(L7) molecular geometry


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(L7) molecular geometry

  1. 1. Molecular Geometry and Hybridization of Atomic Orbitals
  2. 2. Molecular Geometry  Diatomic molecules are the easiest to visualize in three dimensions  HCl  Cl2  Diatomic molecules are linear
  3. 3. Valence Shell Electron Pair Repulsion Theory • The ideal geometry of a molecule is determined by the way the electron pairs orient themselves in space • The orientation of electron pairs arises from electron repulsions • The electron pairs spread out so as to minimize repulsion
  4. 4. 4 VSEPR Theory  Frequently, we will describe two geometries for each molecule. 1. Electronic geometry is determined by the locations of regions of high electron density around the central atom(s). 2. Molecular geometry determined by the arrangement of atoms around the central atom(s). Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used.
  5. 5. The Valence Shell Electron Pair Repulsion model predicts shapes. 1. e- pairs stay as far apart as possible to minimize repulsions. 2. The shape of a molecule is governed by the number of bonds and lone pairs present. 3. Treat a multiple bond like a single bond when determining a shape. Each is a single e-group. 4. Lone pairs occupy more volume than bonds. Predicting Molecular Shapes: VSEPR
  6. 6. Predicting Molecular Shapes 1. Draw Lewis structure 2. Determine the number of electron pairs around the central atom. Count a multiple bond as one pair. 3. Arrange electron pairs as shown in the next slide
  7. 7. Predicting Molecular Shapes: Linear Triangular planar Tetrahedral Triangular bipyramidal Octahedral
  8. 8. Basic shapes that minimize repulsions: If the molecule contains: • only bonding pairs – the angles shown are correct. • lone pair/bond mixtures – the angles change a little.  lone pair/lone pair repulsions are largest.  lone pair/bond pair are intermediate in strength.  bond/bond interactions are the smallest. linear triangular planar tetrahedral triangular bipyramidal octahedral
  9. 9. Examples  Illustrate the geometry of the following molecules: 1. BeH2 2. CH4 3. BF3 4. PCl5 5. SF6
  10. 10. Molecular Geometry
  11. 11. 12 VSEPR Theory 1 Lone pair to lone pair is the strongest repulsion. 2 Lone pair to bonding pair is intermediate repulsion. 3 Bonding pair to bonding pair is weakest repulsion.  Mnemonic for repulsion strengths lp/lp > lp/bp > bp/bp  Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o .
  12. 12. Bond Angles and Lone Pairs  Ammonia and water show smaller bond angles than predicted from the ideal geometry  The lone pair is larger in volume than a bond pair  There is a nucleus at only one end of the bond so the electrons are free to spread out over a larger area of space
  13. 13. The A-X-E Notation  A denotes a central atom  X denotes a terminal atom  E denotes a lone pair  Example  Water  H2O  O is central  Two lone pairs  Two hydrogen  AX2E2
  14. 14. Multiple Bonds
  15. 15. Molecular Geometry Summary with Lone Pairs Included
  16. 16. The steps in determining a molecular shape. Molecular formula Lewis structure Electron-group arrangement Bond angles Molecular shape (AXmEn) Count all e- groups around central atom (A) Note lone pairs and double bonds Count bonding and nonbonding e- groups separately. Step 1 Step 2 Step 3 Step 4
  17. 17. 20 Valence Bond (VB) Theory  Covalent bonds are formed by the overlap of atomic orbitals.  Atomic orbitals on the central atom can mix and exchange their character with other atoms in a molecule.  Process is called hybridization.  Hybrids are common: 1. Pink flowers 2. Mules  Hybrid Orbitals have the same shapes as predicted by VSEPR.
  18. 18. 21 Valence Bond (VB) Theory Regions of High Electron Density Electronic Geometry Hybridization 2 Linear sp 3 Trigonal planar sp2 4 Tetrahedral sp3 5 Trigonal bipyramidal sp3d 6 Octahedral sp3d2
  19. 19. Valence Bond Theory  Unpaired electrons from one atom pair with unpaired electrons from another atom and give rise to chemical bonds  Simple extension of orbital diagrams
  20. 20. Multiple Bonds
  21. 21. Hybrid Orbitals  Hybridization of the s and p orbitals on carbon.  The four sp3 hybrid orbitals have equal energy.  The four valence electrons are distributed evenly across the sp3 hybrid orbitals.  The angle between the sp3 hybrid orbitals is 109.5o.
  22. 22. Hybrid Orbitals The number of hybrid orbitals obtained equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed. Key Points sp sp2 sp3 sp3d sp3d2 Types of Hybrid Orbitals
  23. 23. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2. atomic orbitals hybrid orbitals orbital box diagrams
  24. 24. Figure 11.3 The sp2 hybrid orbitals in BF3.
  25. 25. Figure 11.4 The sp3 hybrid orbitals in CH4.
  26. 26. Figure 11.5 The sp3 hybrid orbitals in NH3.
  27. 27. Figure 11.5 continued The sp3 hybrid orbitals in H2O.
  28. 28. Figure 11.6 The sp3d hybrid orbitals in PCl5.
  29. 29. Figure 11.7 The sp3d2 hybrid orbitals in SF6.
  30. 30. Hybrid Orbitals and Geometry