June 12, 2009 – Class 39 and 40 Overview
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12.5 Van der Waals Forces
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Dispersion (London) forces, polarizability, dipole-dipole
interactions, permanent, instantaneous and induced dipoles,
molecular shape and polarizability, predicting relative boiling
points from relative strength of intermolecular forces.
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12.6 Hydrogen Bonding
–
Properties of hydrogen bonds (strong vs weak hydrogen
bonds), hydrogen bonding in water and living matter,
intermolecular and intramolecular hydrogen bonding, effect on
viscosity.
June 12, 2009 – Class 39 and 40 Overview
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12.1 Intermolecular Forces and Some Properties of Liquids
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12.2 Vaporization of Liquids: Vapour Pressure
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Vaporization(evaporation), enthalpy of vaporization, vapor
pressure curves, boiling and the boiling point, normal boiling
point, measuring vapor pressure data, use of the ClausiusClapeyron equation.
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12.3 Some Properties of Solids
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Melting, melting (freezing) point, heat of fusion, sublimation,
deposition.
Intermolecular Forces and Some Properties of Liquids
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Intramolecular Forces: attraction within a molecule. A covalent
bond is an example of an intramolecular force.
Intermolecular Forces: an attraction between molecules.
•
–
–
Responsible for condenses phases (solids and liquids).
Intermolecular forces explain many properties of liquids.
(a) Molecules at the
surface interact with
other surface
molecules and with
molecules directly
below the surface.
(b) Molecules in the
interior experience
intermolecular
interactions with
neighboring
molecules in all
directions.
Intermolecular forces
•
There are three primary types of intermolecular forces:
1.
(London) Dispersion Forces
2.
Dipole-Dipole Forces (van der Waal's Forces)
3.
Hydrogen Bonding
•
Note:
polarizability: the ease with which a particle’s electron
cloud can be distorted.
Dispersion (London) forces
•
Dispersion (London) forces: Intermolecular forces
associated with instantaneous and induced dipoles.
–
Electron location is expressed in terms of probabilities only
–
For non-polar molecules it is most probable that the electrons
will be evenly distributed in the atomic or molecular orbitals.
–
It is possible for the electrons in one molecule to flicker into an
arrangement that results in partial positive (d+) and partial
negative (d-) charges.
–
When this occurs the molecules acquire an instantaneous (it
lasts less than 10-16 seconds!) dipole.
–
A molecule that has acquired an instantaneous dipole can then
induce a dipole in another molecule.
Dispersion (London) forces
a) Normal Condition: A non-polar molecule has a symmetrical charge distribution
b) Instantaneous Condition: A displacement of the electronic charge produces an
instantaneous dipole with a charge separation represented as d+ and d-.
c) Induced Dipole: The instantaneous dipole on the left induces a charge
separation in the molecule on the right. The result is a dipole-dipole
interaction.
The two dipoles, in the two molecules, will attract each other, and the
result is that the potential energy of the two is lowered.
Dispersion (London) forces
•
Dispersion forces are present in both polar and non-polar
molecules, however, it is the main type of intermolecular
force between non-polar molecules.
•
There are two factors that effect the magnitude of dispersion
forces:
1. Number of electrons: Larger molecules with more
electrons more easily undergo vibrations that lead to uneven
distribution of charge.
Molecules with more electrons will have stronger
dispersion forces.
Dispersion (London) forces
•
There are two factors that effect the magnitude of dispersion
forces:
2. Surface area: Molecules with a larger surface area offer a
greater opportunity for a molecule to induce a dipole in a
nearby molecule.
Spherical molecules with the same number of
electrons as more branched molecules will have weaker
dispersion forces.
Dispersion (London) forces
•
Example: F2 has 18 electrons and Cl2 has 34 electrons. The
dispersion forces for Cl2 are stronger than those of F2.
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Example: Cl2 and C4H10 each have 34 electrons. C4H10 has
a larger, more complex shape therefore it has stronger
dispersion forces than Cl2.
•
When a substance melts or boils, the intermolecular forces
are overcome.
•
Molecules with greater dispersion forces will have higher
boiling points because more energy is required to overcome
the attraction between molecules.
Dipole-Dipole Forces
•
Dipole-Dipole Forces: intermolecular attractions associated
with molecules with permanent dipoles.
•
Recall that the bond dipoles for molecules do not always
cancel. When the bond dipoles do not cancel the resulting
molecule is polar (it has a permanent dipole).
•
This leads to polar molecules trying to line up with the
positive of one dipole directed toward the negative end of
neighbouring dipoles.
Dipole-Dipole Forces
Dipole-dipole forces involve
the displacement of
electrons in bonds, rather
than the displacement of all
the electrons in a molecule,
as in dispersion forces.
Dipole-dipole forces add to
dispersion forces (which are
present for all molecules).
Dipole-dipole forces also
affect physical properties
such as melting and boiling
point.
Dipole-Dipole Forces
Example: The boiling point of N2(l) is -195.81 oC. The
boiling point of O2(l) is -182.96 oC. If only dispersion
forces are considered, one would predict the boiling
point of NO(l) to be in between that of O2(l) and N2(l) .
The actual boiling point of NO(l) is -151.76 oC, much
higher than either that of O2(l) or N2(l) .
NO is a polar molecule, hence it has both dispersion
forces and dipole-dipole forces present. It has a higher
boiling point than O2(l) or N2(l) because extra energy is
required to overcome the dipole-dipole interactions.
Dispersion (London) forces
Problem: Which would you expect to have the higher
boiling point?
(a) C4H10 or (CH3)2CO
(b) C3H8, CO2, CH3CN
Hydrogen Bonding
•
Hydrogen Bonding: an intermolecular force of attraction
in which an H-atom covalently bonded to one highly
electronegative atom is simultaneously attracted to
another highly electronegative atom of the same or a
nearby molecule.
–
The only atoms that are electronegative enough to participate
in hydrogen bonding are fluorine, nitrogen and oxygen.
–
Hydrogen bonding, like dispersion forces and dipole-dipole
forces, also affects physical properties, like melting and boiling
points.
–
Notice the anomalies in the following graph of boiling points for
hydrides of Groups 4A, 5A, 6A and 7A.
Hydrogen Bonding
The values for NH3, H2O and HF are unusually high compared to those of other
member of their groups!
Hydrogen Bonding
•
Consider the molecule HF as our example.
•
Fluorine is highly electronegative, and pulls the bonding pair
of electrons closer to itself, leaving the hydrogen nucleus
unshielded with a partial positive charge.
•
The hydrogen nucleus is then attracted to the lone pair of
electrons on another highly electronegative atom.
•
Recall that the other highly electronegative atom will have a
partial negative charge, acquired by attracting electrons
involved in its molecular bonds.
Hydrogen Bonding
•
The result for gaseous HF is that it often forms cyclic (HF)6.
Covalent Bond
Hydrogen Bond
Hydrogen Bonding
N-H Bond Example:
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Urea is CO(NH2)2 and it undergoes extensive hydrogen
bonding due to its four hydrogen atoms bonded directly
to nitrogen, and also H-bonded to both nitrogen and
oxygen. Several of these interaction are depicted
below.
Hydrogen Bonding
O-H Bond Example:
•
In liquid H2O each hydrogen atom is bonded to at least
four other H2O molecules.
•
A few of these bonds are illustrated below.
The molecules can still move around
because they have enough kinetic
energy to break the hydrogen bonds,
which then reform with another H2O
molecule.
Hydrogen Bonding
•
Hydrogen bonding in H2O also explains why ice floats!
•
To float a solid must be less dense than the liquid.
•
In the liquid there is enough kinetic energy to overcome
some of the hydrogen bonds, but in the solid the kinetic
energy is no longer sufficient to overcome the hydrogen
bonds.
This results in the molecules being
organized into a crystalline
arrangement, which is less dense
than the arrangement of the H2O
molecules in the liquid.
Comparing Melting and Boling Points
1.
Classify your compound as ionic or molecular. Ionic
compounds have greater intramolecular forces than do
molecular compounds, hence they have higher melting
and boiling points.
2.
For your molecular compounds, list the intermolecular
forces present. The more intermolecular forces, the
higher the melting or boiling point for a compound.
Comparing Melting and Boling Points
Melting/Boiling Point
Compound
Highest
1. Ionic
2. Molecular
(H-Bonding + Dipole-Dipole
+ Dispersion Forces)
3. Molecular
(Dipole-Dipole + Dispersion
Forces)
Lowest
4. Molecular
(Dispersion Forces Only)
Intermolecular Forces and Some Properties of Liquids
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Surface Tension (g): the energy or work required to extend
the surface of a liquid, measured in J/m2.
–
Increasing temperature leads to more molecular motion, and
less work is required to extend a surface. Therefore, surface
tension decreases when temperature increases.
Despite being more dense than water, a
needle is supported on the surface due to
surface tension.
At 20 oC the surface tension of water is 7.28
x 10-2 J/m2.
Intermolecular Forces and Some Properties of Liquids
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Problem: What intermolecular forces account for
the following surface tension data?
substance (formula)
surface tension / J m-2 at 20oC
(a) pentane (C5H12)
1.6 x 10-2
(b) diethyl ether (C2H5OC2H5) 1.7 x 10-2
(c) ethanol (C2H5OH)
2.3 x 10-2
(d) butanol (C4H9OH)
2.5 x 10-2
(e) water (H2O)
7.3 x 10-2
(f) mercury (Hg)
48 x 10-2
Intermolecular Forces and Some Properties of Liquids
•
•
Cohesive Forces: intermolecular forces between like molecules,
such as within a drop of liquid.
Adhesive Forces: intermolecular forces between unlike molecules,
such as molecules of a liquid and of a surface with which it is in
contact.
(a) Water spreads into a thin film on
clean glass; adhesive forces are
greater than cohesive, and work done
by a collapsing drop supplies energy
for spreading.
(b) Glass coated in grease or oil; drops
of water stand on glass. Here,
cohesive forces are greater than
adhesive.
Intermolecular Forces and Some Properties of Liquids
Adhesive forces lead to a
concave meniscus for water
(g = 7.28 x 10-2 J/m2).
Cohesive Forces in mercury
consist of strong metallic
bonds (g = 47.2 x 10-2 J/m2) .
It forms a convex menisucus.
Adhesive forces (as
evidenced by the level of the
meniscus) are magnified in
tubes of small diameter,
known as capillary tubes.
Intermolecular Forces and Some Properties of Liquids
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Viscosity: refers to a liquid’s resistance to flow. Its magnitude
depends on intermolecular forces of attraction and in some cases,
on molecular sizes and shapes.
–
When liquid flows, one portion of the liquid moves with respect to
neighboring portions. Cohesive forces are responsible for the fast (ex.
water or ethanol) or slow (ex. honey or motor oil) rate of flow.
Stronger intermolecular forces lead to slower
flow and greater viscosity for a substance.
Viscosity can be determined by measuring the
velocity of a ball dropping through a liquid.
Vaporization of Liquids: Vapour Pressure
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Vaporization (evaporation): the passage of molecules
from the liquid to the gaseous state. Intermolecular
forces must be overcome in order for a sample to
vaporize. Vaporization occurs more readily with:
–
Increased temperature: more molecules have sufficient kinetic
energy to overcome intermolecular forces.
–
Increased surface area: a greater proportion of the molecules
are at the surface of the liquid.
–
Decreased strength of intermolecular forces: less kinetic
energy is required to overcome the forces of attraction.
Vaporization of Liquids
Recall from Chapter 7
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Enthalpy of vaporization (DHvap): quantity of heat that
must be absorbed in order to vaporize a liquid at
constant temperature.
–
Enthalpy of condensation (DHcondensation) is equal in magnitude
but opposite in sign to DHvap.
Vaporization of Liquids
Problem: How much energy is released when 0.459 g of
H2O (g) is cooled from 136 oC to 85 oC?
Vaporization of Liquids: Vapour Pressure
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Vapor pressure: the pressure exerted by a vapor when
it is in dynamic equilibrium with its liquid at a fixed
temperature.
vaporization
liquid
vapor
condensation
Liquid vaporizes in
a closed container.
Condensation begins.
Rate of vaporization >
rate of condensation.
Equilibrium. Rate of
vaporization = rate
of condensation.
Vaporization of Liquids: Vapour Pressure
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Liquids with high vapor pressure at room temperature
are volatile.
•
Liquids with high vapor pressure at room temperature
are nonvolatile.
•
Vapor pressure increases with temperature. This is
illustrated by vapor pressure curves.
•
Vapor pressure curves: a graph of vapor pressure as a
function of temperature.
Vapour Pressure Curves
(a) Diethyl ether, C4H10O
(b) Benzene, C6H6
(c) Water, H2O
(d) Toluene, C7H8
(e) Aniline, C6H7N
The normal boiling points are
the temperatures at the
intersection of the dashed
line (P = 760 mmHg).
Vapour Pressure Curves
Vaporization of Liquids: Vapour Pressure
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Boiling: a process in which vaporization occurs through
a liquid. It occurs when the vapor pressure of a liquid is
equal to barometric pressure.
•
Normal boiling point: the temperature at which the
vapor pressure of a liquid is 1 atm. It is the
temperature at which the liquid boils in a container
open to the atmosphere at 1 atm.
•
Boiling points vary with barometric pressure.
–
At an altitude of 1609 m (Denver, Colorado), barometric
pressure is 630 mmHg and the boiling point of water is 95 oC.
It takes longer to cook food under these conditions!
Vaporization of Liquids: Vapour Pressure
Vapor pressure plotted as ln P vs. 1/T (P measured in mmHg and T in Kelvin) –
note that the relationships are linear.
Vaporization of Liquids: Vapour Pressure
•
Clausius-Clapeyron equation: can be used to predict
the temperature at a certain pressure, given the
temperature at another pressure, or vice versa.
Alternatively, if the corresponding temperature and
pressure is known at two points, the enthalpy of
vaporization can be determined.
P2 DH vap
ln

P1
R
 1
1
  
 T1 T2 
Where P1 and T1 are one pair of pressure and temperature conditions
P2 and T2 are the other pair of pressure and temperature conditions.
DHvap is the heat of vaporization
R is the gas constant in units of 8.314 J mol−1K−1
Vaporization of Liquids: Clausius-Clapeyron
Equation
Problem: The normal boiling point of isooctane is 99.2 oC,
and its enthalpy of vaporization is 35.76 kJ/mol C8H18.
Calculate the vapor pressure of isooctane at 25 oC.
Problem: The vapor pressure of methanol is 100 mmHg at
21.2 oC, and is 121 mmHg at 25 oC. What is the
enthalpy of vaporization for methanol?
Some Properties of Solids
•
Melting: the transition of a solid to a liquid that occurs at
the melting point. The melting point and freezing point
of a substance are identical.
–
Heat of fusion (DHfus)
Some Properties of Solids
•
Sublimation: the passage of molecules from the solid
to the gaseous state.
•
Deposition: the passage of molecules from the
gaseous to the solid state.
•
Enthalpy of sublimation: DH sub  DH fus  DH vap
Sublimation of solid iodine, and the deposition of the vapor
to solid on the cooler walls of the flask.