The shapes of things
Molecular shape determines
properties
Bonding determines shape
Learning objectives
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Write Lewis dot structures for simple
molecules
Predict shape of simple molecules
Predict polarity of simple molecules
Covalent molecular compounds
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Covalent compounds are usually molecular
Bonds between atoms are covalent
Interactions between molecules are very weak
Atoms in a covalent molecule don’t stack like
marbles
Bonds have specific directions
Molecules have specific shapes
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Shape will depend on
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The number of atoms bonded to the central atom
The number of lone pairs around the central atom
Distinguish between
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Electronic geometry (molecular geometry)
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Consider atoms and lone pairs
Molecular shape
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Consider atoms only
Lewis dot structures: doing the dots
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Molecular structure in
simplest terms: arrange
valence electrons as dots in a
2-dimensional figure
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Only valence electrons are
shown
Electrons are either in:
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bonds
lone pairs (stable molecules do
not contain unpaired electrons –
with very few exceptions)
Octet rule is guiding principle:
each atom has 8 dots round it
(H has 2 dots)
Lewis dot structures made easy: the
S = N –A machine
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Start with the skeleton of the molecule
Least electronegative element is the central atom
 S = N - A
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N = total number of electrons required to fill octet for each
atom in the molecule (8 for each element, except 2 for H
and 6 for B)
A = total number of valence electrons
S = total number of electrons in bonds
We are given N and A; we need to find S
Applying the rules
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Calculate N for the molecule
Calculate A (all the dots)
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Determine S (no of dots in
bonds)
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NF3
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N = 8(N) + 3 x 8(F) = 32
A = 5(N) + 3 x 7(F) = 26
S = 32 – 26 = 6
include charges for ions (add
one for each –ve charge and
subtract one for each +ve charge)
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(S = N – A)
Satisfy all octets and create
number of bonds dictated by
S (may be multiple bonds)
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FNF
F
F NF
F
Two tests for dot structures
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Is the number of dots in the molecule equal to
the number of valence electrons?
Are all the octets satisfied?
If both yes structure is valid
If either no then back to the drawing board
Electronic geometry
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Identify central atom.
Many molecules
have more than one.
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Central atom has
more than one atom
bonded to it
Methanol has two central atoms
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O is one central atom –
bonded to H and C
C is another central
atom – bonded to O, H,
H and H
Consider geometry
around each one
separately
H
O
H
H
H
Counting regions of charge
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Count only atoms and lone pairs immediately
bonded to central atom
Count the regions of electrons
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Bonds – single, double or triple count as 1
Lone pairs count as 1
Number will be between 2 and 4 for molecules
that obey octet rule
Counting groups
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OF2 two bonds, two lone pairs
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Total groups = 4
CF4 four bonds, no lone pairs
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Total groups = 4
Double or triple bonds count as one
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CO2 has two groups
HCN has two groups
Total number of groups dictates
electronic geometry
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Octet rule:
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Two – linear
Three – trigonal planar
Four – tetrahedral
Additional possibilities (expand octet):
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Five – trigonal bipyramidal
Six - octahedral
Summary of possible molecular
shapes
Polar bonds and polar molecules
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Not all molecules
containing polar bonds
will themselves be polar.
Need to examine the
molecular shape
Ask the question:
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Do the individual bond
polarities cancel out?
If so, non polar. If not,
polar.
Consider some examples
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In CO2 (linear molecule) the two polar bonds
oppose each other exactly
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In chemical tug-o-war there is stalemate
The most important polar molecule
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In BF3 the three bonds cancel out – tug of war
stalemate
In H2O (bent) the polar bonds do not directly
oppose – no stalemate
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Lone pair also adds some component
Overall net polarity
Consequence of polarity: H2O is a liquid, CO2
is a gas
Symmetry and polarity
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If the molecule “looks”
symmetrical it will be
nonpolar
If the molecule “looks”
non-symmetrical it will
be polar
Rules of thumb for evaluation of
polarity
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Presence of one lone pair of electrons
Only one polar bond
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Always polar molecules
Two or more polar bonds
Do polar bonds perfectly oppose?
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If no, polar molecule