Transcript Intermolecular Forces

Intermolecular
Forces
(rev. 12/15/09)
Objectives
SWBAT:
 Distinguish between different types of
intermolecular forces.
 Complete a heating or cooling curve
calculation.

Intermolecular Forces

Forces that hold solids and liquids
together may be ionic or covalent
bonds or they may involve a weaker
interaction called intermolecular
forces.

All of these forces are van der Waals
forces
Intermolecular Forces

Generally,
the strengths of intermolecular forces
are much weaker than
intramolecular forces
(ionic or covalent bonds).

The stronger the attractive force, the
higher the boiling or melting points.

The Intermolecular Forces
(forces between molecules)
are weaker than
Intramolecular Forces
(The Chemical Bonds within an
Individual Molecule).
Types of Intermolecular Attractive Forces

Ion – Dipole Forces

Dipole – Dipole Forces

Hydrogen Bonding

London Dispersion Forces

Dipoles arise from opposite but equal
charges separated by a distance.
Molecules that possess a dipole
moment are called Polar molecules
(remember the polar covalent bond?).
Ion-Dipole Forces


Ion-dipole forces –
exist between an
ion and the partial
charge on the end
of a polar molecule
http://www.chem.purdue.edu/gchelp/liquids/ions.gif
Electrolytes

When salt is dissolved in water,
the ions of the salt dissociate from
each other and associate with the
dipole of the water molecules. This
results in a solution called an
Electrolyte.
Dipole – Dipole Forces

Dipole-dipole forces – exist between
neutral polar molecules, when dipoles are
close together
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these are weaker than ion-dipole forces
The molecules orient themselves to
maximize the positive/negative interactions
and to minimize the + + and - - interactions.
These forces are typically only about 1% as
strong as covalent or ionic bonds.
These forces rapidly become weaker as the
distance between the dipoles increases.
Dipole-Dipole

http://upload.wikimedia.org/wikipedia/commons/5/59/Dipole-dipole-interaction-in-HCl-2D.png
http://itl.chem.ufl.edu/2041_f97/lectures/lec_g.html


Inductive forces arise from the distortion of the
charge cloud induced by the presence of another
molecule nearby. The distortion arises from the
electric field produced by the charge distribution of
the nearby molecule.
These forces are always attractive but are in
general shorter ranged than electrostatic forces. If a
charged molecule (ion) induces a dipole moment in
a nearby neutral molecule, the two molecules will
stick together, even though the neutral molecule
was initially round and uncharged.
London Dispersion Forces

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London Dispersion forces:
exist primarily between
non-polar atoms or molecules, (including noble gases)
Sometimes called induced dipole-induced dipole attraction.
These forces exist between all molecules to some degree.
http://itl.chem.ufl.edu/2041_f97/lectures/lec_g.html

Inductive forces that result not from permanent charge
distributions but from fluctuations of charge are not
called inductive forces at all but are called London
Dispersion forces.

These forces are everywhere but are most important in
systems that have no other types of molecular
stickiness, like the rare gases (rare gases include the
noble gases, xenon, krypton and neon).
The rare gases may be liquified, and it is dispersion
forces that hold the atoms together (no electrostatic or
inductive forces exits)

London Dispersion Forces


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The constant motion of an electron in an
atom or molecule can create an
instantaneous dipole moment by affecting
the electron distribution of a neighboring
atom
This inter-atomic attraction is relatively weak
and short lived. This is the weakest
intermolecular force.
The strength of these forces increases with
increasing molecular mass
London Dispersion Forces

London forces are the attractive forces that
cause non-polar substances to condense to
liquids and to freeze into solids when the
temperature is lowered sufficiently.

Dispersion forces are present between any
two molecules (even polar molecules) when
they are almost touching (this means they
are found in all substances).
London Dispersion Forces
http://itl.chem.ufl.edu/2045/matter/FG11_005.GIF
London Dispersion Forces
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Dispersion forces are present between all molecules,
whether they are polar or nonpolar.
Larger and heavier atoms and molecules exhibit
stronger dispersion forces than smaller and lighter
ones (outer electrons are shielded from nucleus
positive charge allowing more interactions).
In a larger atom or molecule, the valence electrons
are, on average, farther from the nuclei than in a
smaller atom or molecule. They are less tightly held
and can more easily form temporary dipoles.
The ease with which the electron distribution around
an atom or molecule can be distorted is called the
polarizability.
London Dispersion Forces
London dispersion forces tend to be:
 stronger between molecules that are
easily polarized.
 weaker between molecules that are
not easily polarized.

Hydrogen Bonding

Hydrogen bonding – is a special type of
intermolecular attraction that exists between
the hydrogen atom in a polar bond
(particularly an H-F, H-O or H-N bond) and
an unshared electron pair on a nearby small
electronegative ion or atom (usually an F, O,
or N atom on another molecule).

This is a specific type of dipole-dipole force
Hydrogen Bonding
https://vinstan.wikispaces.com/file/view/800px-Hydrogen-bonding-in-water-2D.png/46631659/800px-Hydrogen-bonding-inwater-2D.png
Hydrogen Bonding

Two factors account for the strengths of these
interactions:
1. large polarity of the bond
2. close approach of the dipoles (allowed
by the very small size of the hydrogen
atom)
Hydrogen Bonding

Each attraction is electrostatic in nature,
(involving attractions between positive and
negative species)

See Brown and LeMay page 403 for a flow
diagram for intermolecular forces.
Polarizability

Polarizability – the ease with which the charge
distribution in a molecule can be distorted by an
external electric field. (see B&L pg. 397)

More polarizable molecules have stronger London
Dispersion forces

Strength increases with increasing size
occurs between all polar and non-polar molecules
Properties of Liquids
viscosity – the resistance of a liquid to flow
 The greater a liquid’s viscosity, the more
slowly it flows.
 Viscosity decreases with increasing
temperature. At higher temperatures, the
greater average kinetic energy of the
molecules more easily overcomes the
attractive forces between molecules.
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Surface Tension
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Surface tension – the energy required to increase
the surface area of a liquid by a unit amount.
Surface tension is due to an increase in the attractive
forces between molecules at the surface of a liquid
compared to the forces between molecules in the
center, or bulk, of the liquid. This property causes
fluids to minimize their surface areas.
see Brown and LeMay page 404
Surface Tension

When a liquid is poured onto a solid surface,
it tends to bead as droplets, which is a phenomenon
that depends on the intermolecular forces.
http://quest.nasa.gov/space/teachers/microgravity/image/66.gif
http://z.about.com/d/physics/1/G/8/0/-/-/SurfaceTension.png
Surface Tension

Although molecules in the interior of the
liquid are completely surrounded by other
molecules, those at the surface are subject
to attractions only from the side and from
below. The effect of this uneven pull on the
surface molecules tends to draw them into
the body of the liquid and causes a droplet
of liquid to assume the shape that has a
minimum surface area (a sphere).
Phase Changes Section
Vocab

The melting process for a solid can be referred to as fusion.

A heating curve is a plot of the temperature versus the amount
of heat added.

A cooling curve is a plot of the temperature versus the amount
of heat removed.
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Critical temperature is the highest temperature at which a
substance can exist as a liquid.
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The critical pressure is the pressure required to bring about
liquefaction at this critical temperature.
Heating Curve of Water
Heat of
fusion
Heat of
vaporization
(1) is ice
(2) is ice and liquid water (melting)
(3) is liquid water
(4) is liquid water and vapor (vaporization)
(5) is water vapor
http://www.greatneck.k12.ny.us/GNPS/SHS/dept/science/Blumberg/worksheets/heating%20curve%20and%20energy_files/
image003.jpg
Heat of Fusion

Heat of Fusion (ΔHvap) is the energy
required to melt one mole of a
substance at constant temperature.
Heat of Vaporization

Heat of vaporization (ΔHvap) is the
energy required to vaporize one mole
of a substance at constant
temperature.
Lines on the Graph
The horizontal lines of a heating curve
represent the heat of fusion and heat
of vaporization.
 Notice that the temperature doesn’t
change during melting or vaporization.
 The nearly vertical lines represent the
heat required to effect the
corresponding temperature change of
a single phase.

Heating Curve Diagram
http://library.thinkquest.org/C006669/media/Chem/img/Graphs/HeatCool.gif
Heating Curve for Water

http://www.bbc.co.uk/schools/ks3bitesize/science/chemistry/physical_changes_4.shtml
Cooling Curve
www.docbrown.info
Heating and Cooling Curves
We need to look in the textbook to see some
heating and cooling curves and how to do the
calculations.
 See B&L page 406
 Try Sample Exercise 11.4 and the Practice
Exercise

Students

See teacher’s webpage for several
heating/cooling curve links

Try
http://chapsipc.wetpaint.com/page/Calculating+Heating+Curve+of+
for an example heating curve
calculation
Water?t=anon
Phase Diagrams
A phase diagram is a graphical way to
summarize the conditions under which
equilibria exist between the different
states of matter.
 The diagram also enables us to
predict the phase of a substance that
is stable at any given temperature and
pressure.
 See the diagrams B&L page 413

Phase Diagram for Water
www.serc.carlaton.edu
The point where the three lines intersect in a phase diagram
shows the pressure and temperature where the solid, liquid,
and vapor all exist in equlibrium. This point, which occurs for
water at 0.01°C (32.02°F), is known as the triple point.
http://encarta.msn.com/media_461541579/phase_diagram_for_water.html
http://www.chem.queensu.ca/people/faculty/Mombourquette/FirstYrChem/colligative/index.htm
On a Phase Diagram
You should be able to label:
Each phase change
(i.e. sublimation, melting, freezing,
etc.)
 Triple point and critical point.
 Direction of curves for H2O and CO2
diagrams.

Triple Point
The “triple point” is where all three
curves intersect on a phase diagram.
 All three phases co-exist at this point.

What is the definition of the
term “critical point” on a phase
diagram?

Phase Diagrams
Each diagram contains 3 curves.
 Each curve represents conditions of
temperature and pressure at which
the various phases can coexist at
equilibrium.

General Phase Diagram
http://images.google.com/imgres?imgurl=http://kramerslab.tn.tudelft.nl/~rob/Courses/PhysicsOfFluids/Figures%2Bmovies/PhaseDiagra
m.jpg&imgrefurl=http://kramerslab.tn.tudelft.nl/~rob/Courses/PhysicsOfFluids/html-lectures/Lecture1.1.html&usg=__kgqG_EuLDsbyFSuKYNB_JMHTQ=&h=315&w=412&sz=19&hl=en&start=5&tbnid=vuK0Mfhoi9N4aM:&tbnh=96&tbnw=125&prev=/images%3Fq%3Dphase%2Bdiagr
am%26gbv%3D2%26hl%3Den
Phase Diagram
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On the previous slide see the liquid/gas curve.
This is the vapor pressure curve.
The point on the graph where the vapor pressure is
1 atm is the normal boiling point of the substance.
The vapor pressure curve ends at the critical point.
Beyond the critical point the liquid and gas phases
becomes indistinguishable.
Phase Diagram
for Carbon Dioxide
http://serc.carleton.edu/images/research_education/equilibria/h2o_phase_diagram__color.v2.jpg
Notice the solid/liquid curve on the
carbon dioxide phase diagram.
 This curve follows the typical
behavior, the melting point increases
with increasing pressure.

Phase Diagram for Water
Water Phase Diagram
Notice the solid/liquid curve on the
water phase diagram.
 The melting point of water decreases
with increasing pressure.
 Water is one of a few substances
whose liquid form is more compact
than its solid form.
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Why does ice float?
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Because ice floats, we can infer that ice must
be less dense than water.
If water is frozen in a glass jar, the glass jar
breaks.
If a soda can freezes, it will also burst.
From both of the above we infer that the
volume of the ice has increased.
Conclusion: The volume of ice must be
greater than the same mass of liquid water.
Why does the volume increase?
Molecular basis for the Volume
Increase of Ice:
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The normal pattern for most compounds is that as the
temperature of the liquid increases,
the density decreases as the molecules spread out
from each other. As the temperature
decreases, the density increases as the molecules
become more closely packed.
This pattern does not hold true for ice as the exact
opposite occurs.
In liquid water each molecule is hydrogen bonded to
approximately 3.4 other water
molecules. In ice each molecule is hydrogen bonded to
4 other molecules.
Compare the structures of Liquid
Water and Solid Ice – See Graphic
 Notice the empty spaces within the ice
structure, as this translates to a more
open or expanded structure.
 The ice structure takes up more
volume than the liquid water
molecules, hence ice is
 less dense than liquid water.
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Ice vs Water Structure
http://www.elmhurst.edu/~chm/vchembook/images/122iceliquid.gif
Question

Explain why the phase diagram for
water is different than the phase
diagram for carbon dioxide.
Ch 11 Problems

5, 7-11, 13, 19, 25, 27, 33, 34, 37, 40, 47, 48,
52-54, 56, 57, 62, 65
AP Problems
 2003 #6
 2005 #8
 2006 #6
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Vapor Pressure

Vapor Pressure is the partial pressure
exerted by a vapor in a closed system
when it is in equilibrium with its liquid
or solid phase.