Liquids
Polar bonds and dipoles
Intermolecular forces
Liquid properties
Phase changes
Evaporation, vapour pressure and boiling point
Clausius-Clapeyron equation
Intermolecular forces
the sequence gas → liquid → solid
 Intermolecular attractions increase
 In



Gases – essentially no interactions
Liquids – movement allowed
Solids – completely rigid
Polarity redux

Electronegativity differences between atoms
creates polar bonds – the more electronegative
atom attracts the electrons
Molecular dipole
 Molecules
are assemblies of several
bonds
 Molecular polarity depends on the
orientation of the individual dipoles
 If the dipoles cancel out, molecular is nonpolar
 If the dipoles don’t cancel, the molecule is
polar
Symmetry and polarity
Studying bonds is an
approximation

We can calculate the centers of gravity of the negative
and positive charges in a molecule


If they do not coincide, the molecule is polar
These calculations are involved, so studying individual
bonds is a good approximation
-
+
Dipole moments

The dipole moment is the charge x length of the
dipole
  Qr

An electron and proton separated by 0.1 nm (a
typical bond length)
  1.6 x10

19
9
x0.1x10  1.6 x10
29
Cm  4.8D
Where 1 D (Debye) = 3.336 x 10-30 Cm
Algorithm for predicting molecular
polarity







Establish molecular skeleton
Draw Lewis dot structure
Count groups of charge around central atom
Establish electronic geometry using VSEPR
Determine molecular shape
Identify polar bonds and lone pairs
Inspect molecule: do polar bonds/lone pairs
cancel out?
Percent ionic character

We have seen that we can calculate the dipole
moment for a given charge separation
 Comparison with experimental values permits
estimation of “ionic character”



In HCl the experimental dipole moment is 1.03 D.
The theoretical dipole given the bond length of 0.127
nm is 6.09 D
Percent ionic character = 1.03/6.09 x 100 % = 16.9 %
May the force be with you
 Covalent
and ionic bonds are the
intramolecular forces that hold the atoms
in molecules together
 Intermolecular forces hold the molecules
together
 Collectively, the intermolecular forces are
called van der Waals forces
 All arise from electrostatic interactions
Name of
force
Ion-dipole
Origin
Strength
Between ions
and
molecules
Between
permanent
dipoles
Quite strong
(10 – 50
kJ/mol)
Hydrogen
bonds
Polar bonds
with H and
(O,N)
Quite strong
(10 – 40
kJ/mol)
London
dispersion
forces
Fluctuating
dipoles in nonpolar bonds
Weak (1 –
10 kJ/mol)
Dipoledipole
Weak (3 –
4 kJ/mol)
Ion - dipole
 Characteristic
of interactions in solutions
of ionic compounds in polar solvents


Negative ion with the positive dipole end
Positive ion with the negative dipole end
Dipole - dipole
 Important
attractive force in polar
substances
 Strength of the order of 3 – 4 kJ/mol
(compared with 200 – 400 kJ/mol for
covalent bonds)
Manifested in boiling points:
 Nonpolar
substances have much lower
boiling points

Acetone (polar) 56ºC butane (nonpolar) 0.5ºC
 Boiling
point increases with dipole strength
London calling


Even molecules with no net dipole moment attract each
other.
Electrons are not static but mobile:


Effect increases with atomic number – as atom becomes
more polarizable


Fluctuation creates dipole in one molecule which induces dipole
in another molecule
Boiling increases with atomic weight
Conventionally, dispersion forces are said to be weaker
than other inter-molecular forces. For large molecules
this is not really true. Large molecules are solids
because of dispersion forces
Hydrogen bonds: the most
important bond?

Key to life
 Between H and O, N or F
 Dipole-dipole bonds of unusual strength (up to
40 kJ/mol)
Hydrogen bonding

The ultimate expression of polarity
 Small positive H atom exerts strong attraction on
O atom
 Other H-bonding molecules: HF, NH3
 H2O is the supreme example: two H atoms and
two lone pairs per molecule
H2O has optimum combination of
lone pairs and H atoms
Compound
Number of lone Number of H
pairs
atoms
HF
3
1
H2O
2
2
NH3
1
3
H bonding generates threedimensional network
Water: the miracle
 All
the properties of water that make it
unique and life sustaining can be traced to
hydrogen bonding




Density of ice lower than water
Anomalous high b.p.
High heat capacity
Universal solvent
Understanding the force
 Predicting
the forces acting between
molecules means understanding the
molecules
 All molecules experience London forces,
but only some will have dipole-dipole or
hydrogen bonds. Where present, the
latter will dominate
Properties of liquids depend on
intramolecular forces

Water flows but syrup is sticky
 Viscosity measures resistance to flow



Small non-polar molecules flow easily
Large or highly polar molecules flow less easily
Units of viscosity are kg/m-s
Surface tension? Take a tablet

Surface tension is the tendency of a liquid to
resist spreading out
 Arises from molecules at the surface
experiencing inward pull
 Walking on water: it’s no miracle, it’s surface
tension
 Surface tension is the energy required to
increase the surface area of a liquid – units are
J/m2
Cohesive and adhesive
 Cohesive
forces are the attractive forces
between like molecules
 Adhesive forces are the attractive forces
between unlike molecules
Meniscus
 Adhesive
forces pull H2O molecules to
maximize coverage
 Cohesive
forces between H2O molecules
drag liquid up
 Gravity pushes liquid down
Capillary action

Combined effects of cohesive, adhesive and
gravitational forces cause liquid to rise towards
edge of container
 In very thin columns the effect of gravity is
diminished and the liquid rises higher
 Originally used as explanation (incorrect) for
transport of water through plants (Osmosis is the
cause)
Just a phase I’m going through
 A phase
change occurs when matter
changes from one state to another
 Solids can exhibit more than one phase
which also undergo phase changes (gray
tin to white tin)
Energetics of phase changes
 In


the series: solid → liquid → gas
Energy is required to break intermolecular
forces
Distribution of molecules is more disordered
(entropy) – greater disorder is more
favourable
Roadmap of changes
 More
condensed to less condensed
means heat absorption and entropy gain
which are opposing
Phase changes involve “latent” heats

With matter in a single phase, heating the
substance gives a T increase depending upon
S.H.
 At a phase change, two phases are in
equilibrium and heat is absorbed to convert one
into the other without a change in T. Hence the
term “latent” heat – a term no longer in popular
use.
Fusion versus vaporization
 For
all substances, the heat of
vaporization is much larger than the heat
of fusion

More bonds are broken in creating the vapour
Vapour pressure

Liquids do not turn into a vapour only at the
boiling point
 At any temperature, there is vapour in
equilibrium with the liquid



A puddle of water on the sidewalk evaporates
A liquid develops a pressure in a manometer
The pressure exerted by the vapour in
equilibrium with the liquid is the vapour pressure
Maxwell, Boltzmann and vapour
pressure

Molecules exhibit a range of energies, which
moves to higher energy as T increases


More molecules have sufficient energy to escape
liquid as T increases
When the vapour pressure = atmospheric
pressure, the liquid boils
Contrasting volatile and nonvolatile liquids
Property
Volatile liquid
Non-volatile liquid
Cohesive forces
Viscosity
Surface tension
Specific heat
Low
Low
Low
Low
High
High
High
High
Vapour pressure
Rate of
evaporation
High
High
Low
Low
Boiling point
Heat of
vaporization
Low
Low
High
High
Clausius – Clapeyron equation
 The
vapour pressure in equilibrium with a
liquid obeys the following equation


Calculate ΔHvap from vapour pressure data
Calculate vapour pressure as f(T) given ΔHvap and one
vapour pressure value
ln Pvap
 H vap  1
  C
  
R T
