Intermolecular forces: Generalizing properties

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Intermolecular forces Generalizing properties

• Low boiling point particles are more likely to
  leave liquid solution
• Weaker IM forces lower boiling point
• Lower boiling point more vapor higher vapor
  pressure
• High boiling point slow evaporation
• If IM forces are the same, look at formula
  weight. Heavier molecules have higher boiling
  points.
• Strength of IM forces
• Hydrogen bondgtdipole-dipolegtLondon dispersion

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Intermolecular Forces
Interacting molecules or ions
No
Yes
Yes
No
No
Yes
No
Yes
London Forces only Ex. Ar(l), I2(s)
Dipole-Dipole Ex. H2S
Hydrogen Bonding Ex. NH3, H2O
Ion-dipole Forces Ex. KBr in H2O
Ionic bonding Ex. NaCl
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Waters Properties

• Hexagonal crystal shape
• Molecule is polar.
• Hydrogen bonding
• Ice floats.
• Expands during freezing until -4.0 º C.
• Solid form is less dense than liquid
• Surface tension
• Water beads on smooth surfaces.
• Insects walk on water surfaces.

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Surface tension

• Force that pulls adjacent parts of a liquid
  surface together.
• The higher the attractive forces between
  particles in the liquid, the higher the surface
  tension.
• Hydrogen bonds make water have higher surface
  tension than most liquids.

Soap
Water droplet
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Phases of matter Comparison
Property Solid Liquid Gas
Particles Closely packed High density (ButWater is different!) More densely packed than in gas Most compressible-least densely packed
Particle movement Vibrate weakly around fixed positions Lowest kinetic energy Can change positions with other particles Can change positions with other particles Highest kinetic energy
Intermolecular forces Most effective (strongest) Stronger than in gases Least effective (weakest)
Shape and volume Both definite Definite volume only No definite shape or volume
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Solids

• Crystalline solids
• Particles are arranged in an orderly, geometric,
  repeating pattern.
• Examples Emerald, diamond, calcite
• Amorphous solids (Without shape)
• Particles are arranged randomly.
• Examples Glass, plastic
• Network solids
• Covalent bonds, usually single element arranged
  in orderly pattern
• Examples Diamond, graphite

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Bonding in Solids

• Molecular solids
• Most are liquids or gases at room temp.
• Ex. H2O, Ar
• Covalent Network solids
• Covalent bonds are stronger than IM forces, so
  substances have relatively high melting points
  and are harder than molecular ones.
• Ex quartz, diamond, graphite, SiO2
• Ionic solids
• Ionic bonds are the strongest of all
• Strength of bond depends on charge Higher
  charges higher melting point.
• Crystal structures Examples
• Face-centered cubic, body-centered cubic,
  hexagonal close-packed structures.
• Metallic Solids (metallic bonds)

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The Crystal Lattice

• 3-dimensional pattern that repeats itself over
  and over again.
• Each ion is bonded with all oppositely charged
  ions that directly surround it.
• NaCl forms a cube shape, called a
  body-centered-cubic structure.
• There are 7 crystal shapes, determined by how the
  ions are arranged in the lattice.

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Crystal Growth

• Crystals grow by adding ions to all sides.
• They grow equally in all directions from the
  outside.
• Crystals form in 2 ways
• Solution containing a dissolved ionic compound
  evaporates.
• An ionic solid is heated until it melts, then
  liquid is cooled. (Igneous rocks)

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Energetics of Ionic Bond Formation

• Recall that heat of formation of NaCl was
  exothermic (?Hf -410.9kJ/mol)
• Separation of NaCl is endothermic
• (?H 788 kJ/mol)
• The energy required to separate 1 mol of ions in
  an ionic lattice into gaseous ions is called
  lattice energy, ?Hlattice .
• Lattice energy depends on the charge on the ions
  and the size of the ions.

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Lattice Energy (Ionic bonds)

• Lattice energy depends on the charge on the ions
  and the size of the ions.
• The stability of the compound comes from the
  attraction between ions of unlike charge.
• The specific relationship is given by Coulombs
  equation
• E kQ1Q2 Q is the charge on the
  particles, d is the distance
  d between their centers and k
  is a constant.
• As Q1 and Q2 increase, E increases and as d
  increases, E decreases.

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Lattice Energies for Some Ionic Compounds
Compound Lattice Energy (kJ/mol) Compound Lattice Energy (kJ/mol)
LiF 1030 MgCl2 2326
LiCl 834 SrCl2 2127
LiI 730
NaF 910 MgO 3795
NaCl 788 CaO 3414
NaBr 732 SrO 3217
NaI 682
KF 808 ScN 7547
KCl 701
KBr 671
CsCl 657
CsI 600
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CrystalsImages created by Daniel Mayer or
Wikimedia Commons and licensed under terms of the
GNU FDL.

• Can examine structure using X-ray diffraction
• Uses Braggs Law to determine distance between
  planes of atoms. Computer instrumentation is
  used to translate wave functions into
  photographic images.
• Braggs Law n? 2dsinT
• n is an integer (1), ? is wavelength, d is
  distance between atoms, T angle of incidence
• 3D link
• http//www.le.ac.uk/eg/spg3/atomic.html

cbc
cfc
hexagonal
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Allotropes (different forms of same element)

• Carbon (C)
• Diamond
• Graphite (pencil lead)
• Charcoal
• Sulfur (S)
• Rhombic (puckered ring) S8
• Phosphorous (P)
• White phosphorous, P4 is most reactive,
  tetrahedral
• Red phosphorous is more stable.

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1. http//www.green-planet-solar-energy.com/the-eleme
   nt-sulfur.html
2. http//www.avogadro.co.uk/structure/chemstruc/netw
   ork/g-molecular.htm
3. http//www.enmu.edu/services/museums/miles-mineral
   /images/diamond_large.jpg

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Silicon Doping(N-type is more conductive when
voltage is applied.)

• OOOO OOOO OOOO
• OOOO OB.OO OPOO
• Silicon (4 e-) P-type N-type
• semiconductor hole created extra e- in
  lattice
• p positive n negative
• To customize
• conductive properties, add a dopant such as B
  (p-type), As or P (n-type)

.. .. .. ..
.. .. .. ..
.. .. ..
.. .. .. ..
.. .. .. ..
.. .. .. ..
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Glass SiO2

• High melting point ( 1700C) but may vary
  depending on the particular structure . Very
  strong Si-O covalent bonds arranged in a
  continuous lattice have to be broken throughout
  the structure before melting occurs.
• Is also very hard due to the need to break the
  very strong covalent bonds.
• Doesn't conduct electricity because there aren't
  any delocalized electrons. All the electrons are
  held tightly between the atoms, and aren't free
  to move.
• Insoluble in water and organic solvents because
  there are no possible attractions which could
  occur between solvent molecules and the silicon
  or oxygen atoms which could overcome the covalent
  bonds in the giant structure.
• (Water is a simple covalent structure Each
  molecule is made of 3 atoms so its melting point
  is low compared to the giant lattice structure
  of solids like SiO2. )

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Allotrope Two or more forms of the same
element that have distinctly different physical
or chemical properties.

• Fullerenes include C60, buckminsterfullerene, a
  hollow sphere resembling a soccer ball.
• Graphite is a black solid that feels soft and
  greasy to the touch. Planar sheets of
  molecules can slip by one another easily. It is
  used as a lubricant and leaves black marks if
  rubbed on a lighter-colored surface. It conducts
  electricity. Selling price lt 0.01/gram.
• Diamond is one of the hardest substances known
  (Mohs hardness 10). Its hardness is due to
  rigid networks of tetrahedrons, carbon atoms
  covalently bound. It does not conduct
  electricity. Selling price 50.00 -
  20,000.00/gram.

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Material modification

• Pencil lead is softened by adding clay to
  graphite.
• Gold jewelry is strengthened by adding copper or
  other metal. 14 karat gold means that 14/24s of
  the material is Au. (The relative proportion of
  gold originated with a medieval coin called a
  mark a mark weighed 24 karats.)
• Ceramics Developed from conventional clay (Si,
  O, Al) and the addition of other minerals to
  improve strength, melting point and brittleness.
  Ceramics can often get much hotter before they
  melt than metals.
• Plastics Synthetic polymers primarily from
  carbon. Disadvantage is that most are made from
  nonrenewable petroleum resources.

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Changing states

• Equilibrium When there is no net change in a
  system.
• Dynamic equilibrium
• When a vapor is in equilibrium with its liquid as
  one molecule leaves the liquid to become a vapor,
  another molecule leaves the vapor to become a
  liquid. In other words, an equal number of
  molecules will be found moving in both
  directions.

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Boiling Point

• Vapor pressure Pressure exerted by a vapor
  Pressure of the liquid at given temperature
• Liquid boils when its vapor pressure equals
  pressure of the atmosphere.
• Boiling is the conversion of a liquid to vapor
  within the liquid as well as at its surface.
• Boiling point is the temperature at which the
  equilibrium vapor pressure of the liquid equals
  the atmospheric pressure.
• Volatile liquids are liquids that evaporate
  readily.

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Boiling Point, cont.

• High elevation Low atmospheric pressure
• Low atmospheric pressure lower boiling point
• High pressure in pressure cooker increased
  boiling point, faster cooking
• If pressure above liquid increases, the liquid
  temperature rises until it matches the new
  pressure and boils again.

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Separation by Distillation

• Distillation is the separation of liquid
  substances according to their different boiling
  points.
• As a liquid mixture is heated, the substance with
  the lower boiling point will vaporize first.
• Distillate Condensed liquid substance

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Kinetic Energy and Equilibrium Vapor Pressure

• In the beginning
• particles condensing to liquid phase
• particles evaporating to gas phase
• Increase temp Increase kinetic energy
• Now, more molecules have enough energy to leave
  the liquid.
• More vapor molecules higher vapor pressure
• Equilibrium will soon be established, but at a
  higher vapor pressure.

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Heat of Vaporization

• Amount of heat necessary to boil (or condense)
  1.00 mole of a substance at its boiling point
• 1) 1.00 mole of a substance2) There is no
  temperature change
• The molar heat of vaporization (?Hvap ) for water
  is 40.7 kJ/mol. It comes from a table.
• q ?Hvap (mass / molar mass) (q ?Hvapn)
• q is total amount of heat involved.

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Heat of Vaporization

• Vapor pressure increases nonlinearly for liquids.
  Mathematically, the relationship is
• ln(Pvap) -?Hvap (1/T) C
• R
• Where C constant characteristic of a given
  liquid.
• m slope -?Hvap and x 1/T and b intercept
  C
• R

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Heat of Fusion

• aka standard enthalpy change of fusion
• Amount of thermal energy absorbed or lost for 1
  gram of a substance to change states from a solid
  to a liquid or vice versa.
• Temperature at which it occurs is called the
  melting point.
• Temperature falls if thermal energy is removed
  from a liquid or solid
• At the transition point between solid and liquid
  (melting point), EXTRA energy is required to go
  from liquid to solid and increase order. For
  molecules to maintain the order of a solid, extra
  heat must be withdrawn.
• In the other direction, to create the disorder
  from the solid crystal to liquid, extra heat must
  be added.
• The molar heat of fusion for water is 6.02
  kJ/mol.
• q ?Hvap (mass / molar mass)

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Phase DiagramA phase diagram is a graph of
pressure vs. temperature that shows the
conditions under which phases of matter exist.
Critical temp (Tc) Above this, the substance
cannot exist in the liquid state.
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Phase Diagrams Density

• Negative liquid/solid slope shows density of
  solid is LESS than
• liquid (like H2O). See previous slide.
• Most substances will have a positive slope of
  this line since most solids are more dense than
  the liquid

http//wine1.sb.fsu.edu/chm1045/notes/Forces/Phase
/Forces06.htm
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Four major "points" on a phase diagram

• Triple point, TP - All three phases can exist in
  equilibrium at this temperature and pressure.
• (The solid-liquid line and the liquid-vapor line
  meet.)
• Normal boiling point, Tb - The temperature at
  which the vapor pressure of a liquid is equal to
  standard atmospheric pressure.
• (Standard atmospheric pressure line crosses the
  liquid-vapor line.)
• Normal melting point, Tm - The temperature at
  which the vapor pressure of the solid and the
  vapor pressure of the liquid are equal.
• (Standard atmospheric pressure line crosses
  the solid-liquid line.)
• Critical temperature, Tc - The temperature above
  which no amount of pressure will liquefy a vapor.
  
• (The liquid-vapor line becomes vertical.)