Week 2 F11 W - How
to
determine
Lewis
structures
...

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Unformatted text preview: How
to
determine
Lewis
structures
 •  Recommended
procedure:
 –  Determine
which
atom
in
the
central
atom
(single
atom,
least
 electronega:f…)
 –  Count
the
total
number
of
valence
electrons
(N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each nega7ve charge; subtract one electron for each posi7ve charge). –  Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required. –  Using lone pairs, complete octets around the noncentral atoms. –  Count the number of electrons depicted (two for each bond and two for each lone pair). –  If this number is less than N, then add electrons to the central atom un7l the total number of electrons depicted is N. –  If the octet rule is not sa7sfied for the central atom and lone‐pair electrons are nearby, use those electrons to make double or triple bonds to the central atom. –  Check each atom to see if it has a formal charge. (Singly bonded oxygen will require a nega<ve charge, for example.)
 Examples
 •  Determine
the
Lewis
structures
for
the
following
molecules
 –  H2O
 –  NH3
 –  NH4+
 –  CH4
 –  CO32‐
 –  F2
 –  O2
 –  BH3
 –  PCl5
 hypervalent
 –  H2SO4
 –  XeO2
 –  XeF2
 –  XeF4
 –  SF6
 hypervalent
 Lewis
structures
 Resonance
Structures
 Molecular
Shape
and
the
VSEPR
 •  VSEPR
(Valence
Shell
Electron
Pair
Repulsion)
 •  TO
PREDCIT
THE
SHAPE
OF
MOLECULES!!!!
 •  Based
on
electron
electron
repulsion.
 Molecular
Shape
 2‐coordinate
 E
:
central
atom
 X
:
coordinated
atom
 L
:
lone
pair
 VSEPR
(Valence
Shell
Electron
Pair
Repulsion)
 CO2
 EX2
;EL3X2
;EL4X2
 Θ
=
180°
 H2O
 ELX2
 Θ
<
120°
 EL2X2
 Θ
<
109.5°
 3‐coordinate
 E
:
central
atom
 X
:
coordinated
atom
 L
:
lone
pair
 BH3
 EX3
 Θ
=
120°
 [XeF3]+
 EL3X3
;EL2X3
 Θ
=
90°
 NH3
 ELX3
 Θ
<
109.5°
 4‐coordinate
 E
:
central
atom
 X
:
coordinated
atom
 L
:
lone
pair
 Seesaw
 EX4
 Θ
=
109°
 ELX4
 Θ
=
90°
 Θ
<
120°
 EL2X4
 Θ
=
90°
 5‐coordinate
 E
:
central
atom
 X
:
coordinated
atom
 L
:
lone
pair
 EX5
 Θ
=
90°
 Θ
=
120°
 ELX5
 EL2X5
 Θ
<
90°
 Θ
~
72°
 Trigonal
bipyramidal
geometry
 •  There
are
two
possible
loca:ons
of
lone
pairs
:
axial
 and
equatorial.
 –  If
there
is
a
single
lone
pair,
the
lone
pair
occupies
an
 equatorial
posi:on
(lone
pair
has
the
most
space
and
 minimize
the
interac:on
between
the
lone
pair
and
 bonding
pairs).
 –  More
than
one
lone
pair:
 •  lone
pair
(lp)
‐
lone
pair
(lp)
interac:ons
are
more
important
 •  lone
pair
(lp)
–
bonding
pair
(bp)
interac:ons
next
in
 importance.
 •  In
addi'on,
interac'ons
at
angles
of
90°
or
less
are
most
 important,
larger
angles
generally
have
less
influence.
 Example
of
ClF3
 Structure
B
can
be
eliminated
quickly
because
of
90°
lp‐lp
angle.
Structure
C
favored
over
A
 because
only
four
90°

lp‐bp
interac:ons
for
C
compared
to
six
such
interac:on
in
structure
A
 Experimental
data
for
ClF3
 E
:
central
atom
 X
:
coordinated
atom
 L
:
lone
pair
 6‐coordinate
 EX6
 Θ
=
90°
 What
is
the
shape
of
the
molecules?
 –  H2O 
 –  NH3 
 –  NH4+ –  CH4 
 –  CO32‐ –  F 2 
 –   O2 
 –  BH3 
 –  PCl5 
 –  H2SO4 –  XeO2 –  XeF2
 –  XeF4
 –  SF6 
 
bent
 
trigonal
pyramidal
 
tetrahedral
 
tetrahedral
 
trigonal
planar
 
linear
 
linear
 
trigonal
planar
 
trigonal
bipyramidal
 
tetrahedral
 
bent 

 
linear 

 
square
planar
 
octahedral
 Mul:ple
bonds
 •  The
VSEPR
model
considers
double
and
 triple
bonds
to
have
slightly
greater
 repulsive
effects
than
single
bonds
 because
of
the
repulsive
effect
of
the
π
 electrons.
 •  Also
mul:ple
bonds
like
lone
pairs
tend
to
 occupy
more
space
than
single
bonds
and
 to
cause
distor:on.
 Trigonal
bipyramidal
geometry
 •  There
are
two
possible
loca:ons
of
lone
pairs
:
axial
 and
equatorial.
 –  If
there
is
a
single
lone
pair,
the
lone
pair
occupies
an
 equatorial
posi:on
(lone
pair
has
the
most
space
and
 minimize
the
interac:on
between
the
lone
pair
and
 bonding
pairs).
 –  More
than
one
lone
pair:
 •  lone
pair
(lp)
‐
lone
pair
(lp)
interac:ons
are
more
important
 •  lone
pair
(lp)
–
bonding
pair
(bp)
interac:ons
next
in
 importance.
 •  In
addi'on,
interac'ons
at
angles
of
90°
or
less
are
most
 important,
larger
angles
generally
have
less
influence.
 Example
of
ClF3
 Structure
B
can
be
eliminated
quickly
because
of
90°
lp‐lp
angle.
Structure
C
favored
over
A
 because
only
four
90°

lp‐bp
interac:ons
for
C
compared
to
six
such
interac:on
in
structure
A
 Electronega:vity
and
Atomic
size
effects
 •  If
the
central
atoms
remains
the
same,
molecules
that
 have
a
larger
difference
in
electronega:vity
values
 between
their
central
and
outer
atoms
have
smaller
 bond
angles.
 •  The
atom
with
larger
electronega:vity
draws
the
 shared
electrons
toward
itself
and
away
from
the
 central
atom,
reducing
the
repulsive
effect
of
these
 electrons.
 Electronega:vity
and
Atomic
size
effects
 •  As
the
central
atom
becomes
more
electronega:ve,
it
pulls
 electrons
in
bonding
pairs
more
strongly
toward
itself,
 increasing
the
concentra:on
of
electrons
near
the
central
 atom.
 •  The
net
effect
is
an
increase
in
bonding
pair‐bonding
pair
 repulsions
near
the
central
atom
increasing
the
bond
 angles.
 •  Molecule
with
the
most
electronega've
central
atom
has
 the
largest
bond
angles.
 Effects
of
size
 •  generally,
electronega:vity
and
size
goes
hand
to
 hand.
 •  the
most
electronega:ve
atoms
have
also
been
the
 smallest.
 •  However
if
the
size
of
the
outer
atoms
or
groups
 becomes
more
important
than
the
 electronega:vity,
opposite
effect.

 trigonal
bipyramidal
geometry

 •  Central
atom
to
axial
distances
are
longer
than
the
distances
to
equatorial
 atoms.
This
effect
has
been
aoributed
to
the
greater
repulsion
of
lone
and
 bonding
pairs
with
atoms
in
axial
posi:ons
(three
90°
interac:ons)
than
 with
atoms
in
equatorial
posi:ons
(two
90°
interac:ons).



 •  tendency
for
less
electronega:ve
groups
to
occupy
equatorial
posi:on,
 similar
to
lone
pairs
and
mul:ple
bonded
atoms.
 Dipole
moment
 •  Homonuclear
diatomic
renders
bond
non‐polar.
 •  For
heteronuclear
diatomic,
difference
in
 electronega:vity,
the
bonding
electrons
are
drawn
 closer
towards
the
more
electronega:ve
atoms.
 δ+
 H
 δ‐
 F
 Molecular
Dipole
Moment
 •  Polarity
is
a
molecular
property.
 •  Determine
the
net
dipole
moment,
it
depends
 on
the
magnitudes
and
rela:ve
direc:ons
of
 all
the
bond
dipole
moments
in
the
molecule.
 •  Example:
CCl4,
H2O,
NH3.
 ...
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This note was uploaded on 10/03/2011 for the course CHEM 113A taught by Professor Professornotknown during the Spring '09 term at San Jose State University .

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