Intermolecular Forces, Liquids, and Solids

1
Intermolecular Forces, Liquids, and Solids

• Chapter 11 Brown LeMay

2
Temperature Review

• Measure of kinetic energy
• What can you say about the KE of salt particles,
  water molecules, and oxygen particles at room
  temperature?
• State determined by strength of forces that keep
  particles together

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Strength

• Compare energy needed for phase change vs.
  decomposition in HCl(l)
• Intermolecular (called weak) because they are
  weaker than ionic or covalent
• Boiling point reflects strength of bonds in
  liquid
• Melting point reflects strength of bonds in solids

4
Kinds of Intermolecular Forces

• Three major kinds dipole-dipole, London
  dispersion, and hydrogen bonding
• In solutions, ion-dipole
• All are electrostatic in nature
• Approximately 15 of covalent or ionic strength

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Ion - dipole

• When?
• Ionic solid polar liquid
• Increases with increasing charge of ion or
  polarity of solvent
• Determines solubility

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Dipole-Dipole forces

• Weaker than previous
• end of one attracts end of another
• If size is equal, more polar has stronger dipole
  attractions. (NH3 vs H2O)
• If polarity is the same but masses differ, than
  smallest is stronger. (Able to orient better)

7
London Dispersion Forces

• All molecules have this
• Only attraction in nonpolar molecules
• How can Iodine be a solid?
• Temporary lopsided charge builds up from random
  motion of electrons - 1930
• Increases with mass we say it has greater
  polarizability
• Straight molecule is more polarizable than a
  curled up molecule why?
• Halogen Family is a great essay

8
Hydrogen Bond

• Strongest of all weak forces
• Is caused when H is bonded to F, O, or N
• These are so electronegative that the H is a
  naked nucleus or bare proton
• Very attractive!
• Will bond to nearby electron pairs

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Importance of Hydrogen Bonding

• Biological systems DNA, proteins
• Water chemistry (MP, BP, specific heat, surface
  tension)
• Density of ice

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Density

• Most solids are more dense than liquid
• Water is less dense because of hydrogen bonding
• At 4C, water becomes less dense
• Important for life in winter
• Causes lake turnover
• Alum example

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Practice

• Look at Flow Chart

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Properties of Liquids

• Viscosity
• Slower than..
• Resistance of a liquid to flow
• Time it as it goes through a small tube with
  gravity acting upon it.
• Poise 1g/cm-s
• Trends same substance decreases with
  increasing temperature
• series (same structure) increases with
  increasing mass

13
Surface Tension

• How many drops on a penny?
• Uneven forces at surface
• Acts like pond scum
• Definition energy needed to increase the
  surface area of a liquid by a certain amount
• Water is high why?
• Called cohesive force together
• Water moving up a stem adhesive force
• Capillary acion rise up a thin tube
• Meniscus!

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Phase Changes

• Solid to Liquid is called Heat of Fusion ?Hfus
• For water, 6 kJ/mol
• Liquid to Gas is called Heat of Vaporization
• ?Hvap
• For water, 40.7 kJ/mol
• ?Hsub is sum of each

15
Heating Curve

• Try a problem
• Remember - flat during phase change, temperature
  change when heating a single phase
• Cooling is opposite

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Supercooling

• Happens with some liquids - remove heat and it
  doesnt freeze when it should
• Very unstable
• May happen during hibernation

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Critical Temperature

• Highest temperature at which a liquid can form
  from a gas when pressure is applied.
• Above this, the substance is called a
  supercritical fluid.
• Gas just becomes more compressed.
• Critical pressure - pressure at the critical
  temperature

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Vapor pressure

• Vapor pressure forms above any liquid if
  container is closed why?
• Equilibrium is reached
• This is vapor pressure
• Higher if forces holding liquid together are weak
  - called a volatile (fleeing) liquid

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Boiling Point

• Temperature at which the VP equals atmospheric
  pressure
• Normal BP - boiling point at 1 atm
• Everest? Autoclave?

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Phase Diagram

• Handout
• Look at lines
• Look at slope of AB
• Freeze-drying - library book example

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Water vs. CO2
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Structure of Solids

• Amorphous (rubber, plastics) - large or mixtures
  - no true structure
• Crystalline - highly ordered structure
• Crystalline solids have true melting points

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Unit Cell

• Repeating unit of a solid
• 7 types (6-sided parallelograms)
• Ni, Na, NaCl
• Array of points in the crystal lattice

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3 cubic unit cells
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Total Atoms for each unit cell
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Packing

• Spheres naturally pack hexagonally
• Animation

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Bonding

• Shown by x-ray diffraction
• Molecular - low MP
• If unit packs well, mp can be high
• Covalent Network Solid - very strong
• Many covalent bonds in 3-D
• Diamond, graphite, SiO2, SiC, BN
• Ionic - greater charge, greater MP
• Metallic solids - hexagonal close packed, mp
  varies