Chapter 10 Review (sections 1-6 and 11)

  1. Polar Bonds vs. Polar Molecules (section 10.1)
    1. A molecule is polar if the polar bonds do not "cancel out" when you add them up (vector addition). 
    2. CF4 is tetrahedral in electron geometry and tetrahedral in molecular shape. All four bonds are C-F bonds and are very polar YET the bond dipole arrows all cancel out.  The overall molecule is nonpolar despite the polar bonds.
    3. NH3 is tetrahedral in electron geometry and trigonal pyramid in molecular shape.All three N-H bonds are polar AND the bond dipole arrows do NOT cancel out.  The overall molecule is polar and has a dipole moment (a north and south pole)
    4. Check for yourself that CO is polar and CO2 is nonpolar despite having polar bonds.
    5. Look at example 10.2 and try problems 10.2 and 10.4
    6. Basically polar molecules are those who have polar bonds arranged asymmetrically.
    7. Great animation and example problems - make sure you click next (IMPORTANT)
  2. Intermolecular Forces (IMF) (section 10.2)
    1. Atoms are held together in a molecule by covalent bonds.  But what holds molecules with each other?  IMF!  IMF are all electrostatic attractions (+ attracted to -)  Check out this tutorial - scroll down to Intermolecular Forces Tutorial
    2. Check out Figure 10.1 in the text.  The IMF of gases are <<< liquids < solids.  Solids are held together very well and it takes lots of energy to pull them apart. Liquids are held closely together, just not in a rigid pattern, but the forces are still significant. Gases are already apart - hardly any forces exist between the molecules and atoms of a gas. 
    3. Remember than ionic and covalent bonds are far stronger than IMF.
    4. The IMF from strongest to weakest are: (ALL IMPORTANT)
      1. Ion-dipole force:  This occurs when ions and polar molecules are together.  Cations (+) attract the d- end of the polar molecules and anions (-) attract the d + end of the polar molecules.  This is why some ionic solids dissolve in water.  The sum of the ion-dipole forces of the ions surrounded by water is greater than the original ionic bond. So the ionic bond breaks and the water molecules surround the ion.  Why does salted water boil at a higher temperature than plain water???  Think about it.  Figure 10.3
      2. H bonding force.  H bonds are not really bonds but a force. This is a special kind of dipole-dipole attraction.  When an H is covalently bonded to an N, O or F it can be attracted to a lone pair on another molecule.  This is because the H-F, H-N, and H-O bonds are very polar with the H being d + and the N, O, F being d-.  So the H is attracted to the negative lone pair on another molecule. Figure 10.7.  Thank goodness for the H bond force in water.  Because of the H bond force water is held together stronger than expected and it has a higher mp and bp than expected.  In general the more mass the higher the mp and bp. Look at Table 10.4  
        1. bp of H2O is actually higher than the more massive H2S which is abnormal
        2. bp of NH3 is actually higher than the more massive PH3 which is abnormal
        3. bp of CH4 is lower than SiH4 which is normal (methane can’t H bond – you should know why!)
        4. bp of HF is actually higher than the more massive HCL which is abnormal
        5. Water, ammonia and HF have abnormally high mp and bp despite small mass because they have H bond forces!!!
      3. Dipole-dipole force.  When polar molecules are attracted to each other.  Consider H-Cl.  The Cl is more EN than H so it pulls the 2 electrons in the polar covalent bond closer and is d- thus the H is d +.  When another H-Cl gets close they arrange themselves so the Cl of the first molecule is close to the H of the second. Figure 10.4.  As polarity increases, the IMF increase and the mp and bp increase.  This is because the stronger the IMF the harder it is to pull these molecules apart so a higher temperature is needed for melting and boiling.
      4. London force.  This force is in nonpolar molecules and neutral atoms.  It is weak and temporary.  Consider Xe atoms which are big and fat.  Let’s say the electrons for a moment are mostly on one side – then that side would be temporarily d- and the side with less electrons would be d+.  Now an atom next to this one would have it’s electrons attracted to the d+ side so would also become temporarily polar – we say it has an induced dipole.  Then it could induce a dipole in the next atom and so forth.  Figure 10.5.  All atoms and molecules have London forces, but the stronger forces overshadow them.  These induced dipoles are short lived.  London forces are best in big fat atoms and molecules with lot of electrons. Polarizability = how easy it is to induce a dipole.  Polarizability thus increases the more massive an atom.  Consider the table below.  The halogens all have only London forces since they are nonpolar.  Thus as the polarizability increases (mass increases) the London forces increase and the bp increases. Again this is because the more London forces, the harder it is to pull them apart.  Fluorine and chlorine with the least London forces are gases, then iodine with the most force is actually a solid. One last thing to mention – long chain molecules can wrap around each other thus making them harder to separate than rounder branched chains. Figure 10.6

 

F2

Cl2

Br2

I2

State

gas

gas

liquid

solid

bp (oC)

-188

-35

59

185
    1. Overall London forces < dipole-dipole < H bond force < ion-dipole
    2. Example 10.3 and problems 10.5 and 10.6
    3. Put in increasing bp order:  HBr, Ar, CH3OH, He    Answer is He < Ar < HBr < CH3OH
  1. Liquids (section 10.3)
    1. Viscosity = thickness of a liquid (oil is thick).  As IMF increases viscosity increases because the molecules will hold together stronger making the liquid thicker.
    2. Surface tension = how the molecules hold each other at the top of a liquid.  As IMF increases the surface tension also increases because the molecules will hold together stronger making the liquid surface tension higher.
    3.  Figure 10.8 and Table 10.6
  2. Phase Changes (section 10.4)
    1. There are six phase changes or changes in states of matter
      1. s g  l = melting
      2. l g  g = boiling, evaporation, vaporization
      3. s  g g = sublimation
      4. g g  s   = deposition
      5. g g  l  = condensation
      6. l  g s = freezing
    2. As we go from solid to liquid to gas, we are breaking the IMF that hold the molecules together.  Because this is just a physical change, the molecule itself is not changing - ie the BONDS holding the molecule together are staying intact. 
    3. As we add heat to melt or boil something the enthalpy is increasing, DH is + and feels cold as the heat is going into the substance.  As we remove heat to freeze or condense something the enthalpy is decreasing, DH is - and the heat is exiting. A freezer sucks the heat out of something to cool it down. Figure 10.9 (IMPORTANT)  Don't worry about DG or DS at this time. 
    4. Heating Curves (IMPORTANT) are graphs that show heat on the x axis and Temperature on the y axis. Watch this movie, go to activities and click on Changes of State Movie. Refer to the curve at the right. As heat is added a solid will warm up (A) until it melts, then the T stays constant while all the solid turns to liquid (B), finally the liquid heats up (C) until it boils, then the T stays constant while all the liquid turns to gas (D), then finally the gas heats up (E).
      1. Why does the T stay constant during a phase change?  Well all the energy is going into breaking or forming the IMF instead of kinetic energy which affects Temperature.  For example if melting, the energy is going into breaking apart the solid so a liquid can form. 
      2. At (B) solid is in equilibrium with liquid
      3. At (D) liquid is in equilibrium with gas
      4. DHfus is the heat of fusion = energy to convert from solid to liquid
      5. DHvap is the heat of vaporization = energy to convert from liquid to gas
      6. DHfus is always less than DHvap because it takes more energy to completely break the IMF to turn a liquid into a gas. 
  3. Evaporation, Vapor Pressure (VP) and Boiling Point (bp) (section 10.5)
    1. Boiling = when liquid turns to gas due to heat being applied.  Boiling is a phase change that occurs at the boiling point, which is a temperature.
    2. Evaporation is when liquid turns to gas slowly at a temperature less than the bp.  When a high kinetic energy molecule at the surface of a liquid gets bumped by other molecules, it can pop out of the liquid and join the atmosphere then float away.  Evaporation thus proceeds slowly.  When the higher KE molecules leave, the lower KE molecules are left behind so it feels cool.  This is why when you get out of the pool or shower you feel cold at first - all the evaporation from your skin. 
    3. Vapor Pressure = the pressure in a closed container due to the liquid evaporating and setting up equilibrium with its vapor. Watch this movie. Consider an empty container - put a liquid inside - close the lid. What is above the liquid?  Well air and pressure due to the air Pair. Now let it sit. After a while some of the liquid has evaporated and joined the air so the pressure is greater than it starter.  The additional pressure is the VP.  Now the total pressure in the container is Pair + VP.  Equilibrium is reached:  everytime a liquid molecule joins the vapor somewhere else a vapor molecule joins back with the liquid. (l)  D  (g)  equilibrium. (IMPORTANT) 
    4. Compare all three with pictures (IMPORTANT)
    5. As IMF increase, the evaporation rate will decrease because it will be harder to pull them apart from the liquid. As IMF increase the VP will decrease for the same reason, less will go into the vapor phase if they are held together with stronger force.  Finally as IMF increase the boiling point will increase because it will take higher energy at higher T to pull them apart.
    6. As T increases, evaporation rate will increase, VP will increase because with more KE the liquid molecules are more likely to join the vapor.  However bp stays constant. 
    7. Normal bp is at 1.0 atm (atmosphere) of pressure which is like sea level pressure - normal Earth surface pressure.  On super high mountains the pressure is less than 1.0 atm, and the bp will decrease.  With less atmosphere pushing down on liquids, they are more likely to pop out into the air. It totally sucks to cook up high in the mountains because water will not boil as hot as normal. Instant coffee is disgusting.
  4. Solids (section 10.6)
    1. Table 10.9 has the first four summarized.  There are a total of 5 types of solids:  (IMPORTANT)
      1. Ionic - held together by ionic bonds in all directions (sometimes called ion-ion forces), high mp and bp, brittle.  Example: NaCl and other salts.
      2. Molecular - held together by IMF, low mp and bp, insulators.  Examples:  ice, dry ice, solid iodine
      3. Covalent Network - held together by covalent bonds in all directions, high mp and bp, super hard, only molecule you can see with your bare eyes, they are one giant molecule.  Example:  diamond, ruby
      4. Metallic - held together by metallic bonds, medium to high mp and bp, conductors, malleable, ductile.  Example:  gold, copper, sodium, zinc
      5. Amorphous, polymers - held together by covalent bonds, high mp but usually decompose instead.  Example:  plastics, rubber
  5. Phase Diagrams (section 10.11)  (IMPORTANT)
    1. These are plots of P vs T.  At high pressures and cold T solids exist.  At low pressures and hot T gases exist.  Liquids are in between.  Now look at these points. Check out this tutorial - scroll down to Phase Diagrams Tutorial
      1. This is the triple point.  All three phases of matter can exist here in equilibrium. This combination of P and T is the only way all three states can exist at once and is unique for each chemical.
      2. This point is on the boundary between solid and liquid and is at 1 atm so must be the normal melting point.
      3. This point is on the boundary between gas and liquid and is at 1 atm so must be the normal boiling point.
      4. This is the critical point where no matter how hard you smush (increase pressure) you can't make a liquid at this temperature. Everything past point 4 is a super critical fluid.
    2. Note the segments between the phases are where phase changes occur.  There is the melting/freezing boundary of solid and liquid, there is the boiling/condensing boundary of liquid and gas, and there is the subliming/depositing boundary of solid and gas. 
    3. The phase diagram for CO2 looks like the one above except the 1 atm line is below the triple point.  What does that mean?  Well than at normal pressure here on Earth, the solid goes straight to the gas (ie it sublimes) and bypasses the liquid states completely.  You would have to increase pressure to above the triple point in order to see liquid CO2.
    4. Water's phase diagram is weird.  You should know exactly what is weird and what that means for water. Go to activities and click on Phase Diagram of Water Movie.
    5. Try Problems 10.17, 10.18, 10.19.

 

End of chapter 10 Review.  Study Hard. 

Recommended problems in the text:  26, 32, 34, 36, 40, 42, 44, 66, 82, 84, 86