Transcript GASES - University of the Witwatersrand

INTERMOLECULAR FORCES
LIQUIDS AND SOLIDS
Chapter 11
STATES OF MATTER
Weak attractive forces
between molecules
Intermolecular
forces stronger
Strong
intermolecular
forces
The state of a substance depends largely on the balance between the
kinetic energies of the particles and the interparticle energies of
attraction.
INTERMOLECULAR FORCES
intramolecular force
- covalent bond between atoms
in a molecule
(can also be ionic bond)
weak
16 kJ/mol
Strong
431 kJ/mol
HCl boils at
-85oC at 1 atm
intermolecular force
- attraction between molecules
When a substance melts or boils the intermolecular forces are
broken (not the covalent bonds).
 Many properties are governed by the strength of the
intermolecular forces e.g. boiling point, melting point, vapor
pressure, viscosity, etc.
A liquid boils when:
A solid melts when:
4 types of intermolecular forces:
• Ion-Dipole Forces
• Dipole-Dipole Forces
• London Dispersion Forces
• Hydrogen Bonding
van der
Waals
forces
electrostatic forces
Can you arrange them in order of increasing strength?
What is a dipole?
+
-
Ion-Dipole Forces
• Interaction between an ion and a dipole
• Strongest of all intermolecular forces.
• Ion-dipole interactions make it possible for ionic substances to
dissolve in polar solvents.
Dipole-Dipole Forces
• Interaction between two dipoles
• Weaker than ion-dipole forces.
• Polar molecules need to be close together.
Molecules in liquids are free to move
 results in both attractive and
repulsive forces.
 overall and stronger attractive
force
If two molecules have about the same mass and size, then
intermolecular attractions (dipole-dipole forces) increase with
increasing polarity.
London Dispersion Forces
• Interaction between nonpolar atoms or molecules, but can also
occur between polar molecules
• Occur only when atoms or molecules are close together
• Weakest of all intermolecular forces.
While the electrons in the 1s
orbital of helium would repel each
other (and therefore tend to stay
far away from each other), it does
occasionally happen that they
wind up on the same side of the
atom.
Instantaneous
dipoles formed!
These dipoles are
temporary
Polarisability: the ease with which the distortion of the charge
distribution occurs.
Greater polarisablility  stronger dispersion
In general, larger molecules tend to have greater polarisability as
they have more electrons which are further away from the nuclei.
Exercise
Explain why n-pentane has a higher boiling point than neopentane?
Both have the molecular formula C5H12.
Hydrogen Bonding
Special case of dipole-dipole forces
Boiling pints
polar
nonpolar
Hydrogen bonds are dipole-dipole interactions between the H-atom
in a polar bond (usually H-F, H-O or H-N) and an unshaired e- pair
on a nearby small electronegative ion or atom (usually F, O, N)
Hydrogen
bond
Why are we specifying F, O and N?
Small and electronegative!
Hydrogen only has 1e-  +ve nucleus is rather exposed.
Hydrogen can approach the small electronegative atom closely and
interact strongly.
Exercise
Why does ice float on water?
Exercise
Which of the following molecules can hydrogen bond with itself?
O
CH2F2
NH3
CH3OH
H3C– C–CH3
SUMMARY
PROPERTIES OF LIQUIDS
Viscosity
The resistance of a liquid to flow
- related to the ease with which
molecules can move past each other
Viscosity increases with molecular weight and
decreases with higher temperature.
Which intermolecular forces also increase with molecular weight?
Surface tension
A measure of the inward forces that must
be overcome in order to expand the
surface area of a liquid
Surface molecules experience
a net inward force.
Molecules at the surface can
pack more closely together
Molecules in the interior are attracted
equally in all directions.
Surface tension of water at 20oC:
0.0729 J/m2
PHASE CHANGES
Every phase change is accompanied by a change in energy of the
system.
H (change in enthalpy)
Hvap
Hcond
Hsub
Hfus
(fusion)
Hfreez
Hdepos
Hvap
Hfus
Hvap > Hfus
 In the transition from liquid to vapour phase, the molecules must
essentially sever all intermolecular interactions.
 In melting, many of these interactions remain.
Hsub =
Plot of temperature versus heat added is a heating curve
e.g. Ice initially at -25oC is heated (constant P = 1 atm)
While evaporating,
heat added is used to
break intermolecular
forces
Recall: Hvap >
Hfusmelting, heat
While
added is used to break
intermolecular forces
From the graph determine:
Hfus
Hvap
HAB
HCD
HEF
?
n = (100 g)/(18.016 g/mol)
n = 5.55 mol
Exercise
Calculate the enthalpy change, ΔH, when 100 g of water at 50 oC is
cooled down to ice at -30 oC.
A
water
o
50 C
Given:
Specific heat capacities:
Water = 4.18 J/g K, Ice = 2.09 J/g K
Δ Hfus = 6.01 kJ/mol
HAB
H = CmT
T = Tf - Ti
C
0oC
-30oC
ice
HCD
B
HBC =
Hfreez
D
HAB = (4.18 J/g K)(100 g)(0 - 50 K) = -20.9 kJ
HBC = ΔHfreez = -ΔHfus = -(6.01 kJ/mol)(5.55 mol) = -33.4 kJ
HCD = (2.09 J/g K)(100 g)(-30 - 0 K) = -6.27 kJ
HAD = HAB + HBC + HCD = -60.6 kJ
What is the enthalpy change when 100 g of ice at -30 oC is heated up to
water at 50 oC?
VAPOUR PRESSURE
Closed
container
• Some of the molecules on the surface of a liquid have enough
energy to escape the attraction of the bulk liquid  these
molecules move into the gas phase.
• As the number of molecules in the gas phase increases, some of
the gas phase molecules strike the surface and return to the
liquid.
• After some time the pressure of the gas will be constant at the
vapour pressure  when liquid and vapour reach dynamic
equilibrium.
As the temperature increases, the fraction of molecules that have
enough energy to escape increases.
The boiling point of a liquid is the temperature at which its vapor
pressure equals atmospheric pressure.
Normal boiling point  boiling point at 1 atm.
Why does water boil at a higher temperature at the coast than here
in Jhb?
PHASE DIAGRAMS
Phase diagrams display the state of a substance under various
pressure and temperature conditions.
The solid lines show the conditions P,T conditions under which
equilibrium exists between phases.
Line AB: liquid-vapour interface
- Each point along this line is the boiling point of the substance at
that pressure.
Critical point B: above this critical temperature and critical
pressure the liquid and vapour are indistinguishable from each
other.
Line AD: solid -liquid interface
- Each point along this line is the melting point of the substance at
that pressure.
Line AC: solid-vapour interface
- Each point along this line is the sublimation point of the
substance at that pressure.
- Note: the substance cannot exist in the liquid state below A
Triple point A: the temperature and pressure condition at which all
three states are in equilibrium.
Phase Diagram of Water
Note the high critical
temperature and critical
pressure:
 due to the strong van der
Waals forces between water
molecules.
The slope of the solid–liquid line
is negative.
 The melting point decreases
with increasing pressure.
WHY?
Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in
the liquid state at pressures
below 5.11 atm.
CO2 sublimes at normal
pressures.
BONDING IN SOLIDS
Solids can be:
crystalline
Particles are in highly ordered
arrangement
or
amorphous
No particular order in the
arrangement of particles.
The physical properties of crystalline solids (e.g. m.p., hardness)
depend on the arrangement of particles and on attractive forces
between particles.
There are 4 types of crystalline solids:
•
•
•
•
Molecular
Covalent network
Ionic
Metallic
Molecular Solids
Atoms or molecules are held together by intermolecular forces
 dipole-dipole forces, London dispersion forces, hydrogen bonds
Because of these weak forces they are soft and have relatively low
melting points (<200oC)
Most substances that are gases or liquids at room temperature
form molecular solids at low temperature.
Explain the m.p.’s and b.p.’s observed below:
m.p./oC
b.p./oC
Benzene
5
80
Toluene
-95 CH
3
111
Phenol
43 OH
182
Covalent-Network Solids
Atoms are held together in large networks or chains by covalent
bonds.
Covalent bonds much stronger than intermolecular forces
 Harder solids and higher melting points than molecular solids.
Diamond
Graphite
Which one is harder and has the higher m.p.? Explain.
Ionic Solids
Ions are held together by ionic bonds
 strength of the ionic bond depends on the charges of the ions
Ions pack themselves so as to maximize the attractions and
minimize repulsions between the ions
 Depends on relative size and charge of ions
Metallic Solids
Consist of entirely metal atoms.
Metals are not covalently bonded, but the attractions between
atoms are too strong to be van der Waals forces.
Bonding due to valence
electrons delocalized
throughout the solid.
In general, the strength of
bonding increases as the
no. of electrons available
for bonding increases
The m.p. for sodium is 97.5oC and for chromium is 1890oC. Explain.
SUMMARY