Chapter 11: Liquids and Solids and
Intermolecular Forces
Principles of Chemistry:
A Molecular Approach,1st Ed.
Nivaldo Tro
Dr. Azra Ghumman
Memorial University of Newfoundland
States of Matter: Liquids and Solids
11.2
A molecular comparison of Gases, Liquids, and
Solids
11.3 Intermolecular Forces: The forces that hold
condensed phase together
11.5 Vapourization and Vapour Pressure (Excluding
the Clausius-Clapeyron Equation)
11.6 Sublimation and Fusion
11.7 Heating Curve for Water
11.11 Crystalline Solids: The Fundamental Types
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The Phases of Matter
Three states of matter-Gas, Liquid and Solid
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A Molecular comparison of Gases,
Liquids, and Solids
Gases are compressible fluids.
Molecules are widely separated and are in
constant random motion throughout mostly empty
space.
Assume the shape of container
Negligible forces of interaction
Negligible molecular size compared with total
volume (K.T)
Liquids- incompressible fluids.
Molecules are more closely spaced than gases
Assume the shape of container
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Solids
Solids- Molecules or ions are in close contact and
only oscillate or vibrate about fixed positions
definite shape, incompressible and rigid.
Crystalline: particles are arranged in an orderly
geometric pattern e.g. salt and diamonds
Amorphous: particles have no regular geometric
pattern over a long range e.g. plastic and glass
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Changes between Phases
Phase transition: Change of a substance from one
physical state to another
• Caused by change in temperature and or pressure
melting
freezing
vapourization
condensation (LPG)
Increase in pressure favour the dense state i.e.
Condensation of gases
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11.3 Intermolecular Forces
Intermolecular Forces(IMF)- The interactive forces
between the molecules
hold the condense phase together
Normally weak forces than chemical bonds
Determines Physical properties of liquids and
solids
Three types of forces are known to exist between
neutral molecules
Dispersion forces (Fritz W.London) van-der Waals
Dipole- dipole forces
forces
Hydrogen bonding
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Intermolecular Forces
Intermolecular attractions are due to attractive forces
between opposite charges (q1 and q2) separated by a
distance r
Coulomb’s Law
larger the charge, stronger the attraction
longer the distance, weaker attraction
H-bonding especially strong
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Intermolecular Forces
London dispersion forces- Result from temporary
dipole in the molecules due to unequal electron
distribution.
Present in all substances
Dipole–Dipole attractions- Permanent polarity in
the molecules due to their structure leads to
especially strong dipole–dipole attraction
Hydrogen bonding- Occurs in substances
containing hydrogen atoms bonded to certain very
electronegative atoms (O, N, F).
Stronger than van der Waals forces (Dispersion
and dipole-dipole)
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Instantaneous Dipole in He
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Dispersion Forces
The magnitude of the dispersion forces depends on several
factors.
1. Polarizability: the tendency for charge separation to occur in
a molecule.
The larger a molecule, more easily it can be distorted to give
an instantaneous dipole
Polarizability increases with molar mass resulting in strong
dispersion forces
2. Shape of the molecule- Elongated and branched molecules
are easily polarizeable compared to compact structures
more surface-to-surface contact = larger induced dipole
stronger attraction
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Sample Problem
n-Pentane and neopentane (C5H12) has same molar
mass 72.15 g mol-1. But boiling point of n-pentane is
higher than neopentane. Why?
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Effect of Molecular Size
on Size of Dispersion Force
Effect of size on the boiling points
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Dipole–Dipole Attractions
Dipole- dipole force- is an attractive intermolecular
force resulting from the tendency of polar
molecules to align themselves positive end to the
negative end
H Cl
H Cl
Miscibility
Like dissolve like
The permanent dipole adds to the attractive forces
between the molecules.
Note- All molecules have dispersion forces
• Strong IMF’s increase the bp and mp of the
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Practice problem
Choose the substance in each pair with the highest
boiling point.
b)
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Hydrogen Bonding
Hydrogen bonding is a force that exists between a hydrogen
atom covalently bonded to a very electronegative atom, X (F, O
and N) and a lone pair of electrons on a very electronegative
atom,Y in the vicinity.
X-H----:Yone of the following three structures must be present
Only N, O, and F are electronegative enough to leave the hydrogen
nucleus exposed
Hydrogen bond
is not a chemical bond
Type of strong dipole-dipole force
strongest of three intermolecular forces
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Boiling point versus molecular weight for
Hydrides
If London dispersion forces are dominant, The boiling
points of hydrides should increase with increasing
molar mass but H2O, NH3 and HF do not follow this
trend. Why?
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Hydrogen Bonding in water
Properties of water and IMF
The e- s in the O-H bond are
drawn to the O atom, leaving the
dense positive charge of the
hydrogen nucleus exposed.
The strong attraction of this
exposed nucleus for the lone pair
on an adjacent molecule accounts
for the strong hydrogen bond.
+
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H-Bonding in Water and properties
On freezing the molecules are arranged in open
hexagonal pattern that gives ice lower density at 0°C
(0.917g.cm-3) than water(1.000 g. cm-3).
Moderate temperatures near lakes
Water expands on freezing (equatic life under ice)
H-bonding in Biological molecules
A similar mechanism explains the attractions in HF
and NH3
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Identifying intermolecular forces
1. What kind of intermolecular forces are expected in
the following substances? Arrange them in order of
increasing boiling points.
CH4 , CHCl3 , C2H6 and CH2CH2OH.
2. Choose the substance in each pair that is a liquid at
room temperature (the other is a gas).
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Ion–Dipole Attraction
Ion–Dipole forces-occurs when an ionic compound
is mixed with a polar compound
ions from an ionic compound are attracted to the
dipole of polar molecules.
E.g. Solubility of ionic compounds in water.
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Phase transition revisted
Phase Transition:A change of a substance from
one state to another is called a change of state or
phase transition
Heat for phase transition q= n H
Hsub = Hfus + Hvap
Hsub = Hfus + Hvap
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11.5 Vapourization and Vapour
Pressure
The average kinetic energy is proportional to the
temperature.
Vapourization- Some molecules with more kinetic
energy than the average, overcome the attractive
forces and escape the liquid to gas phase as vapour
Rate of vapourization depends;
Temperature
Surface area
Strength of IMF
Condensation: Opposite of vapourization
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Distribution of Thermal Energy
The higher the temperature, the greater the average
energy of the collection of the molecules
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Effect of IMF on Vapourization
Volatile substances -Liquids and solids that
evaporate easily at normal temperatures
Relatively high vapour pressure
Weak IMF e.g., gasoline, fingernail polish remover
Nonvolatile - Liquids that that do not vapourize
easily
Low vapour pressure
Strong IMF e.g. H2O compared to acetone or
motor oil
Vaporization is an endothermic process. Why?
Condensation is an exothermic process. Why?
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Heat of Vaporization
Heat of vaporization, H vap-The amount of heat
energy required to vaporize one mole of the liquid
sometimes called the enthalpy of vaporization
always endothermic, H vap is +ve
somewhat temperature dependent e.g. for water
H2O(l)
H2O(g) H vap = 44.7 kJ mol-1 (25 °C)
= 40.7 kJ mol-1 (100 °C)
H cond. = - H vap
= - 40.7 kJ mol-1 (100°C)
Heats of vapourization of several liquids (Table 11.7)
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Heat of Vapourization
Using heat of vapourization in calculation (recall
stoichiometry of H section 6.5)
Calculate the mass of water that can be vaporized
with 155 kJ of heat at 100 °C.
Solution-
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Dynamic Equilibrium
Dynamic Equilibrium-The condition at which two
opposite processes, evaporation and condensation
occur at same rate
rate of vapourization = rate of condensation
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Vapour Pressure
Vapor pressure- The pressure exerted by the vapor
when it is in dynamic equilibrium with its liquid
The weaker the attractive forces between the
molecules, the higher the vapor pressure.
The higher the vapor pressure, the more volatile is
the liquid
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Vapor–Liquid Dynamic Equilibrium
Le Chatelier’s Principle-When a system in dynamic
equilibrium is disturbed the system responds so as to
minimize the disturbance and return to a state of
equilibrium
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Boiling Point
Boiling Point- The temperature at which the vapour pressure
of a liquid equals the external pressure ( by the atmosphere)
Boiling point varies with external pressure e.g.
bp of H2O =100 C
at 1.00 atm
= 95 C
at 0.83 atm
Normal boiling point -The temperature at which vapour
pressure of a liquid equals 1.00 atm (760 mmHg = 101.3kPa)
Normal bp of H2O at 1.00 atm = 100 C
The normal bp is related to vapor pressure and is lowest for
liquids with the weakest intermolecular forces.
When intermolecular forces are weak, little energy is
required to overcome them resulting in low boiling points
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Vapour Pressure of Several liquids at
different temperature
Which liquid has the strongest intermolecular
attractions?
• Vapour pressure increases
with increasing temperature
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Heating Curve of a Liquid
As you heat a liquid, its temperature increases
linearly until it reaches the boiling point.
q = mass × Cs × T
During boiling the
temperature remains
constant
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Critical Point: Another phase of
matter
•Supercritical fluids- have properties of both gas and
liquid states
•Critical point-at which gas and liquid phases mingles; the
meniscus between gas and liquid disappears
Critical Temperature(Tc)-The temperature at which this
transition occurs
Critical Pressure(Pc)- The pressure at which this transition
occurs
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11.6 Sublimation and fusion
Sublimation- The phase transition from solid to gas
Deposition- The phase transition from gas to solid
The opposite of sublimation.
The solid and vapor phases exist in dynamic
equilibrium in a closed container.
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Melting or Fusion
Melting point (fusion)- The temperature at which a
crystalline solid changes to a liquid (endothermic)
E.g. H2O(s)
H2O(l)
Hfus= 6.02kJ /mol
Freezing- Freezing Point-The temperature at which pure
liquid changes to a crystalline solid or freezes (Exothermic)
opposite to melting
Hcryst = - Hfus
H2O(l)
H2O(s)
Hcryst= - 6.02kJ /mol
can be used to identify the substance
Heat of fusion, H fusThe amount of heat energy required to
melt one mole of the solid
Hfus is always +ve
somewhat temperature dependent
Hsub = Hfus + Hvap
(Heat of fusion for several substances See Table11.8)
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Heating curve for one mole of Water
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Crystalline solids: Fundamental
Types
There are four types of solids.
Molecular solids (Van der Waals forces) composite particles
are molecules.
E.g. solid water (ice), and solid carbon dioxide (dry ice)
Ionic solids(Ionic bond)- composite particles are ions
Atomic solids- composite particles are atoms
Metallic (Metallic bond) e- - sea model
Nonbonding atomic solids are held together by
dispersion forces.
Network covalent atomic solids –are held together by
covalent bonds
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Classification of Solids
A molecular solid- composed of molecules held together by
intermolecular forces
dispersion forces, dipole–dipole attractions, and H-bonds
Many solids are of this type e.g. solid neon, ice, solid sulfur
(S8) and solid carbon dioxide (dry ice)
Weak IMF leads to low melting points. generally <300 °C
An ionic solid consists of cations and anions held together by
electrical attraction of opposite charges (ionic bond) e.g. NaCl
Strength of Ionic bonds depends on
The magnitude of ionic charge
Size of the ions
the m.p. of NaCl 801°C compared to MgO is 2800 °C
Lattice energy(energy needed to separate a crystal into isolated
ions in gaseous form)
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Atomic Solids
Nonbonding Atomic Solids- A group that only consists of
noble gases in solid form e.g. Xe
held together by weak dispersion forces
very low melting points which increases uniformly with
molar mass
e.g. mp of Ar is -189°C and of Xe is -112°C
tend to arrange atoms in closest-packed structure
maximizes attractive forces and minimizes energy
A metallic solid consists of positive cores of atoms held
together by a surrounding “sea” of electrons (metallic bonding).
positively charged atomic cores are surrounded by
delocalized electrons e.g. Fe, Cu, Au etc
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Network Covalent Solids
A covalent network solid consists of atoms held
together in large networks or chains by covalent
bonds.
Examples include different forms of carbon, as
diamond (3-D) or graphite (2-D sheets) asbestos
(long chains), and silicon carbide.
very high melting points
generally >1000 °C
The dimensionality of the network affects other
physical properties.
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The crystalline Structure of Diamond
and Graphite
Covalent bonds extend throughout a crystalline solid
3-D Network
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Two-dimensional
structure 43
The Structure and Properties of
Diamond
Diamond(bonding) -Each C is sp3 hybridized and
covalently bonded to four other C atoms
tetrahedrally (very stable 3–D structure)
each crystal is one giant molecule held together
by covalent bonds
Propertiesvery high mp ~3800 °C
very rigid (directionality of the covalent bonds)
Very hard (used as abrasives)
electrical insulator & nonreactive
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The Structure and properties of
Graphite
Graphite(bonding)- Each C is sp2 hybridized and bonded to
three other C atoms in trigonal planar geometry (TwoDimensional Network)
– has three bonds and one bond.
– form flat sheets fused together
each sheet is a giant molecule
The sheets are then stacked and held together by dispersion
forces.
Properties- slippery feel, used as lubricants
high mp
electrical conductor due to delocalized pi electrons
chemically nonreactive
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Silicates
Most common network covalent atomic solids
~90% of Earth’s crust
Basic compound silica (SiO2)- extended arrays of Si O
Al substituted for Si— aluminosilicates
Quartz - most common crystalline form SiO4
three-dimensional array of Si covalently bonded to four O in
tetrahedral geometry
impurities add color
mp ~1600 °C.
very hard
Glass is SiO2the amorphous form (Fig 11.49b)
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Operational Skills
Calculating the heat required for a phase change of a
given mass of substance.
Identifying intermolecular forces and describing
physical properties in terms of IMF’s
Identifying types of solids.
Determining the relative melting points based on
types of solids.
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