Unit 7
Lewis diagrams
molecular geometry
bond and molecular polarity
IMFAs
Lewis dot diagrams
add up the total number of valence
electrons for all atoms in the molecule
arrange the atoms to pair up the separate
atoms’ single electrons as much as
possible
confirm that:
the total number of electrons exactly matches
the total valence electrons of the original atoms,
and
each atom has an octet of electrons (8), except
H and He have a duet of electrons (2)
structural formulas
also called “Lewis structures” or “Lewis
diagrams” (but not “Lewis dot structures”)
replace each shared pair of electrons with
a solid line representing a covalent bond
consisting of two shared electrons
continue to show the lone pairs of
electrons (which are unshared)
double-check that the lone pairs plus bond
pairs still add up to the correct total
number of valence electrons
multiple bonds
additional bonds may need to be added to
a Lewis structure if
single electrons remain
atoms do not have octets
in simple cases, you may be able to pair
up single electrons on adjacent atoms to
form additional bonds, e.g.
CO2
N2
C2H4
multiple bonds
in other cases, you cannot strictly keep
electrons with their original atoms; the
electrons are free to move elsewhere in
the molecule as needed to complete
octets, e.g.
carbon monoxide, CO
ozone, O3
in these cases, atoms may not form their
“normal” number of bonds
but the total number of valence electrons
must not change; they are just rearranged
multiple bonds
computational approach
you can also calculate exactly how many
bonds are in a molecule in the following
way
add up the valence electrons that the atoms in
the molecule actually have
separately add up the valence electrons those
atoms need in order to have noble gas
configurations
calculate the difference, need – have
that difference is the number of shared
electrons the molecule must have
multiple bonds
computational approach
O2
CO
O O
have: 6 + 6 = 12
need: 8 + 8 = 16
4 shared ethus 2 bonds
C O
have: 4 + 6 = 10
need: 8 + 8 = 16
6 shared ethus 3 bonds
after building the basic skeleton with bonds
add remaining electrons as needed to complete
octets
double-check that the total number of electrons is
exactly the number of valence electrons (“have”)
general hints for Lewis structures
if a given molecule can be drawn with both
symmetrical and asymmetrical structures,
the symmetrical one is more likely to be
correct
central atoms are often
written first in the formula
the least electronegative element
the element that can form the most bonds
hydrogen and halogens
only form one bond, thus are terminal atoms
are generally interchangeable in molecules
exceptions to octet “rule”
most atoms have octets (8 valence
electrons) when in molecules, but there
are
exceptions
group
number of electrons
number of bonds
examples
column 1
duet (2)
1
H2, LiH
column 2
quartet (4)
2
BeH2 , MgI2
column 3
sextet (6)
3
BH3 , AlCl3
octet (8)
4 bonds
3 bonds + 1 lone pair
2 bonds + 2 lone pairs
1 bond + 3 lone pairs
CH4
NH3
H2O
HCl
columns 4-8
molecular shapes: VSEPR model
valence shell electron-pair repulsion
groups of electrons naturally find positions
as far apart from each other as possible
different molecular shapes result based on
how many groups of electrons are present
each of the following counts as one “set”
of electrons around the central atom
a lone pair
a single bond (2 shared e-)
a double or triple bond (4 or 6 shared e-)
VSEPR model—central atom with:
2 sets of e–
3 sets of e–
4 sets of e–
5 sets of e–
6 sets of e–
linear
trigonal planar
tetrahedral
trigonal
bipyramidal
octahedral
e.g. BeF2
e.g. BF3
e.g. CF4
e.g. SF5
e.g. XeF6
electron geometry vs. molecular shape
 each set of electrons occupies a position around
the central atom
 the number of sets defines the electron
geometry
 but lone pairs are essentially transparent
 even though they are invisible, lone pairs make
their presence known by distorting the positions
of the bonds around them (since lone pairs repel
the electrons in the bonds)
 this results in several related molecular shapes
within each general class of electron geometry
tetrahedral electron geometry
4 electron sets
bonds
lone pairs
molecular shape
example
4 single
0
tetrahedral
CH4
3 single
1
triangular pyramid
NH3
2 single
2
bent (~109°)
H2O
1 single
3
linear
HCl
tetrahedral electron geometry
tetrahedral electron geometry
triangular planar electron geometry
3 electron sets
bonds
lone pairs
molecular shape
example
3 single
0
triangular planar
BH3
2 single + 1 double
0
triangular planar
CH2O
1 single + 1 double
1
bent (~120°)
O3
linear electron geometry
2 electron sets
bonds
lone pairs
molecular shape
examples
2 single
0
linear
BeH2
2 double
0
linear
CO2
1 single + 1 triple
0
linear
HCN
in addition, any diatomic molecule must be
linear (since any two points lie on a line)
triangular planar and linear
electron geometry
bond polarity
two electrons shared between two atoms
form a covalent bond
if those electrons are shared equally (or nearly
equally), it is a non-polar covalent bond
if one atom attracts the electrons much more
strongly than the other atom, it is a polar covalent
bond
if one atom completely removes an electron from
the other atom, the result is an ionic bond
bond polarity
the electronegativity difference between
the two atoms determines how polar a
bond
ΔEN, electronegativity difference
bondis
type
non-polar
polar
ionic
Cℓ2
0.0 – 0.4
0.5 – 1.7
> 1.7
HCℓ
LiCℓ
bond polarity
dipole moment is the actual measureable
quantity related to bond polarity
the size of the dipole moment is affected
by
electronegativity difference
bond length
we will focus on ΔEN and a qualitative
sense of bond polarity
molecular polarity
 the overall polarity of a molecule depends on the
combined effect of the individual polar bonds
individual bonds polar
individual bonds polar
overall molecule
nonpolar
overall molecule
polar
molecular polarity
what allows bond dipoles to cancel?
geometric symmetry of the molecule
having identical terminal atoms (or
atoms with the same electronegativity)
what prevents bond dipoles from
canceling?
geometric asymmetry (due to lone
pairs)
having different terminal atoms
molecular polarity
molecular polarity
 inherently symmetrical
shapes (if all
surrounding atoms are
the same)
tetrahedral
triangular planar
linear
 inherently asymmetrical
shapes
bent
triangular
pyramid
 even symmetrical shapes
become
asymmetrical if different terminal atoms are
IMFA: intermolecular forces of
attraction
“mortar”—
holds the
separate
pieces
together
(i.e. the
IMFA)
“bricks”— individual atoms, ions, or molecules of a solid
IMFA: intermolecular forces of
attraction
types of IMFA
strongest
covalent network
occurs between
atoms such as C, Si, & Ge
(when in an extended grid or network
ionic bond
cations and anions
van der Waals forces
(metals with non-metals in a salt)
metallic bond
metal atoms
hydrogen bond
ultra-polar molecules
(those with H–F, H–O, or H–N bonds)
weakest
dipole-dipole attraction
polar molecules
London forces
non-polar molecules
IMFAs – Trends and Characteristics
melting points and boiling points
Solubility
Conductivity (type specific)
Strength (type specific)
Melting Points and Boiling Points
stronger IMFAs cause higher m.p. and b.p.
atoms and molecules that are heavier
and/or larger generally have higher m.p.
and higher b.p.
larger e– clouds can be distorted (polarized)
more by London or dipole forces, causing
greater attraction
CH3CH2CH2CH2CH3 > CH3CH2CH3
Polarity (for dipole dipole and H-bonds)
more polar=higher b.p, m.p.
HCl > HI
Melting Points and Boiling Points
strategy to predict m.p. and b.p.
first sort atoms/molecules into the six IMFA
categories
then sort within each IMFA category from
lightest to heaviest (or least polar to most
polar)
same IMFA: sort by molar mass
melt
boil
+184.4
(257)
+150
I2
+113.7
(257)
+100
+58.8
–7.2
(160)
Cℓ2
–101.5
(71)
–200
F2
(38)
–219.62
–34.04
Cℓ2
(71)
–100
–150
–250
°
C
 ex: halogen family
 all are non-polar (London force)
 lowest to highest m.p. and b.p.
matches lightest to heaviest
 thus at room temperature:
0
–50
Br2
(160)
+50
Br2
I2
–182.95
F2
(38)
 F2 (g)
 Cℓ2 (g)
 Br2 (ℓ)
 I2 (s)
same mass: sort by IMFA type
+97.4
 ex: organic
(can form twice as many H-bonds)
molecules
 all are ~60 g/mol
 different types of
1-propanol (ultra-polar = H-bonds) IMFA
+56.2
acetone (more polar)
+10.8
methyl ethyl ether (slightly polar)
butane (non-polar)
+198
+150
+100
+50
0
–50
°
C
–0.5
ethylene glycol
 the stronger the
IMFA, the higher the
boiling point
Solubility
substances generally mix best with others
of similar IMFAs
”like dissolves like”
non-polar mixes well with non-polar
polar mixes well with polar (and ultra-polar / ionic)
other physical properties
strength, conductivity, etc. are related to
the type of IMFA
details about each IMFA
strongest
covalent network
ionic bond
metallic bond
hydrogen bond
dipole-dipole attraction
London forces
weakest
London (or dispersion) forces
 non-polar molecules (or single atoms)
normally have no distinct + or – poles
 electron clouds are slightly distorted by
neighboring molecules
sort of like water sloshing in a shallow pan
 form temporary dipoles
 Low MP and BP
 soluble
London dispersion forces in action
δ+
δ-
1. temporary
polarization
due to any
random little
disturbance
2. induced
polarization
caused by
neighboring
molecule
3. induced
polarization
spreads
4. induced
polarization
non-polar molecules, initially with uniform charge distribution reverses
dipole-dipole attractions
 polar molecules have permanent dipoles
 the molecules’ partial charges (δ+, δ-) attract the
oppositely-charged parts of neighboring
molecules
 this produces stronger attraction than the
temporary polarization of London forces
therefore polar molecules are more likely to be liquid at
a temperature where similar non-polar molecules are
gases
hydrogen bonding (or ultra-dipole attractions)
 H—F, H—O, and H—N bonds are more polar than
other similar bonds
 Very small atoms, particularly H
 F, O, and N are the most electronegative elements
 particularly polar
 molecules containing these bonds have much
higher m.p. and b.p than otherwise expected for
non-polar or polar molecules of similar mass
 the geological and biological systems of earth
would be completely different if water molecules did
not H-bond to each other
hydrogen bonding (or ultra-dipole attractions)
ultra-polar molecule
(much higher boiling
point)
non-polar molecules
(lower boiling points)
hydrogen bonds
(between molecules,
not within them)
hydrogen bonding (or ultra-dipole attractions)
Beware!!
H
H
H
O
H
H
O
H
O
H
H
O
These are not
hydrogen bonds. They
are normal covalent
bonds between
hydrogen and oxygen.
These are hydrogen
bonds. They are
between separate
molecules (not within a
molecule).
metallic bonding
 structure
nuclei arranged in a regular grid or
matrix
“sea of electrons”—delocalized valence
electrons free to move throughout grid
 resulting properties
conductive (electrically and thermally)
strong, malleable, and ductile shiny
Form alloys = mixture of metals
 Bronze = copper and tin
 Brass = copper and zinc
 Steel = iron and carbon
metallic “bond” is generally weaker
than covalent bond since there are
not specific e– pairs forming bonds
ionic bonding (salts)
structure: orderly 3-D array (crystal)
of alternating + and – charges
made of
cations (metals from left side of periodic table)
anions (non-metals from right side of periodic
table)
properties
non-conductive when solid
conductive when melted or dissolved
hard but brittle (why?)…
why are salts hard but brittle?
1. apply some force
2. layer breaks off and shifts
3. + repels +
– repels –
4. shifted layer
shatters away
from rest of
crystal
covalent networks
 strong covalent bonds hold together millions of
atoms (or more) in a single strong particle
 properties
very high melting temperatures
usually non-conductive (except graphite)
very hard, very strong
 examples
carbon (two allotropes: diamond, graphite)
pure silicon or pure germanium
SiO2 (quartz or sand)
other synthetic combinations averaging 4 e– per atom:
SiC (silicon carbide), BN (boron nitride)
m.p. = 3550°C
m.p. = ~1600°C
C60
buckminsterfullerine
“bucky ball”
summary of properties
strongest
strength
m.p. & b.p.
conductive?
network
extremely hard
very high
ionic
hard but brittle
medium to high if melted or dissolved
usually not
(mobile ions)
van der Waals forces
metallic
strong,
malleable,
ductile
medium to high very
soft and brittle
low
(delocalized e–)
hydrogen
dipole
London
weakest
no
Metallic
London
Ionic
---------Hydrogen
Metallic
Ionic
Metallic
Network
Network
Metallic
Network
Metallic
Hydrogen
Ionic
Covalent
D
F
H
B
G
J
C
I
E
A
To do’s:
Objective:
Prepare for test tomorrow.
Review HW
Characterize each of the 6 IMFA
Practice questions
Independent practice
Practice tests/homework
Ask questions
MP
BP
Covalent network
C, Si, & Ge
ionic
Metal + nonmetal
Metallic
Elements/alloys left
side of periodic table
Hydrogen
bonded
Conductivity
Not
conductive
Conductive in
solution or
molten state
(mobilized ions)
Good
conductor
of heat and
electricity
Not
conductive
Solubility
mostly
Insoluble
Soluble
in polar
solvents
Strength
Very
strong,
hard
Hard
brittle
mostly
insoluble
Strong,
malleable,
ductile
Soluble
in polar
solvents
Soft,
brittle
Soluble
in polar
solvents
Soft,
brittle
H bonded to O, N, F
Dipole-dipole
Polar molecules
London Forces
non-polar molecules
Not
conductive
Not
conductive
Soluble in
non-polar
solvents
Soft,
brittle
soaps and emulsifiers
some molecules are not strictly polar or non-polar, but
have both characteristics within the same molecule
oil
polar
region
water
this kind of molecule can function
as a bridge between molecules
that otherwise would repel each
soaps and emulsifiers
with a soap or emulsifier present to surround it, a drop of
non-polar oil can mix into polar water
 Practice Problems…
 Draw 3 examples of molecules with polar
bonds that overall are not considered to
be polar molecules.
Draw the Lewis structure for AsH3.
Determine the shape and overall polarity.
Which of the following is most likely to
dissolve NaCl?




BH3
CH4
NH3
BeH2
Draw 3 isomers of C5H12
Review for test tomorrow:
Quiz 2
HW # 13, 14, 16,17
Practice test
Review questions from slides