Lewis Structures & Molecular Geometry by F45J2n6

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Lewis Structures & Molecular
Geometry
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
together
 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
VSEPR
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
Linear

 BeCl2
 CO2
Trigonal Planar

BF3
Bent

   SF2
Bent

   H2O
Tetrahedral

   CH4
Trigonal Pyramidal

   NH3
Quick Review: Rules for Drawing Lewis
Structures
   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
molecules.
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
molecules
 Hydrogen Bonding – between molecules that
contain H bonded to O, N or F
 London Dispersion Forces – between large
nonpolar molecules
MC

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

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

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.
IMF

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?
Polarity
   Predict which molecule has the more
polar bonds:
   PCl3 or AsCl3
   SnO or SrO
   SF2 or GeF4
   SiCl4 or SCl2
   Br2 or HBr
Review

   Why is K2S ionic but H2S molecular?

   Write the electron configuration for the Zn2+
ion.
Bonds
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.
Bonds

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

Identify the polar molecules
H – C ≡ N:

S=C=S
   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
MC

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

 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
MC

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.
MC

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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