 70. each C atom is sp2 hybridized,
  71. 71. and there are 6 unhybridized p orbitals on each C atom.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />Delocalized  Bonding<br />
  72. 72. Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />Delocalized  Bonding<br /><ul><li>In benzene there are two options for the 3  bonds
  73. 73. localized between C atoms or
  74. 74. delocalized over the entire ring (i.e. the  electrons are shared by all 6 C atoms).
  75. 75. Experimentally, all C-C bonds are the same length in benzene.
  76. 76. Therefore, all C-C bonds are of the same type (recall single bonds are longer than double bonds).</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Multiple Bonds<br />General Conclusions<br /><ul><li>Every two atoms share at least 2 electrons.
  77. 77. Two electrons between atoms on the same axis as the nuclei are  bonds.
  78. 78. -Bonds are always localized.
  79. 79. If two atoms share more than one pair of electrons, the second and third pair form -bonds.
  80. 80. When resonance structures are possible, delocalization is also possible.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Molecular Orbitals<br /><ul><li>Some aspects of bonding are not explained by Lewis structures, VSEPR theory and hybridization. (E.g. why does O2 interact with a magnetic field?; Why are some molecules colored?)‏
  81. 81. For these molecules, we use Molecular Orbital (MO) Theory.
  82. 82. Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.</li></li></ul><li>PROBLEM:<br />Describe the types of bonds and orbitals in acetone, (CH3)2CO.<br />PLAN:<br />Use the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of the multiple bonds and their orbital overlaps.<br />sp3 hybridized<br />sp3 hybridized<br />sp2 hybridized<br />SAMPLE PROBLEM 11.2<br />Describing the Bond in Molecules<br />SOLUTION:<br />bond<br />bonds<br />
  83. 83. The Central Themes of MO Theory<br />A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals.<br />Atomic wave functions are summed to obtain molecular wave functions.<br />If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei).<br />If wave functions cancel each other, an antibonding MO is formed (a node of zero electron density occurs between the nuclei).<br />
  84. 84. Amplitudes of wave functions subtracted.<br />Figure 11.14<br />An analogy between light waves and atomic wave functions.<br />Amplitudes of wave functions added<br />
  85. 85. Prentice Hall © 2003<br />Chapter 9<br />Molecular Orbitals<br /><ul><li>Molecular orbitals:
  86. 86. each contain a maximum of two electrons;
  87. 87. have definite energies;
  88. 88. can be visualized with contour diagrams;
  89. 89. are associated with an entire molecule.</li></ul>The Hydrogen Molecule<br /><ul><li>When two AOs overlap, two MOs form.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Molecular Orbitals<br />The Hydrogen Molecule<br /><ul><li>Therefore, 1s (H) + 1s (H) must result in two MOs for H2:
  90. 90. one has electron density between nuclei (bonding MO);
  91. 91. one has little electron density between nuclei (antibonding MO).
  92. 92. MOs resulting from s orbitals are  MOs.
  93. 93.  (bonding) MO is lower energy than * (antibonding) MO.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />Molecular Orbitals<br />The Hydrogen Molecule<br />
  94. 94. Prentice Hall © 2003<br />Chapter 9<br />Molecular Orbitals<br />The Hydrogen Molecule<br /><ul><li>Energy level diagram or MO diagram shows the energies and electrons in an orbital.
  95. 95. The total number of electrons in all atoms are placed in the MOs starting from lowest energy (1s) and ending when you run out of electrons.
  96. 96. Note that electrons in MOs have opposite spins.
  97. 97. H2 has two bonding electrons.
  98. 98. He2 has two bonding electrons and two antibonding electrons.</li></li></ul><li>Prentice Hall © 2003<br />Chapter 9<br />1s<br />1s<br />AO of H<br />AO of H<br />Figure 11.15<br />The MO diagram for H2.<br />Filling molecular orbitals with electrons follows the same concept as filling atomic orbitals.<br />*1s<br />Energy<br />H2 bond order = 1/2(2-0) = 1<br />1s<br />MO of H2<br />
  99. 99. Prentice Hall © 2003<br />Chapter 9<br />Second-Row Diatomic Molecules<br /> Electron Configurations and Molecular Properties<br /><ul><li>Two types of magnetic behavior:
  100. 100. paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule;
  101. 101. diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule.
  102. 102. Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:</li>