Lecture 41
Molecular Structures
Ozgur Unal
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Molecular formula for compounds do not show how atoms
are bonded together in a molecule.
In order to show the structure of a molecule we can use
different representations:
Space-filling molecular model
Ball-and-stick molecular model
Structural formula
Lewis structure
Check out Figure 8.13
Structural formula and Lewis structure are very useful to
show how atoms are bonded.
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Steps to draw the Lewis structure of molecules:
1- Predict the location of certain atoms.
2- Determine the total number of electrons available for
bonding.
3- Determine the number of electron pairs available for
bonding.
4- Place the bonding pairs.
5- Determine the number of remaining electron pairs.
6- Determine whether each atom satisfies the octet rule.
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Example: Ammonia is a raw material used in the
manufacture of many materials, including fertilizers, cleaning
products, and explosives. Draw the Lewis structure for
ammonia.
Example: A nitrogen trifluoride molecule contains
numerous lone pairs. Draw its Lewis structure.
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Example: Carbon dioxide is a product of all cellular
respiration. Draw the Lewis structure of carbon dioxide.
Example: Draw the Lewis structure of ethylene, C2H4.
Example: A molecule of carbon disulfide contains both lone
pairs and multi-covalent bonds. Draw its Lewis structure.
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Although the unit acts as an ion, the atoms with polyatomic
ion are covalently bonded.
Calculate the total number of valence electrons
Subtract the charge of the ion in order to find the total
number of electrons available for bonding.
Example: PO4-3
5 valence electrons from P
4*6 = 24 valence electrons from O4
3 electrons from the charge
 Total 32 electrons available for bonding.
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Example: Draw the Lewis structure for NH4+.
Example: Draw the Lewis structure for ClO4-.
Example: Draw the Lewis structure for HCO3-.
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Lecture 42
Resonance Structures
Ozgur Unal
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Using the same sequence of atoms, it is possible to have
more than one correct Lewis structure when a molecule or
polyatomic ion has both double bond and a single bond.
Example: NO3-  Check out Figure 8.14
Resosnance is a condition that occurs when more than one
valid Lewis structure can be written for a molecule or ion.
Example: NO2-, SO3-2, CO3-2, O3
Each molecule or ion that undergoes resonance behaves as if
it has only one structure.
Experimentally measured bond lengths show that the bonds
are identical to each other.
They are shorter than single bonds but longer than double
bonds.
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Example: Draw the Lewis resonance structure for NO2-.
Example: Draw the Lewis resonance structure for SO2.
Example: Draw the Lewis resonance structure for O3.
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Some molecules do not obey the octet rule.
There are several reasons for these exceptions:
Odd number of valence electrons
Suboctet and coordinate covalent bonds
Expanded octets
Odd number of valence electrons:
Example: NO2, ClO2 and NO
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Suboctet and coordinate covalent bonds:
Suboctet  Stable configurations with fewer than 8
electrons present around an atom.
Example: BF3
A coordinate covalent bond forms when one atom donates
both of the electrons to be shared with an atom or ion that
needs two electrons.
Example: BF3 + NH3
Atoms or ions with lone pairs often form coordinate
covalent bonds with atoms or ions that need two more
electrons.
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Expanded Octets:
These are molecules with central atoms that contain more
than 8 valence electrons.
This electron arrangement is referred to as an expanded
octet.
An expanded octet can be explained by considering the d
orbital that occurs in the energy levels of elements in period
three or higher.
Example: PCl5, SF6
Example: Draw the Lewis structure for Xenon tetrafluoride.
Example: Draw the Lewis structure for ClF3.
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