Lewis Structures & Molecular Geometry by F45J2n6

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									Lewis Structures & Molecular
    Honors Theoretical Chemistry
             Chapter 6

 Link to Shapes of Molecules Website
Electron Distribution
8  electrons = full outer shell = stable
  (except H and He which only hold a duet
  of 2)
 Shared pair – a pair of electrons shared
  by 2 atoms and thus bonding them
 Lone pair – nonbonding electrons
    Rules for Writing Lewis Structures

    1. Count the valence electrons. Put them in your bucket.
   2. ID the central atom. Look for an atom that can form lots of
    bonds -- NOT hydrogen.
   3. Bond each of the other atoms to the central atom with
    single bonds.
   4. Complete valence shells of outside atoms.
   5. Use remaining electrons to complete valence shell of
    central atom.
   6. If you run out of electrons, form double or triple bonds.
   7. If you have extra electrons, put them on the central atom.
Draw Arsenic Triiodide

 Arsenic Triiodide has the formula AsI3
 How many total valence electrons?
  Let’s draw the Lewis structure
 How many shared pairs of electrons does
  As have? How many unshared pairs does
  As have?
Draw these Lewis Structures

   Cl2
   O2
   N2
   CH3I
   C2H4
   CH2O
valence shell electron pair repulsion

 Repulsions between valence electrons
  causes them to be as far apart as possible in
  three dimensional space.
 These spatial relationships can be used to
  determine the shape of a given molecule.
    5 Shapes To Know (and love?)

 See Table 5, p.200 in Modern Chemistry
 Linear
 Trigonal Planar
 Bent
 Tetrahedral
 Trigonal Pyramidal

 BeCl2
 CO2
Trigonal Planar


   SF2

   H2O

   CH4
Trigonal Pyramidal

   NH3
Quick Review: Rules for Drawing Lewis
   Count all the valence electrons
   Place the most electronegative atom in the center
    and add the other atoms around it
   Draw bonds between the atoms
   Add pairs of electrons around each atom until the
    atom has 8 electrons; do not exceed the total
    number of valence electrons
Identify the geometry of these molecules:

   PCl3
   ClO3-
   CO2
   SF2
   CCl2F2
Polar Molecules (Dipoles)

 In a bond, the atom with the higher EN will
  attract the electrons more strongly
 The will cause that end of the molecule to
  have a partial negative charge (-), the other
  end will have a partial positive charge (+)
 Lewis structures which have a central atom
  with lone pairs of electrons indicate polar
Intramolecular forces

 Bonds  within a molecule that hold the
  atoms together
 Covalent Bonds
 Ionic Bonds
Intermolecular Forces

   Forces between molecules that hold them near
    each other
      Dipole-Dipole Forces – between polar
      Hydrogen Bonding – between molecules that
       contain H bonded to O, N or F
      London Dispersion Forces – between large
       nonpolar molecules

Which molecule contains a double
a. COCl2
b. C2H6
c. CF4
d. SF2

Which molecule is polar?
a. CCl4
b. CO2
c. SO3
d. None of these

What is the molecular geometry of CH3+?
a. tetrahedral
b. pyramidal
c. bent
d. trigonal planar
Lewis Structures

The Lewis structure for HCN contains one
 double bond and one single bond. Draw it.

Naphthalene, C10H8, is a nonpolar
molecule with a boiling point of 208 oC.
Acetic acid, CH3COOH, is a polar
molecule with a boiling point of 118 oC.
Which has stronger intermolecular forces?
   Predict which molecule has the more
    polar bonds:
   PCl3 or AsCl3
   SnO or SrO
   SF2 or GeF4
   SiCl4 or SCl2
   Br2 or HBr

   Why is K2S ionic but H2S molecular?

   Write the electron configuration for the Zn2+
     Certain forces hold the water molecules in ice
    together. However, at 0 °C ice melts. When water
    reaches 100 °C, boiling occurs and the water
    molecules finally break free of each other. Even in
    the vapor state forces persist between the oxygen
    and hydrogen atoms. At several thousand degrees
    Celsius, the hydrogen atoms break free and water
    molecules no longer exist. Finally, at tens of
    thousands of degrees Celsius, a new and highly
    charged state of matter emerges.
    Retell the above story using specific names for the
    forces and bonds that are involved at all levels and
    include the name for the last state of matter.

 When ice melts to form water, do chemical
  bonds break? ____________
 When NaCl dissolves in water, do chemical
  bonds break? ____________

Identify the polar molecules
 H – C ≡ N:

   Why is H2O a dipole but CO2 is not, yet
    they both have 2 polar covalent bonds?

   Which intermolecular forces are at work
    between the noble gas atoms?
  Would these most likely be formed by
  polar or nonpolar molecules?
 a) solid at room temperature
 b) gas at room temperature
 c) liquid with high boiling point
 d) liquid with low boiling point

 Molecular compounds are easy to melt
 a. covalent bonds are generally weak.
 b. forces between molecules are generally
 c. they are soluble in water
  d. none of the above

 Low electronegativity is characteristic of
   a. metals        b. Group V
    c. metalloids d. nonmetals
    When an atom loses one or more electrons a(n)
    ______ is produced.
 In which compound are the bonds most polar?
    a. SbBr3 b. SbCl3 c. SbF3 d. SbI3

  How does the strength of an intramolecular
  bond compare to the strength of
  intermolecular attractions?
     a. Intramolecular bonds are weaker.
     b. Intramolecular bonds are stronger.
     c. They are about the same.
     d. No generalization is possible.

In order for arsenic (As) to form a stable ion
it must
a. lose 5 electrons.
b. lose 3 electrons.
c. gain 5 electrons.
d. gain 3 electrons.

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