Properties of Liquids
College Chemistry
KINETIC MOLECULAR THEORY OF GASES

Back when we studied gases, we used the KMT to
explain the behavior of gases


Distance between molecules are so great that at
ordinary temperature and pressure, there is no
interaction between molecules
Lots of empty space gives gases lots of their
properties
Compressibility
 Expandability
 Low densities

Kinetic-Molecular Theory Applied to
Liquids and Solids
Condensed states – liquids and solids, referred to
as this because these states have much higher
densities than they do in the gaseous state
 We have already looked at how the kinetic-molecular
theory can be applied to gases, how can it be applied
to liquids and solids?

Kinetic-Molecular Theory Applied to
Liquids and Solids
Gas
Liquid
Solid
Highly
compressible
Slightly
compressible
Very little
compressibility
Low density
High density
High density
Fills container
completely
Does not expand to
fill – has a definite
volume
Rigidly retains its
volume
Assumes shape of
container
Assumes shape of
container
Retains own shape
Rapid diffusion
Slow diffusion
Extremely low
diffusion
High expansion on
heating
Low expansion on
heating
Low expansion on
heating
KINETIC-MOLECULAR THEORY APPLIED
TO SOLIDS AND LIQUIDS
Densities of solids are usually (except for water)
higher than liquids and it is possible for multiple
phases of matter to exist at the same time
 Phase – homogeneous part of the system in
contact with other parts of the system but
separated by a well-defined boundary


Ex: a glass of water with ice in it
FORCES AND BONDS BETWEEN
GASES AND LIQUIDS

Intramolecular forces– chemical bonds, hold
atoms together
WITHIN a molecule
 Both covalent and ionic bonds


Intermolecular forces – exist BETWEEN
molecules
General name: van der Waals forces
 very weak, but very important for life

INTERMOLECULAR VS. INTRAMOLECULAR
FORCES
Intramolecular forces – hold atoms together
(ionic, covalent, or metallic bonds)
 Intermolecular forces – attractive forces that
hold molecules together



Generally much weaker than intramolecular forces,
but define the properties of the liquids nonetheless
We can define these intermolecular forces overall
as van der Waals forces, named after the
Dutch physicist Johannes van der Waals

All the following intermolecular forces we will talk
about are van der Waals forces, except for hydrogen
bonding
DIPOLE - Review
In simplistic terms, the shift of electron density
to one side of the compound (the most
electronegative side)
 Polar covalent bonds usually exhibit dipole
behavior fairly well

Not ionic enough to pull all electrons towards them
like in ionic
 Not covalent enough to fully share electrons

http://www.grandinetti.org/Te
aching/Chem121/Lectures/VS
EPR/
ION-DIPOLE ATTRACTIVE FORCES
An ion is attracted to a polar molecule
 An anion is attracted to the partial positive side
 A cation is attracted to the partial negative side
 This “splits up” the ionic compound  soluble
 This force becomes stronger as either the ion has
a larger charge or the polar molecule becomes
stronger (dipole)

http://www.chem.purdue.edu/gche
lp/liquids/iondip.html
ION-DIPOLE FORCE
One example of an ion-dipole interaction is
hydration (interaction of cations and/or anions in
water)
 Ions with a larger charge interact stronger with
water because they have a stronger dipole
 Ex: Na+ vs. Mg+2

DIPOLE-DIPOLE FORCES
Occurs between polar covalent molecules with
permanent dipoles
 Larger the dipole, the larger the force
 The dipole of the molecule forms an attractions
with the dipole of the water molecule
 Opposites attract!
 some covalent molecules will still not form
conducting solutions (too strong of bonds)
 Very weak

http://www.chem.purdue.edu/gchelp/liquids/dipdip.html
DISPERSION FORCES
A normal atom is purely symmetrical and the net
charge of the atom is zero
 A temporary dipole can form that causes a
separation of the positive and negative charges of the
atom
 This temporary dipole can force the neighboring
atoms to become distorted and form induced
dipoles



A dipole created by the presence of a neighboring dipole
Results in a weak dispersion force that can hold
atoms together

Ex: at high temperatures, dispersion forces hold noble
gases together and become condensed to liquids
DISPERSION FORCES
• Dispersion forces are
THE only force
between nonpolar
molecules
• Dispersion forces
happen between ALL
MOLECULES!
• WEAKEST type of
force

http://www.chem.purdue.edu/gchelp/liquids/disperse.html
DISPERSION FORCES

At low enough temperatures (and reduced atom
speeds), these dispersion forces are enough to
hold atomic gases together causing the gas to
condense
DISPERSION FORCES
Strength of dispersion forces
depends on the charge of the ion
(or the strength of the dipole)
just like ion-dipole interactions
 However, polarizability of the
molecule also plays a huge role



Polarizability – ease with which
the electron distribution in the
atom can be distorted (think of a
balloon)
Usually the larger the # of
electrons and the more diffuse
the electron cloud is, the more
polarizable it is (larger molar
mass)
DISPERSION FORCES

Dispersion (or London) forces usually increase
with molar mass b/c molecules with larger molar
masses tend to have more electrons
Dispersion forces increase in strength with the
number of electrons
 Also, electrons are more easily disturbed b/c outer
electrons aren’t held as strongly


Usually the larger the molar mass (stronger the
dispersion forces), the greater the melting/boiling
point
DISPERSION FORCES
Melting Points of Similar Nonpolar Compounds
Compound
Melting Point (°C)
CH4
-182.5
CF4
-150.0
CCl4
-23.0
CBr4
90.0
CI4
171.0
LONDON DISPERSION FORCES
EXAMPLE 11.1

What types of intermolecular forces exist
between the following pairs: (a) HBr and H2S, (b)
Cl2 and CBr4

(a) HBr and H2S
All molecules have dispersion forces
 Ion-dipole? No
 Dipole-dipole? Yes, it is polar


(b) Cl2 and CBr4
All molecules have dispersion
 Ion-dipole? No
 Dipole-dipole? No, not polar

HYDROGEN BONDING

Hydrogen bonding occurs
when an extremely strong
dipole-dipole force occurs
between H and either F, O,
or N

Must occur in very
electronegative atoms!
Why?




Separation of dipoles
It has a misleading name –
NOT a bond, only an
intermolecular force
Hydrogen bonding is why
water has a very high
boiling point compared to
other compounds
Strongest type of bond!
DIPOLE RELATED FORCES
Between
Ion – Dipole
Ion charge and polar molecule
H - bonding
Polar bonded hydrogen and dipole charge
Dipole - Dipole
Ion – induced dipole
Dipole – induced dipole
Induced dipole –
induced dipole
(dispersion)
Decreasing Strength
Type
Two polar molecules
Ion charge and nonpolar molecule
Polar and nonpolar molecule
2 nonpolar molecules
Remember, ALL molecules have this!
Stronger than dipole-related forces: Ionic, metallic, and covalent bonding
Weaker than dipole-related forces: London dispersion forces
Can mix and match terms!
VISCOSITY
Viscosity – “friction” or resistance to motion that
exists between molecules of a liquid when they
move past one another
 Stronger the attraction between molecules
(intermolecular forces) = stronger resistance
(greater viscosity)
 Viscosity increases as temperature increases



This is why people heat up maple syrup – so it flows
quicker
Ex: alcohol and gasoline both “run” a lot faster
than water, and much faster than syrup

Alcohol and gasoline both have weaker
intermolecular forces
SURFACE TENSION
Surface tension – imbalance of intermolecular
forces at the surface of a liquid; these uneven
forces make the surface behave as if it had a tight
film stretched across it
 Surface tension explains the “beading” of
raindrops, the ability of certain insects to walk
across the water, etc.
 Surface tension increases in liquids with stronger
intermolecular forces


water has a VERY strong intermolecular forces
SURFACE TENSION

One example of surface
tension is capillary action


Water rising spontaneously up
a tube (or water traveling up
your jeans after you step in a
puddle)
Two types of forces bring
about capillary action
1. cohesion
 2. adhesion

SURFACE TENSION

Cohesion – intermolecular attraction between
like molecules
Think “co” – “cooperate”
 Example, water molecules


Adhesion – attraction between unlike molecules
If we think of water in a tube, this would be the
attraction between water and the glass tube
 If adhesion is stronger than cohesion, the liquid will
be pulled up
 Mercury however, will not do this

STRUCTURE OF WATER

Most interesting aspect: liquid water has a
higher density than ice

This is why ice floats
Liquid water can form two hydrogen bonds, with
each oxygen approximately tetrahedral bonded to
four hydrogen (two covalent and two hydrogen)
 Ice cannot hydrogen bond because it has a highly
ordered 3D structure and molecules can’t get to
close to one another

STRUCTURE OF WATER
But let’s look at ice as it melts…
 At its melting point, the water has enough KE to
pull free of intermolecular hydrogen bonds

These molecules become trapped in the space of the
3D ice structure
 As a result, there is more molecules per unit volume
in liquid water than ice
 At the same time, thermal expansion takes place
(water expands as its been heating), so the density
decreases
 Therefore, max density of water is actually at 4°C

WHY DO LAKES FREEZE TOP DOWN?
As we just talked about, as temperature of the
water drops near the surface, density increases
 The colder water than sinks to the bottom and
the less dense, warmer water rises to the top
 This continues until the overall water
temperature is 4°C
 Below this temperature, density of water begins
to slowly decrease with decreasing temperature
and it no longer sinks
 Ice will form, but since it is less dense than
water, it will not sink
