Course Grades Check for mistakes! (December 14, 2017) MCTC Principles of Chemistry 1 Hybrid Fall 2017 Instructor: Kirk Boraas Email Kirk Course Lectures D2L Lab Manual Online Office Hours: click On-Campus Office Hours  (S-3580):  Monday: 11AM - 2PM Wednesday: 4-5 PM Syllabus Click Here CHEM 1152 Website This is NOT an online course.      Campus visits are required (weekly labs and periodic exams). Withdrawing? Click HERE first. (NEW!!!)  Exam #2 Fall 2017 (NEW!!)   Exam #4 Fall 2017 Course Announcements   9/13   OR   9/14     7 - 9 PM  D2L Exam 1                 10/12    Exam 2  6-9 PM   Rm S-2500         11/8  OR 11/9        7 - 9 PM  D2L Exam 3                  12/8    Exam # 4  6-9 PM  Rm L-3000       12/13 Optional Cumulative Final Exam 6-9 PM Rm S-1400 PREVIOUS BLOG LINKS   Week 1 & 2            Week  3 & 4             Week 5 & 6          Week 7 & 8      ||      Week 9 & 10       Week 11 & 12         Week 13 & 14      Week 15 & 16

 December 6,  2017    Wednesday   Week 16 Chapter 11: Water: An extraordinary substance Study Questions 1. What's so special about water? 2. Non-polar molecules are .....    Polar molecules are ....     a. hydrophilic           hydrophilic     b. hydrophobic        hydrophobic     c. hydrophobic        hydrophilic     d. hydrophilic          hydrophobic 3.  The density of solid water  (a.k.a. ice) is much different from that of most       other materials.  What is this difference and how is it observed? 4. How does the spacing between water molecules for liquid water compare     to the molecular spacing for solid water (ice)? 5. Local large bodies of water help control the temperatures of areas around     them.  How is this possible? 6. Why do you feel cooler when sweat evaporates? Water-Liquid AwesomeWater-Liquid Awesome Click and drag below for answer 1. It forms hydrogen bonds.  It is the only substance known to exist in all 3 forms (solid, liquid and gas) naturally on the surface of the earth.     High cohesion.  Water is an "Amazing Solvent!" 2. c 3.  Unlike most substances, the density of solid water (a.k.a. ice) is LESS than the density of liquid water.  Consequently ice floats.  We observe this      when as ice cubes floating in a drink and ice forming on the top of a lake whilst the majority of the lake under the ice is still liquid water. 4. The spacing between molecules for ice is GREATER than the spacing between molecules for liquid water.      This makes the ice less dense than the liquid. 5. Water has a high heat capacity.  This means it can absorb a lot of heat energy without changing its temperature.  Thus, a lot of heat can be     present but if  absorbed by water, the temperature doesn't change much. 6. As water evaporates, hydrogen bonds are broken and this energy is removed.  Removing energy lowers the temperature of the skin.

 December 5,  2017    Tuesday   Week 16 Chapter 11: Phase Diagrams Phase Diagram BasicsPhase Diagram Basics Phase Diagrams: DetailsPhase Diagrams: Details Study Questions The following questions refer to the CO2 phase diagram at right. 1. Identify the physical states corresponding to i.  , ii.  , iii.   and iv. 2. CO2 is a ________ at room temperature and atmospheric pressure.      a. solid       b. liquid        c. gas       d. super critical fluid 3. What it the meaning of the "triple point" and what are the     temperature and pressure corresponding to CO2's triple point? 4.  A CO2 sample at a pressure of 80 atm and a temperature of      40oC exists  as a ...      a. gas       b. liquid         c. solid       d. supercritical fluid 5. A CO2 sample at a pressure of 60 atm and a temperature of      -20oC is warmed to a temperature of 40oC while keeping the       pressure constant.  What is the order of phase changes that occur? 6. Consider a CO2 sample at  -60oC .   If the pressure is gradually increased from 10 atm, what physical changes would you expect to see? 7. Consider a CO2 sample at  -50oC.    If the pressure is gradually increased from 10 atm, what physical changes would you expect to see? 8.  What is a supercritical fluid? Click and drag below for answers.. 1. i. Solid     ii. Liquid    iii. gas    iv. Supercritical fluid 2.  gas 3. The triple point is that set of conditions that CO2 can exist simultaneously as a solid, liquid and gas.  For CO2, this corresponds      to T = -56.57 oC and P = 5.11 atm. 4. Supercritical Fluid 5. solid -> liquid -> gas 6. No changes.  The CO2 begins as a solid under these conditions and increases in pressures when T= -60oC  always lie in the "solid" region     of the phase diagram 7. The liquid form of CO2 is what's observed when T = -56.57 oC and P = 5.11 atm.  However, increases in pressure lead to the formation of a     solid as the line separating solid and liquid is crossed. 8. A supercritical fluid exists at pressures and combinations where liquid and gas phases become indistinguishable.  See the following link     for more information on supercritical fluids....  Click here.

 December 4,  2017    Monday   Week 16 Chapter 11: Heating Curve for Water Phase Changes KB Phase Changes KB Heating Curves and Calculations KB Heating Curves and Calculations KB Study Questions 1. How does the temperature change when ice is melting? 2. Which of the following equations is used when calculating     heat the heat required to change liquid water's temperature from     10oC to 60 oC?      a. q =  m c Δ T              b. q =  n ΔHf            c. E = mc2 3. How many grams of steam can be produced if 125 grams of liquid     water is vaporized?      a. Zero grams      b. 50 grams       c. 100 grams       d. 125 grams 4.  Refer to the diagram at right:       a.  How much heat is required for  50.0 grams of water at 20.00oC to be             heated to 100.00oC.       b. How much heat is required for 50.0 grams of water at 100.00oC          to be completely vaporized into steam at 100.00oC?      c. How much heat is required to convert steam at 100.00oC into          steam at 120.00oC?      d. What is the total amount of heat required to convert 50 grams          of water at 20.00oC into steam at 120.00oC?  (...add them up!) 5.  How many joules (J) of energy are released when 6.80x103g of steam at 100.0°C are completely frozen to ice at 0.0°C: (Source) Click and drag below for answers 1. For phase changes like ice melting or water boiling, there is no change in temperature 2. Since there is a change in temperature, we need to use it in the calculation and only "a" has a delta T term 3. Mass is conserved and so you should have the same mass of steam as you had water:  "d"   125 grams 4. a.  q = m c Δ T = (50.0 g)   (4.184 J/goC)   (100.00oC - 20.00oC)          = 16736 Joules     b.  q = n ΔHv   = (50.0g/18.01g/mol)   (40.7 kJ/mol)                               = 112.99 kJ     c.  q = m c Δ T = (50.0 g)  (1.84 J/goC)   (120.00oC - 100.00oC)            = 1840. Joules     d.   (Remember to convert all units to either kJ or J before adding up)    131.6 kJ 5. 2.05 x 107 Joules

 December 1,  2017    Friday   Week 15 Chapter 11: Sublimation  and fusion Clausius-Clapeyron Equation Clausius-Clapeyron Equation Clausius-Clapeyron Equation: Boiling Point Determination Clausius-Clapeyron Equation: Boiling Point Determination Study Questions 1. The vapor pressure of water is 1.0 atm at 373 K, and the enthalpy of vaporization is 40.7 kJ mol-1.      Estimate the vapor pressure at temperature 363 and 383 K respectively. (Source) 2. The vapor pressures of ice at 268 K and 273 K are 2.965 and 4.560 torr respectively. Estimate the heat of sublimation of ice. (Source) 3. Calculate ΔHvap$\mathrm{\Delta }{H}_{}$vap for ethanol, given vapor pressure at 40 oC = 150 torr. The normal boiling point for ethanol is 78 oC. (Source) Click and drag below for answers 1. 1.409 atm 2. 52.37 kJ/mol 3. 39 kJ/mol

 November 30,  2017    Thursday   Week 15 Chapter 11: Vaporization Vapor Pressure and boiing Vaporization and Vapor Pressure Vaporization and Vapor Pressure Atmospheric Pressure and Boiling Atmospheric Pressure and Boiling Study Questions 1. Dynamic equilibrium exists in a closed container that is partially filled with water.  What does this mean?     a. liquid molecules become gas molecules at the same rate as gas molecules become liquid molecules.     b. liquid molecules become gas molecules at a rate slower than gas molecules become liquid molecules.     c. liquid molecules become gas molecules at a rate faster than gas molecules become liquid molecules. 2. What is "Vapor Pressure"      a. the pressure of the gas formed by a liquid evaporating in an open container.      b. the pressure of the gas formed by a liquid in a sealed container.      c. the pressure of a gas formed when liquid and gas are in dynamic equilibrium 3. What prevents many (but not all)  liquid molecules from entering the vapor phase?      a. adhesive forces between the liquid and the walls of the container.      b. molecules in the vapor phase that get in the way of escaping liquid molecules.      c. cohesive intermolecular forces that pull a molecule back into the liquid phase 4. How do stronger intermolecular forces affect vapor pressure?      a. Stronger intermolecular forces produce no change in a liquid's vapor pressure      b. Stronger intermolecular forces produce an increase in a liquid's vapor pressure.      c. Stronger intermolecular forces produce a decrease in a liquid's vapor pressure. 5. As a liquid is heated, how do the intermolecular forces change?     a. Intermolecular forces get stronger     b. Intermolecular forces get weaker.     c. Intermolecular forces don't change with temperature Click and drag below for answers 1. a    2. b and c    3. c     4. c    5. c   However, kinetic energies do increase making it possible to overpower intermolecular forces.

 November 29,  2017    Wednesday   Week 15 Chapter 11: Surface Tension, Viscosity and Capillary Action Surface Tension Surface Tension Viscosity Viscosity Study Questions 1. What is cohesion? 2. What is adhesion? 3. Which of the following best describes how surface tension develops?      a. Intermolecular forces attract all molecules equally.      b.  Weak intermolecular forces permit some molecules to escape            the liquid phase.      c. Unbalanced intermolecular forces pull surface molecules           downward at the surface of liquids.      d. Intermolecular forces between different materials producing          an adhesive affect. 4. How would you expect the viscosity of non-polar CCl4 to compare        the viscosity of H2O? Capillary Action and the Meniscus Capillary Action and the Meniscus Questions 5, 6, and 7 refer to the figure at right.that demonstrates capillary action differences                 between water and mercury. 5. Adhesive forces exist between ...      a. adjacent liquid molecules      b. liquid molecules and the sides of the capillary tube      c. liquid molecules inside the tube and outside the tube      d. liquid molecules and gravity 6. For which material are the cohesive forces stronger than the adhesive forces?      a. Water     b. Mercury    c. in either case adhesive and cohesive forces are equal. 7. Why doesn't the water in the capillary climb completely out of the tube?     a. the capillary tube is sealed at the top     b. adhesive forces aren't capable of supporting a taller column and its additional weight     c. cohesive forces are greater for taller water columns and prevent the water from going further up the capillary column     d. water will climb out of the tube if you wait long enough. Click and drag below for answers 1. Attraction a material has for itself. 2. Attraction a material has for different materials 3. c 4. CCl4 is non polar and experiences very weak London Dispersive intermolecular forces. Water experiences much stronger hydrogen bonding.     The more strongly connected water molecules you would expect have higher viscosity and be "thicker" when pouring than CCl4. 5. b 6. b 7. b.

 November 28,  2017    Tuesday   Week 15 Chapter 11: Intermolecular Forces Intermolecular Forces Intermolecular Forces Hydrogen Bonding and Common Mistakes Hydrogen Bonding and Common Mistakes Study Questions 1. What is dipole dipole attraction? (Source) 2. What is hydrogen bonding? (Source) 3. How does the strength of the dipole-dipole intermolecular force compare to that of typical covalent bonds? (Source) 4. How do London Dispersion forces work? (Source) 5. How do intermolecular forces affect each of the following? (Source)    a. When intermolecular forces increase  the boiling point temperature _______.    b. When intermolecular forces decreases the melting point temperature _______.    c. When intermolecular forces increase the vapor pressure _______. 6. Which of the following is the weakest attractive force? (Source)     a. polar covalent bond         b. ionic bond        c. hydrogen bond        d. dipole-dipole force 7. What is the strongest intermolecular force present for each of the following compounds?  (Drawing structures will help identify the type of      intermolecular force). (Source)     a. H2O      b. CCl4     c. NH3     d. CO2      e. PCl3      f. N2       g. C2H6      h. CH3COCH3     i.  CH3OH      j. BH3    Click and drag below for answers 1. For molecules that have positive and negative ends (i.e. they're polar), they are attracted to one another (unlike charges attract) 2. Hydrogen bonding is the strongest type of intermolecular force. It occurs when the highly electronegative elements N, O, F bond with H.     These attractions are stronger because as N, O, and F pull the bonding electrons closer to themselves all that is remaining on the H side of     the compound  is a “naked” proton. This causes the N, O, or F side to be greatly negative and the H side to be extremely positive. 3. All intermolecular forces are significantly weaker than intramolecular forces; this includes dipole-dipole (an intermolecular force) and      covalent bond (an intramolecular force). 4. London dispersive forces are the result of momentary shifts in electrons within a molecule that leave it with a brief positive and negative end.     These brief charges are enough for molecules to have a net attraction for eachother.  All molecules experience London Dispersive Forces although      there may also be stronger intermolecular forces at work. 5. a. When intermolecular forces increase the boiling point increases.     b. When intermolecular forces decrease the melting point temperature decreases.     c. When intermolecular forces  increase the vapor pressure decreases. 6. d 7. a. Hydrogen bond    b. LDF    c. Hydrogen bond   d. LDF   e. dipole-dipole    f. LDF     g. LDF    h. dipole-dipole     i. Hydrogen bond    j. LDF

 November 27,   2017    Monday   Week 15 Chapter 11: Solids Gases and liquids: A Molecular Comparison Study Questions 1. What is the name given to forces that exist between molecules?      a. Intermolecular forces      b. Intramolecular forces      c. Covalent bonding      d. Ionic bonding 2. In which physical state are the atomic or molecules moving with the     greatest kinetic energy?      a. Solid       b. Liquid      c. Gas 3.  Atoms and molecules in the solid state are unable to move at all. (T/F) 4.  As the temperature of a solid increases, the atoms or molecules      break free from their locations within the crystalline lattice      and begin to move around.  What is this called?      a. vaporization    b. deposition  c. freezing   d. melting    e. condensation Solids, Liquids and Gases: A Molecular Comparison Solids, Liquids and Gases: A Molecular Comparison 5. Which of the following is exothermic and releases heat into the surroundings?     a.  freezing       b.  boiling        c. subliming         d. melting          e. condensing     f. depositing (as in "deposition"  ...  gas ->  solid) 6.  Which of the following is true for crystalline solids?       a.  Crystalline solids only exist at very low temperatures       b.   Crystalline solids have their particles arranged randomly       c.   Crystalline solids have highly ordered structures       d.   Crystalline solids are usually very soft       e.   Crystalline solids exist only at high temperatures 7. In liquids, which of the following best describes the intermolecular forces?       a.  strong enough to hold molecules relatively close together but not strong enough to keep molecules from  moving past each other       b.  strong enough to keep molecules from moving past each other       c.  very weak compared with kinetic energies of the molecules       d.  not strong enough to hold molecules relatively close together       e. strong enough to keep the molecules confined to vibrating about their fixed lattice points 8.  How do the intermolecular forces change as a material changes from solid to liquid to gas?       a. Intermolecular forces increase      b. Intermolecular forces decrease     c. Intermolecular forces don't change Click and drag below for answers p. 1. a     2. c     3. False (Vibrate)    4. d     5. a, e and f      6. c   7. a     8. c.  Intermolecuar forces don't change.  However, the kinetic energies of atoms/molecules increase with temperature and this affects how well the          intermolecular forces are able to keep molecules close together.