Chapter 5

Molecular Compounds

Chapter 5 suggested problems: 28, 36, 46, 52, 54, 56, 60, 66, 68, 67, 78, 80, 82, 84, 92, 94

  1. Covalent bonds and molecular compounds

    1. General

      1. Covalent bonds - between NM-NM, bonding electrons shared more or less equally

      2. Covalent compounds gain an octet of electrons through sharing electrons, i.e. through forming chemical bonds

      3. Covalent bonds result from the overlap of orbitals; atoms must be close enough to attract each other and share electrons but not so close that they repel each other

      4. Not all covalent bonds are equally strong or of equal length; these parameters depend on various factors

    2. Covalent bonds and the Periodic Table

      1. The number of bonds a nonmetal is likely to form can be predicted based on its position in the Periodic table

      2. There are a number of exceptions to the Octet Rule that will be discussed later this chapter

    3. Multiple covalent bonds

      1. It is possible for nonmetals to share more than one pair of electrons

      2. Nonmetals can form double and triple bonds

      3. These bonds are stronger and shorter than single bonds (units: kJ/mol, numbers from Atkins:67)

        1. C-C: 348 (1), 612 (2), 820 (3)

        2. Bond lengths for C-C bonds (pm): 154 (1), 133 (2), 120 (3)

        3. N-N: 163 (1), 409 (2), 945 (3)

        4. O-O: 146 (1), 497 (2)

        5. C-N: 305 (1), 613 (2), 890 (3)

      4. Lewis structures of molecules with multiple bonds

    4. Coordinate covalent bonds

      1. Bonds in which one atom donates both of the bonding electrons

      2. Even though one atom donates both bonding electrons, they are shared more or less equally by the bonding atoms

      3. Most common in ammonium ion and in transition metal complexes

      4. Lewis acid-base behavior

    5. Naming binary molecular compounds

      1. Molecular compounds: covalently bonded compounds

      2. See Ch. 4 notes

    6. Characteristics of molecular compounds

      1. Held together by covalent bonds

      2. Many molecular compounds dissolve but most of them do not dissociate since they are not ionic, i.e. solutions of molecular compounds do not conduct electricity

      3. The bonding forces between molecules are weaker than between ions, so as a rule molecular compounds have lower MP/BP than ionic compounds

  2. Molecular formulas and Lewis structures

    1. Why do we care about molecular geometry?

    2. Molecular formulas, structural formulas, condensed structural formulas, and skeleton (backbone) formulas

      1. As examples: 2-hexanol and 2,4-dimethyl-3-pentanone

      2. Remember constitutional isomers

    3. Lewis structures: show bonds between atoms and the presence of unpaired electrons

      1. To correctly estimate molecular geometry a knowledge of both bonding pairs and nonbonding pairs of electrons is necessary

    4. Drawing Lewis structures of atoms

      1. Review

    5. Drawing Lewis structures of molecules

      1. Rules

        1. Sum the valence electrons from all atoms

        2. Draw the skeleton structure: the central atom will usually be the first named atom(s)

        3. Subtract 2 e-/bond from the total electrons and divide the remainder by 2

        4. Assign nbp to terminal atoms

        5. Assign remaining nbp to central atom

        6. Form multiple bonds if necessary

        7. Resonance structures if appropriate

      2. Examples

        1. CH4

        2. PCl3

        3. HCN

        4. BrO3-

        5. NH4+

        6. SO3

        7. CO32-

    6. Exceptions to the octet rule

      1. Molecules with an odd number of electrons (ClO2, NO, NO2)

      2. Molecules with an atom with less than an octet (H, Be, B in BH3)

      3. Molecules with an atom with more than an octet

        1. Most common: P (10 electrons) and S (12 electrons)

        2. Other elements in the 3rd Period and higher take advantage of empty d orbitals to accommodate the extra electrons

          1. ICl4-

          2. PF5

  3. The shapes of molecules and VSPER - Valence Shell Paired-Electron Repulsion

    1. There is an excellent correlation between the type of hybridization that occurs in a molecule, the number of bonds formed, the number of nonbonding pairs of electrons on the central atom, and the geometry of the molecule

    2. The mutual repulsion of pairs of electrons cause them to arrange themselves as far apart spatially as possible

    3. Lone pairs take up slightly more space than bonding pairs and are more repulsive as a consequence

    4. Multiple bonds count as one bonding pair
      total pairs
      bonding pairs
      nonbonding pairs
      geometry
      bond angles
      hybridization
      example
      2
      2
      0
      linear
      180
      sp
      BeH2
       
      3
      3
      0
      trigonal planar
      120
      sp2
      BF3
       
      2
      1
      bent
       
      sp2
      NO2
       
      4
      4
      0
      tetrahedral
      109.5
      sp3
      CH4
       
      3
      1
      pyramidal
        
      sp3
      NH3
       
      2
      2
      bent
        
      sp3
      H2O
       
      5
      5
      0
      trigonal bipyramidal
      120, 90
      dsp3
      PCl5
       
      4
      1
      seesaw
       
      dsp3
      SF4
       
      3
      2
      t-shaped
       
      dsp3
      ClF3
       
      2
      3
      linear
       
      dsp3
      XeF2
       
      6
      6
      0
      octahedral
      90
      d2sp3
      SF6
       
      5
      1
      square pyramidal
       
      d2sp3
      BrF5
       
      4
      2
      square planar
       
      d2sp3
      XeF4

  4. Polar covalent bonds and polar molecules

    1. Electronegativity: a brief review

    2. Polar covalent bonds - bonding electrons are not shared equally due to differences in electronegativity between the bonding atoms

      1. Show Cl2 and HCl as examples

      2. If the bonding atoms are the same (e.g. H2, N2, O2, F2, etc.) the bonding atoms are shared equally (50%/50%)

      3. If the bonding atoms are not the same the bonding electrons spend more time around the more electronegative atom

      4. This causes partial charges to arise on the bonding nuclei

      5. If the difference in electronegativity is sufficiently large the bond is ionic (usually dEN greater than abt. 1.7)

      6. Molecules with polar covalent bonds are generally considered to be polar, which affects their interactions with other molecules

    3. Polar molecules

      1. As a general rule, a molecule with one or more polar bonds is itself polar

      2. This affects the strength of its interactions with neighboring molecules (Ch. 8)

      3. Molecular polarity is measured in Debye units (D) and range from 0.0 for nonpolar molecules to values of 2-5

      4. Exceptions: highly symmetric molecules with polar bonds may be nonpolar, depending on circumstances
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