Laboratory 06
MOLECULAR GEOMETRY AND POLARITY
The three dimensional (3D) structure of a molecule
-How its atoms are connected and arranged in space
Molecular geometry Polar
Boiling point, Freezing point, Chemical
reactivity
Non Polar
Reactivity
Toxicity
Determination of the correct 3D structure of a
molecule
Strategy
Lewis structure
Electron
domain
geometry
Molecular
geometry
Background- Lewis structure
Diagrams that show the bonding between atoms of a
molecule and the lone pairs of electrons that may
exist in the molecule
Construction of Lewis Structures
Two Rules
1. Total # of valence electrons – the total number
of valence electrons must be accounted for, no
extras, none missing.
2. Octet Rule – every atom should have an octet
(8) electrons associated with it.
Hydrogen should only have 2 (a duet).
Drawing Lewis Dot Structures
1.
2.
3.
4.
5.
6.
7.
Determine the total number of valence electrons.
Determine which atom is the “central” atom.
Stick everything to the central atom using a single
bond.
Fill the octet of every atom by adding dots.
Verify the total number of valence electrons in the
structure.
Add or subtract electrons to the structure by
making/breaking bonds to get the correct # of valence
electrons.
Check the “formal charge” of each atom.
Formal Charge of an Atom
Formal charge = number of valence electrons – number of
bonds – number of non-bonding electrons.
Determination of the electron domain geometry
around the central atom
A) Lewis structures do not indicate the three dimensional
shape of a molecule. They do not show the arrangement
space of the atoms, what we call the molecular geometry or
molecular structure.
B) Molecules have definite shapes and the shape of a
molecule controls some of its chemical and physical
properties.
Valence Shell Electron Pair Repulsion Theory - VSEPR
- predicts the shapes of a number of molecules and
polyatomic ions.
VSEPR Theory
• Predicts the molecular shape of a bonded molecule
• Electrons around the central atom arrange themselves as
far apart from each other as possible
• Unshared pairs of electrons (lone pairs) on the central
atom repel the most
lone pairs of electrons are spread out more broadly than
bond pairs.
(repulsions are greatest between two lone pairs,
intermediate between a lone pair and a bond pair, and
weakest between two bonding pairs of electrons)
• Repulsive forces decrease rapidly with increasing
interpair angle
(greatest at 90o, much weaker at 120o, and very weak at
180o)
Electron domain geometry
-Describe the geometric arrangement around the atom
-Compose, Single lone pair of electrons or chemical bond; a
single, double, or triple bond
-There are five basic electron domain geometries possible.
Molecules with no lone pairs
Trigonal planar
Trigonal bipyramidal
Trigonal planar
Trigonal bipyramidal
Molecules with 3 electron groups
Trigonal planar
Trigonal planar
Trigonal planar
Bent
Molecules with 4 electron groups
Trigonal pyramidal
Bent
Molecules with 5 electron groups
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Trigonal
Bipyramidal
Molecules with 6 electron groups
Bond Polarity
•
•
•
•
Due to differences in electronegativities of the bonding atoms
If Den = 0, bond is nonpolar covalent
If 0 < Den < 2, bond is polar covalent
If Den > 2, bond is ionic
m
Polarity
• Covalent bonds and molecules are either
polar or nonpolar
• Polar:
– Electrons unequally shared
– More attracted to one nuclei
• Nonpolar:
– Electrons equally shared
• Measure of polarity: dipole moment (μ)
Molecular Polarity
•
•
•
•
•
Overall electron distribution within a molecule
Depends on bond polarity and molecular geometry
Vector sum of the bond dipole moments
Lone pairs of electrons contribute to the dipole moment
Consider both magnitude and direction of individual bond
dipole moments
• Symmetrical molecules with polar bonds = nonpolar
Resonance structures
In chemistry, resonance is a way of describing delocalized
electrons within certain molecules or polyatomic ions, where
the bonding cannot be expressed by one single Lewis
formula.
A molecule or ion with such delocalized electrons is
represented by several contributing structures also called
resonance structures.
Resonance structures
1. Do not violate the octet rule!!! (DO NOT HAVE 5 BONDS TO CARBON!!!)
2.The overall charge of a molecule should not change—atoms may have
charges, but the net charge of the entire molecule should not change
3. Place resonance structures inside brackets ([ ]) and use
to separate
each structure
4. Do not break σ bonds! (e.g. do not break C-C, C-H, C-O, or C-N single
bonds)
5.Charges will be preferentially located on atoms of compatible
electronegativity. For example, oxygen is more electronegative than carbon;
therefore, a negative charge will preferentially be placed on oxygen rather
than carbon in the dominant resonance structure.
6. Unless starting with a radical, move electrons in pairs, using a doubleheaded arrow.
7. Do not “jump” lone pairs from one atom to another. Lone pairs can
become π bonds or π bonds can become lone pairs. π Bonds can migrate
from one side of a carbon atom to
another.