Intermolecular Forces

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Intermolecular Forces

• (rev. 12/15/09)

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Objectives

• SWBAT
• Distinguish between different types of
  intermolecular forces.
• Complete a heating or cooling curve calculation.

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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

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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.

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• The Intermolecular Forces
• (forces between molecules)
• are weaker than
• Intramolecular Forces (The Chemical Bonds
  within an Individual Molecule).

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Types of Intermolecular Attractive Forces

• Ion Dipole Forces
• Dipole Dipole Forces
• Hydrogen Bonding
• London Dispersion Forces

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• 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?).

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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
  

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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.

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Dipole Dipole Forces

• Dipole-dipole forces exist between neutral
  polar molecules, when dipoles are close together
• 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.

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Dipole-Dipole

• http//upload.wikimedia.org/wikipedia/commons/5/59
  /Dipole-dipole-interaction-in-HCl-2D.png

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http//itl.chem.ufl.edu/2041_f97/lectures/lec_g.ht
ml

• 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.

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London Dispersion Forces

• 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.

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http//itl.chem.ufl.edu/2041_f97/lectures/lec_g.ht
ml

• 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)

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London Dispersion Forces

• 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

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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).

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London Dispersion Forces
http//itl.chem.ufl.edu/2045/matter/FG11_005.GIF
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London Dispersion Forces

• 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.

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London Dispersion Forces

• London dispersion forces tend to be
• stronger between molecules that are easily
  polarized.
• weaker between molecules that are not easily
  polarized.

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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

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Hydrogen Bonding
https//vinstan.wikispaces.com/file/view/800px-Hyd
rogen-bonding-in-water-2D.png/46631659/800px-Hydro
gen-bonding-in-water-2D.png
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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)

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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.

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Polarizability

• Polarizability the ease with which the charge
  distribution in a molecule can be distorted by an
  external electric field. (see BL pg. 397)
• More polarizable molecules have stronger London
  Dispersion forces
• Strength increases with increasing size
• occurs between all polar and non-polar molecules

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Properties of Liquids

• viscosity the resistance of a liquid to flow
• The greater a liquids 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

• 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

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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
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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).

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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.
• Critical temperature is the highest temperature
  at which a substance can exist as a liquid.
• The critical pressure is the pressure required to
  bring about liquefaction at this critical
  temperature.

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Heating Curve of Water
Heat of vaporization
Heat of fusion
(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/scien
ce/Blumberg/worksheets/heating20curve20and20ene
rgy_files/image003.jpg
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Heat of Fusion

• Heat of Fusion (?Hvap) is the energy required to
  melt one mole of a substance at constant
  temperature.

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Heat of Vaporization

• Heat of vaporization (?Hvap) is the energy
  required to vaporize one mole of a substance at
  constant temperature.

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Lines on the Graph

• The horizontal lines of a heating curve represent
  the heat of fusion and heat of vaporization.
• Notice that the temperature doesnt change during
  melting or vaporization.
• The nearly vertical lines represent the heat
  required to effect the corresponding temperature
  change of a single phase.

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Heating Curve Diagramhttp//library.thinkquest.or
g/C006669/media/Chem/img/Graphs/HeatCool.gif
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Heating Curve for Water
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Cooling Curvewww.docbrown.info
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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 BL page 406
• Try Sample Exercise 11.4 and the Practice Exercise

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Students

• See teachers webpage for several heating/cooling
  curve links
• Try http//chapsipc.wetpaint.com/page/Calculating
  HeatingCurveofWater?tanon for an example
  heating curve calculation

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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 BL page 413

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Phase Diagram for Waterwww.serc.carlaton.edu
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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
0.01C (32.02F), is known as the triple point.
http//encarta.msn.com/media_461541579/phase_dia
gram_for_water.html
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http//www.chem.queensu.ca/people/faculty/Mombourq
uette/FirstYrChem/colligative/index.htm
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On a Phase DiagramYou 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.

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Triple Point

• The triple point is where all three curves
  intersect on a phase diagram.
• All three phases co-exist at this point.

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• What is the definition of the
• term critical point on a phase diagram?

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Phase Diagrams

• Each diagram contains 3 curves.
• Each curve represents conditions of temperature
  and pressure at which the various phases can
  coexist at equilibrium.

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General Phase Diagramhttp//images.google.com/img
res?imgurlhttp//kramerslab.tn.tudelft.nl/rob/Co
urses/PhysicsOfFluids/Figures2Bmovies/PhaseDiagra
m.jpgimgrefurlhttp//kramerslab.tn.tudelft.nl/r
ob/Courses/PhysicsOfFluids/html-lectures/Lecture1.
1.htmlusg__kgqG_EuLDs-byFSuKYNB_JM-HTQh315w
412sz19hlenstart5tbnidvuK0Mfhoi9N4aMtbnh
96tbnw125prev/images3Fq3Dphase2Bdiagram26
gbv3D226hl3Den
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Phase Diagram

• 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.

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Phase Diagram for Carbon Dioxidehttp//serc.carl
eton.edu/images/research_education/equilibria/h2o_
phase_diagram_-_color.v2.jpg
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• Notice the solid/liquid curve on the carbon
  dioxide phase diagram.
• This curve follows the typical behavior, the
  melting point increases with increasing pressure.
  

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Phase Diagram for Water
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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?

• 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?

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Molecular basis for the Volume Increase of Ice

• 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.

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• 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/122i
celiquid.gif
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Question

• Explain why the phase diagram for water is
  different than the phase diagram for carbon
  dioxide.

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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.