 Why does water melt at 0 degrees Celsius and vaporize at 100 degrees Celsius?  e_viewer.php?mid=120 e_viewer.php?mid=120

INTRAmolecular Forces: Bonds WITHIN molecules (Ionic, Covalent, Metallic) INTERmolecular Forces: Bonds BETWEEN molecules Give molecules its properties (i.e. melting/boiling points) Intramolecular forces are stronger than Intermolecular forces Strongest intermolecular forces are less than 5% as strong as intramolecular forces.

 London Dispersion Forces (or just Dispersion Forces)  Dipole-Dipole Forces  Hydrogen Bonding  Ion-Dipole Forces (Mixtures)

 Exist between ALL particles  TEMPORARY shifts of electrons (temporary dipole) within an e- cloud  Electrons go back to evenly distributed when atoms/molecules move away from each other  The higher the mass, the more e- mobility, the stronger the LD force.

 Molecules that are POLAR (uneven distribution of charge) have PERMANENT shift of electrons (i.e. they have dipoles)  Some parts are slightly + and some are slightly –  The larger the dipole, the stronger the force  Opposite charges will align themselves.

 Special type of dipole-dipole  STRONGEST intermolecular force (relatively speaking)  Occurs between molecules containing:  A hydrogen atom bonded to  A Nitrogen, Oxygen, or Fluorine atom  Ex: H 2 O NH 3 HF

Animation of H-bond formation, from Robert Wyatt, Western Kentucky Univ.

 Attractions between ions and polar molecules  Occur in MIXTURES containing ions and polar molecules  Ex: Aqueous solution of Sodium Chloride (i.e. salt water)  The stronger the ion, the stronger the force

 Strong intermolecular forces:  A lot of energy needed to vaporize  High boiling points  Relatively favor solid state  Weak intermolecular forces  Little energy needed to vaporize  Low boiling points  Relatively favor liquid or gas state

 Identify what intermolecular force(s) are present in the following: 1) O 2 5) CCl 4 2) HF6) H 2 3) NH 3 7) HBr 4) NaCl8) H 2 O 2

 Forces on surface are stronger than forces inside  Strong cohesive forces between surface molecules need to compensate for lack of forces from above Cohesion is the attraction between like molecules If intermolecular forces are strong, surface tension is higher

 Measure of resistance to flow  Molasses = higher viscosity  Acetone = lower viscosity  Viscosity decreases with increasing temperature  Breaks/weakens intermolecular forces  If forces are high, viscosity is high

 Amorphous Solids: Solids lacking regular geometric structure  Crystalline Solids: Solids containing rigid structure and order  Atoms, molecules, or ions occupy specific positions

 Structure of crystalline solids depend on  Size  Nature of molecule

 Ionic  Covalent  Molecular  Metallic

 Composed of charged spheres (cations & anions)  Anions typically larger than cations  Do not conduct electricity in solid state, but do in molten (melted state)  Ions can move around when melted

 Atoms are all held together entirely by covalent bonds  Ex: diamond and graphite (see next slide)  Because of the strong covalent bonds, covalent crystals experience high melting points and boiling points  Can be treated as one big molecule

 Molecules held together by intermolecular forces  Most molecular crystals have lower melting points and break easier than ionic/covalent  *Remember: INTERmolecular weaker than INTRAmolecular

 Same metal atom  Usually very dense  Electrons are delocalized (“sea of electrons”)  Cohesion makes metals strong Good conductors of heat & electricity

 A change in the phase (or state) of matter Endothermic— Requires heat Melting (solid to liquid) *melting point Vaporization (liquid to gas) *boiling point Sublimation (solid to gas) Exothermic— Releases heat Freezing (liquid to solid) *freezing point Condenstation (gas to liquid) Deposition (gas to solid)

 The Boiling Point is the temperature where vapor pressure equals external pressure  The Normal Boiling Point is the temperature where vapor pressure is 1 atm (atmosphere)  STP = Standard Temperature and Pressure = 1.00 atm and 0.00°C

 ations/waterphases/status_water.htm ations/waterphases/status_water.htm

 nges/HeatingCurve.html nges/HeatingCurve.html

TIME

 Heat of fusion: how much energy is needed to melt one mole of a substance  Heat of vaporization: how much energy is needed to vaporize one mole of a substance  Larger than heat of fusion  Reverse the sign if wanting to know how much energy needs to be removed to freeze and condense

 Which phase changes are exothermic? Endothermic?  If you have 15.0 g of ice, how much energy is require to melt it? (water ΔH fus = 6.01 kJ/mol)  If 2.45 kJ of energy are released when methane (CH4) is condensed, how many grams of methane do you have? (methane ΔH vap = 8.17 kJ/mol)

 States of matter depend on:  Temperature  Pressure  Phase Diagrams show what state of matter a substance will be under certain conditions of pressure and temperature  Different for every compound.

 Triple Point: Point where all three states of matter coexist  Critical Point: Point where liquid and gas states become indistinguishable.

 When carbon is at a low temperature but a high pressure, what form is it in?  Diamond  What would happen if you had carbon at a pressure of 10 3 and raised the temp from 2000 to 4000?  It would melt

 Which state of matter tends to exist under:  High Temp, High Pressure ▪ Liquid  High Temp, Low Pressure ▪ Gas  Low Temp, High Pressure ▪ Solid

 What would happen to a substance if its pressure was approx. 3.0 atm and its temp. increased from -78 to 29?  It would go from solid to liquid (melt)

 What would happen to a substance if its temp was -78 and its pressure changed from 2.0 atm to 0.5 atm?  It would go from solid to gas (sublime)

 What state of matter is water in at 50.0 degrees C and 1 atm?  Liquid  What is the critical temp and pressure for water?  C & atm