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Chapter 5. Covalent Compounds (Molecular Compounds)

Chapter 5. Covalent Compounds (Molecular Compounds). heat. NaCl (ionic compound). Na + + Cl - (gas). Heat. H 2 O (liquid). H 2 O molecules (gas). NH 3. NH 3 (molecules). etc. A molecular formula tells the # of atoms of each element in a molecule of the compound.

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Chapter 5. Covalent Compounds (Molecular Compounds)

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  1. Chapter 5. Covalent Compounds (Molecular Compounds) heat NaCl (ionic compound) Na+ + Cl- (gas) Heat H2O (liquid) H2O molecules (gas) NH3 NH3 (molecules) etc. A molecular formula tells the # of atoms of each element in a molecule of the compound C2F4 C2H6O

  2. A. Covalent bonds Example H + H H2 Sharing of electrons H H H H or H H

  3. H + Cl HCl H Lewis structure Octet rule Draw the Lewis structures of H2O

  4. NH3 CH4

  5. Consider O2 Consider N2 Draw Lewis structures for the following compounds CH2O C2H4

  6. B. Coordinate Covalent Bonds (less common) BH3 Electron deficient compound NH3 = Coordinate covalent bond

  7. Draw Lewis structure for each of the following molecular formulas in the most stable form (by pure sharing of electrons). a) PCl3 b) C2F6 c) CH2O2 d) CH3N e) C2H2Cl2 f) N2O2

  8. Common elements in covalent compounds: C, O, N

  9. For compounds containing C, H, O, N (the big 4), and F, try this

  10. HCN C3H4 CO2

  11. C. Compounds not following the Octet Rule NO PCl5

  12. E. Lewis structures of Polyatomic ions or molecules with a central atom. 1. Calculate the total number of valence electrons. 2 Draw a single bond between the central atom and each of the surrounding atoms. 3. Add nonbonding electrons to surrounding atoms such that each has an octet of electrons (2 on H). 4. Place the remaining electrons on the central atom. 5. If the central atom does not have octet of electrons, use one or two pairs of nonbonding e’s from the surround atoms to form double or triple bonds with the central atom. 6. Check the total number of electrons. NO2-

  13. - Resonance O N O NO2- ? - - ? Resonance structures or resonance contributors The real molecule or ion is a resonance hybrid of the resonance structures. Each resonance structure is less stable than the resonance hybrid.

  14. Neutral molecules with a central atom

  15. Polyatomic ions Examples: NO3- SO32-

  16. Lewis dot structures of ionic compounds: K2SO3 Ca(NO3)2

  17. F. Electronegativity (EN) Electronegativity of an element = the relative tendency of its atoms to attract the bonding electron pair. or HCl HCl EN of Cl > EN of H

  18. Fig.5.11 Pauling Electronegativity Values

  19. G. Polar covalent bond Figure 5.12: (a) (b) Nonpolar and Polar Covalent Bond d+ d- H Cl H Cl Polar covalent bond

  20. The relative E.N. determines the bond type Examples:

  21. Exercise Arrange the following bonds from most to least polar:  a) N-F O-F C-F b) C-F N-O Si-F c) Cl-Cl B-Cl S-Cl

  22. Molecular Geometry • Valence shell electron pair repulsion (VSEPR) theory CH4 All 4 bonds are equivalent Lewis structure Electron pair arrangement Tetrahedral 109.5o Molecular geometry C s p p p hybridize sp3 sp3 sp3 sp3 Four sp3 hybrid orbits

  23. o 180o 120o 109.5o hybrid orbitals: sp sp2 sp3

  24. Lewis structure NH3 sp3 Electron pair arrangement: tetrahedral Molecular geometry: trigonal pyramidal H2O sp3 Electron pair arrangement: tetrahedral Molecular geometry: angular

  25. BH3 sp2 Lewis structure Trigonal planar B s p p p s p p p sp2 sp2 sp2 Three sp2 hybrid orbitals Electron pair arrangement: trigonal planar Molecular geometry: trigonal planar

  26. SO2 sp2 Electron pair arrangement: trigonal planar Molecular geometry: bent BeH2 s p s p Two sp hybrid orbitals sp sp sp Electron pair arrangement: linear Molecular geometry: linear

  27. o 180o 109o 120o hybrid orbitals: sp sp2 sp3

  28. The Shape (Geometry) of Molecules

  29. Summary # of groups electron pair makeup molecular hybrid of electrons (density) of e- groups geometry orbitals around arrangement central atom 4 tetrahedral 4 bonding tetrahedral 3 bonding trigonal 1 nonbonding pyramidal sp3 2 bonding angular 2 nonbonding (bent) 3 Trigonal 3 bonding Trigonal planar planar sp2 2 bonding Angular 1 nonbonding 2 Linear 2 bonding Linear sp

  30. Examples: Electron pair arrangement Molecular geometry Hybird orbitals HCN

  31. I. Polarity of Molecules H Cl One polar bond Polar molecule net Polar molecule Non polar molecule The 2 polar bonds cancel each other

  32. Summary: • Draw the Lewis structure • If all electron groups around the central atom are • connected to the same atom – nonplar • otherwise - polar

  33. J. Naming of binary molecular compounds mono- 1 di- 2 tri- 3 tetra- 4 penta- 5 hexa- 6 hepta- 7 octa- 8 ennea-(neno) 9 deca- 10 P2O5 diphosphorus pentaoxide N2O4 dinitrogen tetraoxide CO2 SO2 NO

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