TOPIC 4: CHEMICAL
BONDING & STRUCTURE
ESSENTIAL IDEA
The physical properties of molecular
substances result from different types of
forces between their molecules.
NATURE OF SCIENCE (2.2)
Obtain evidence for scientific theories by making and testing
predictions based on them – London dispersion forces and
hydrogen bonding can be used to explain special interactions. For
example, molecular covalent compounds can exist in the liquid
and solid state. To explain this, there must be attractive forces
between their particles which are significantly greater than those
that could be attributed to gravity.
INTRAMOLECULAR FORCES
 Bonding forces that exist within each molecule.
 Affect molecular geometries, physical properties and
reactivity.
INTERMOLECULAR FORCES
 Non-bonding force that exist between the
molecules.
 The strength of the intermolecular forces
influence physical properties such as volatility,
melting and boiling points
 The stronger the forces, the higher the melting and
boiling points.
 Intermolecular forces are much weaker than
intramolecular forces!
UNDERSTANDING/KEY IDEA
4.4.A
Intermolecular forces include
London dispersion forces, dipoledipole forces and hydrogen
bonding.
TYPES OF INTERMOLECULAR FORCES
 London Forces (Dispersion Forces)
 Dipole-dipole Forces
 Hydrogen bonds
 Van der Waals Forces – include London and dipole-
dipole forces
LONDON DISPERSION FORCES
 These are the weakest of the intermolecular forces.
 Occurs between ALL molecules, polar and non-polar
 Substances held together with London dispersion forces
have low melting and boiling points and are often gases at
room temperatures.
 These forces occur when electron clouds shift to form
temporary dipoles which then induce dipoles in
neighboring molecules so that they can attract each other.
 The larger the electron cloud, the stronger the London
dispersion forces because there is a higher probability of
temporary dipoles forming.
LONDON (DISPERSION) FORCES
www.studybite.com
LONDON DISPERSION FORCES
 Why does the temporary dipole occur?
 Electrons move randomly within the electron cloud
 When electrons move to one region of the cloud they may
form a temporary dipole where one region of that
atom/molecule has a area of negative charge and the end
having a positive charge
 The temporary dipole will cause the electrons in another
molecule to be repelled (or move away). This will cause an
induced dipole (since the electrons moved away, that end
of the molecule is more +, while the other end becomes
more -)
 The net result is that there is on average a stronger attractive
force between molecules than repulsive.
LONDON FORCES
 London dispersion forces are responsible for the fact
that non-polar molecules can be condensed to form
liquids and sometimes solids.
 London dispersion forces are also present in polar
molecules but are often overlooked because they are so
much weaker than dipole-dipole forces.
umdberg.pbworks.com
FACTORS AFFECTING STRENGTH
OF LONDON FORCE
Number of electrons
How will more electrons affect polarizability?
More electrons  lower electrostatic attraction between
valence e- and nucleus
Easier to polarize  stronger London force
Stronger intermolecular forces meaning higher
boiling points
Size (volume) of Electron Cloud
Larger e- cloud ( molecule ∴  molecular mass)
Less attraction of e- to nucleus
Easier to polarize  stronger London force
Larger e- cloud (molecular mass) = higher boiling pts
FACTORS AFFECTING STRENGTH
OF LONDON FORCE
Shape of Molecules
How will the shape of a molecule affect it’s interactions
with other molecules?
• More linear  easier to access
surface so more interactions
• More clustered molecules (balllike) allows less surface area =
less interactions between
molecules
• More interactions between
molecules = higher boiling
points
http://www.slideshare.net/Hoshi94/states-of-matter-11767927
DIPOLE-DIPOLE FORCES
 These are stronger than London dispersion forces and
the strength depends upon the degree of polarity.
 These intermolecular forces are caused when
molecules with permanent dipoles (areas of pos/neg
charge) attract each other.
 The stronger the dipole attraction – the higher the
boiling point.
DIPOLE-DIPOLE FORCES
 Dipole = polar molecule
 Dipoles will change their
direction so that their
oppositely charged ends are
near to one another.
 The electrostatic attraction
between the ends is dipoledipole force
Dipole-dipole
force
HYDROGEN BONDING
 This is the strongest of the intermolecular forces .
 Hydrogen bonding is a type of dipole-dipole
attraction.
 Hydrogen bonds form when hydrogen bonds to either
nitrogen, fluorine or oxygen.
 The strength of the hydrogen bond is due to the small
size of hydrogen and the large electronegativity of N,
O and F.
 Water is a unique substance with hydrogen bonding in
that the solid is less dense than the liquid.
HYDROGEN BONDING
HYDROGEN BONDING
 Hydrogen bonding results between the interaction of the
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lone pair of one atom (small and highly electronegative,
like N, O or F) with the hydrogen atom of another molecule
The molecule involved with hydrogen bonding should have
at least one lone pair.
More lone pairs= stronger hydrogen bond
The more electronegative the atom the stronger the
hydrogen bond
Part way between a dipole-dipole bond and a dative
covalent bond
H- bonds have large effect on properties and structure of
materials
UNDERSTANDING/KEY IDEA
4.4.B
The relative strengths of these
interactions are London dispersion
forces < dipole-dipole forces <
hydrogen bonds.
Intermolecular Forces Summary
IM Force
General Description
Relative
Strength
(kJ/mol)
Effect on BP
London
All molecules,
Temporary dipoles
from uneven edispersion
Weak
1-10
Increases with # e, volume of ecloud, Shape (
S.A)
DipoleDipole
Electrostatic attraction
between dipoles
Weak to
Moderate
3-25
Increases with
larger polarity
(E.N. diff.)
Hydrogen
Bond
H bonded to N,O, or F
creating high dipole
moment. H+ attracted
to lone pairs of another
molecule
Moderate to
Strong
(10-40)
Increases with
more lone pairs
APPLICATION/SKILLS
Be able to deduce the types of
intermolecular forces present in
substances, based on their structure
and chemical formula.
GUIDANCE
The term “London dispersion forces”
refers to instantaneous induced dipoleinduced dipole forces that exist between
any atoms or groups of atoms and should
be used for non-polar entities. The term
“van der Waals forces” is an inclusive term,
which includes dipole-dipole, dipoleinduced dipole and London dispersion
forces.
Intermolecular Forces
 All molecules will have some type of van der Waal’s
force.
 Non-polar molecules have only London dispersion
forces.
 Polar molecules have dipole-dipole forces and London
dispersion forces.
 Hydrogen bonding exists when the positive hydrogen
bonds with lone pairs of electrons on nitrogen, oxygen
and fluorine.
Van der Waals’ forces
 The umbrella term “van der Waals’ forces is used to
include both London dispersion forces and dipoledipole attractions. It also covers the less common type
of attraction known as the dipole-induced dipole.
 It refers to all forces between molecules that do not
involve electrostatic attractions between ions or bond
formation.
 London Dispersion Force: Cl2 --- Cl2
 Dipole – Dipole attraction: HCl --- HCl
 Dipole-induced dipole:
HCl --- Cl2
APPLICATION/SKILLS
Be able to explain the physical
properties of covalent compounds
(volatility, electrical conductivity,
and solubility) in terms of their
structure and intermolecular forces.
Using Forces to Predict Properties
 Volatility, solubility and conductivity can all be
predicted and explained from knowledge of the nature
of the forces between molecules.
VOLATILITY
 How easily a substance evaporates (becomes a gas)
 The weaker the intermolecular force, the higher the
volatility (easier for atoms to move apart)
 London dispersion > dipole-dipole > hydrogen
bonding
 Ionic compounds and giant covalent compounds have
low volatility.
 Ionic, giant covalent< polar covalent < non polar
ELECTRICAL CONDUCTIVITY
 Covalent compounds do not contain ions; therefore,
they do not conduct electricity in the solid or liquid
state.
 Some polar covalent compounds, such as HCl which
can ionize in water, will conduct electricity.
 Ionic compounds conduct electricity in the molten or
aqueous state.
 Giant covalent structures are generally non-conductors
except for graphite, graphene, fullerene and Si.
SOLUBILITY
 Non-polar compounds dissolve in non-polar solvents.
 Polar compounds dissolve in polar solvents.
 Solubility is reduced in larger molecules where the
polar bond is only part of the total structure.
 Ionic compounds are very soluble in water and nonsoluble in non-polar solvents.
 Giant covalent structures are non-soluble in both polar
and non-polar solvents due to the very strong covalent
bonds within their structure.
BOILING/MELTING POINT COMPARISON
 The stronger the forces between atoms/molecules, the
more energy required to break atoms/molecules apart,
the higher the boiling and melting points.
 It is easiest to compare boiling points for similar
substances.
 To predict boiling points, you must consider the type of
intermolecular forces involved.
 To identify the force:
 Draw Lewis structures
 Determine if molecule is polar or non-polar
 If polar, determine if there are any H-bonds
BOILING/MELTING POINT COMPARISON
 After the forces are identified you:
 Compare the size of the molecule (Mr) for
London Forces
 For Dipole-dipole look at difference in EN of
atoms
 If H-bonds, how many lone pair electrons?
 Hydrogen bonding > dipole-dipole > London
dispersion forces
 The more intermolecular forces, the stronger the
interaction and the higher the boiling points
EXAMPLE PROBLEMS
Identify the intermolecular forces in the following
substances: He, CH3(CH2)4CH3, NF3, (CH3)2O, CH3F,
CH3CH2OH
Draw Lewis structures as needed
 He – London forces only
 CH3(CH2)4CH3 symmetrical = non-polar so
London forces only
EXAMPLE PROBLEMS
Identify the intermolecular forces in the following
substances: He, CH3(CH2)4CH3, NF3, (CH3)2O, CH3F,
CH3CH2OH
 NF3 Polar so London and Dipoledipole forces
 (CH3)2O Polar so London and
Dipole-dipole forces
EXAMPLE PROBLEMS
Identify the intermolecular forces in the following
substances: He, CH3(CH2)4CH3, NF3, (CH3)2O, CH3F,
CH3CH2OH
 CH3F Polar so London and Dipoledipole forces
 CH3CH2OH Polar w/ Hbond so London, Dipoledipole forces, and Hbond
EXAMPLE PROBLEMS
Place sulfur (S8), chlorine, and argon in
order of increasing boiling points. Explain
your order.
 Argon<Chlorine<Sulfur
 All non-polar so only force is London force
 Sulfur (S8) has highest B.P. due to it being the
largest molecule (molecular mass) and ∴ the
strongest London force, followed by Cl2 and then
Ar
EXAMPLE PROBLEMS
Compare the boiling points of propane (CH3CH2CH3),
methoxymethane (CH3OCH3) and ethanol (CH3CH2OH)
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CH3CH2CH3< CH3OCH3 < CH3CH2OH
CH3CH2CH3 non-polar London only
CH3OCH3 polar – London and dipole-dipole
CH3CH2OH London, dipole-dipole and H-bond
EXAMPLE PROBLEMS
 Melting point is a measure of the difference in strength of
the forces between particles between the solid and liquid
state. Order the following from lowest to highest melting
point: HI, HCl, HBr, HF.
 HF<HCl<HBr<HI
 M.P for HF is -118 oC, for HCl is -114.2 oC, for HBr is -86.8
oC, for HI is -50.80 oC
 The larger the atomic size, the more energy it requires to melt the
molecule.
 The larger the size, the more electrons there are moving randomly
around, so the more likely to polarize and probably stronger London
forces occurring between each molecule.
Let’s See If You Have It
1. Which attractions are stronger: intermolecular or
2.
3.
4.
5.
intramolecular?
Suggest some ways that London forces are different
from dipole-dipole forces.
Which would have a lower boiling point: O2 or F2?
Explain.
Which would have a lower boiling point: NO or O2?
Explain.
Which would you expect to have the higher melting
point (or boiling point): C8H18 or C4H10? Explain.
Let’s See If You Have It
6. What two factors causes hydrogen bonds to be so
much stronger than typical dipole-dipole bonds?
7. What kind(s) of intermolecular forces are present in
the following substances: NH3, SF6, PCl3, LiCl, HBr,
CO2 (hint: consider EN and molecular
shape/polarity)
Let’s See If You Have It
1. Intramolecular are stronger.
2. London forces
Are present in all compounds
Can occur between atoms or molecules
Are due to electron movement not to EN
Are transient in nature (dipole-dipole are more
permanent).
 London forces are weaker
F2 would be lower because it is smaller. Larger
atoms/molecules can have their electron clouds more
easily deformed and thus have stronger London
attractions and higher melting/boiling points.
O2 because it has only London forces. NO has a small
EN, giving it small dipoles.
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3.
4.
Let’s See If You Have It
5.
6.
7.
C8H18 would have the higher melting/boiling point. This is
a result of the many more sites available for London
forces to form (larger molecule)
a large EN and the small sizes of atoms
NH3: Hydrogen bonding (H + N), dipole-dipole London.
SF6: London only (it is symmetrical & non-polar).
PCl3: EN=2.9-2.1= 0.8. Dipole-dipole, London.
LiCl: EN=2.9-1.0=1.9 Ionic, (London).
HBr: EN=2.8-2.1=0.7. Dipole-dipole, London.
CO2: London only (it is symmetrical & non-polar)
Let’s See If You Have It
5.
6.
7.
C8H18 would have the higher melting/boiling point. This is
a result of the many more sites available for London
forces to form (larger molecule)
a large EN and the small sizes of atoms
NH3: Hydrogen bonding (H + N), dipole-dipole London.
SF6: London only (it is symmetrical & non-polar).
PCl3: EN=2.9-2.1= 0.8. Dipole-dipole, London.
LiCl: EN=2.9-1.0=1.9 Ionic, (London).
HBr: EN=2.8-2.1=0.7. Dipole-dipole, London.
CO2: London only (it is symmetrical & non-polar)
Reading
Pages 122 - 132
Look over worked examples