Phases _ Energy Guided Inquiry

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					Chemistry – Unit 7: Energy
Phases & Energy Guided Inquiry
The Kinetic-Molecular Theory of Matter
The kinetic molecular theory is based on the idea that particles of matter are always in motion. In
the late 19th century scientists developed this theory to describe the behavior of atoms and
molecules that make up matter. The theory can be used to explain the solid, liquid and gas phases
in terms of the energy of the molecules and the attractive intermolecular (aka attractive) forces
that act between them.
1. Why is the KMT important?




Important Relationships among Temperature, Mass, Average Velocity & Average Kinetic Energy
We learned earlier in this unit that kinetic energy increases as
temperature increases. Molecules at the same temperature
have the same average kinetic energy. This does not mean           Number of Molecules
that in a group of molecules at the same temperature every
molecule has exactly the same kinetic energy. The individual
molecules at a given temperature have a range of kinetic
energies. The chart on the left graphs the distribution of
kinetic energies for each group of molecules. Each group has
an average kinetic energy, which is shown on the curves as
the symbol μ. The curve with the higher “hump” represents
                                                                                         Kinetic Energy
the molecules that are at a lower temperature than the molecules represented by the second curve. The
group of molecules at the higher temperature has a higher average kinetic energy.

Kinetic energy is the energy of motion. It is represented by the equation:



where, KE = kinetic energy, m = mass and v = velocity. This equation tells us that a molecule’s
kinetic energy depends on its mass and its velocity.
These relationships among temperature, mass, average kinetic energy and average velocity are
important in understanding and explaining the different properties of gases, liquids and solids.
Look at the kinetic energy equation and the kinetic energy chart to answer the following
questions:
2. There are two beakers of water. The water temperature in beaker A is -5 °C and the water
   temperature in beaker B is 20 °C. Which beaker has the water with the higher average kinetic
   energy? Explain your answer.




3. Will every molecule in the beaker with the higher average kinetic energy have more kinetic
   energy than every molecule in the beaker with the lower average kinetic energy? Explain your
   answer.




4. Beaker C contains ethanol, which has a molar mass of 46 grams, and beaker D contains water,
   which has a molar mass of 18 grams? The temperature of both liquids is 30 °C.
  a. Which beaker contains the liquid with the higher average kinetic energy? Explain your
     answer.




  b. Which beaker contains the liquid with the higher average molecular velocity? Explain your
     answer.




Phases: Kinetic Energy vs. Intermolecular Forces
A phase is any part of a system that has uniform composition and properties. There are 4 phases:
solid, liquid, gas and plasma. In this lesson we are going to investigate the first three phases, but
not the plasma phase.
Plasma Phase – This is the highest energy phase. There is so much energy in this phase that the
negatively charged electrons have been stripped away from their positively charged nuclei. On
earth, naturally occurring examples of the plasma phase are flames, lightning and auroras. The
most common example of artificial plasma is the electrified gas in a fluorescent light bulb.
Phase (aka States of Matter) Simulation
Go to the University of Colorado’s PhET site and run the “States of Matter – Basics” simulation:
http://phet.colorado.edu/en/simulation/states-of-matter-basics .
Click on the “Solid, Liquid, Gas” tab and explore what happens as you switch among the phases for
each type of molecule. Pay particular attention to the temperature and the motion of the
molecules.
5. Complete the table based on your observations.
                           Neon                Argon               Oxygen               Water
Solid

Temperature

Speed
(fast, med, slow)
Motion
(random or fixed)

Definite Shape?

Definite Volume?

Liquid

Temperature

Speed
(fast, med, slow)
Motion
(random or fixed)

Definite Shape?

Definite Volume?

Gas

Temperature

Speed
(fast, med, slow)
Motion
(random or fixed)

Definite Shape?

Definite Volume?

6. Pick a substance and set the temperature to 0 K:
  a. What happens to the molecules?




  b. What phase are the molecules in?


7. For the same substance add heat to raise the temperature to 400 K:
  a. What happens to the molecules?




  b. What phase(s) are the molecules in?


  c. How can you tell if the molecules change from a solid to a liquid?




  d. How can you tell if the molecules change from a liquid to a gas?




8. Pick a different substance, click the gas button and describe the motion of the molecules.




9. For the same substance cool the temperature to 0 K:
  a. What happens to the molecules?




  b. What phase(s) are the molecules in?


10. Which phase has the highest temperature?


11. Which phase has the lowest temperature?


12. Rank the phases in order from lowest to highest average kinetic energy.
Intermolecular (aka Attractive) Forces
We learned in an earlier unit that intermolecular forces such as London dispersion forces, dipole-
dipole attractions, hydrogen bonding and ionic bonding attract molecules to each other. The
strength of these intermolecular attractions compared to the kinetic energy of the molecules
determines the phase of a substance.
13. Based on your observations of molecular motion during the simulation:
   a. In which phase are the molecules’ intermolecular forces strong enough to overcome its
      average kinetic energy?




   b. In which phase are the molecules’ intermolecular forces and average kinetic energy about
      the same?




   c. In which phase does the molecules’ average kinetic energy overcome its intermolecular
      forces?




14. Use the internet and your understanding of intermolecular forces to complete the table:

                           Neon                Argon              Oxygen                Water
Intermolecular
Forces

Melting Point

Boiling Point

15. Write a statement that describes the relationship between melting and boiling points of a
    substance and the strength of its intermolecular forces.
Energy & Phase Changes - There are 6 types of phase changes:




16. Write a definition and give an example of each phase change.
  a. Vaporization (aka boiling)




  b. Condensation




  c. Melting




  d. Freezing




  e. Sublimation




  f. Deposition
Phase changes (changes of state) either require energy or give off energy when they occur. When
a molecule goes from a lower energy phase to higher energy phase energy is required.
Conversely, when a molecule goes from a higher energy phase to lower energy phase energy is
released. When energy is required during a phase change the potential energy of the substance is
increased. When energy is released during a phase change the potential energy of the substance
is reduced. The absorption or release of energy must occur during a phase change because of the
conservation of energy law.
The amount of energy required or released is the enthalpy of the phase change:
                    Energy required to change liquid to gas = enthalpy of vaporization = ΔHv
                    Energy required to change solid to liquid = enthalpy of fusion = ΔHf
                    Energy required to change solid to gas = enthalpy of fusion + enthalpy of vaporization
                     = ΔHf + ΔHv

17. Look up the enthalpy of fusion and enthalpy of vaporization valued for these substances.
    Include the proper units

                                                       Enthalpy of Fusion, ΔHf         Enthalpy of Vaporization, ΔHv

                          Neon

                          Oxygen

                          Water


18. Rank the phase changes based on the amount of energy they require or release. Write a phase
    change equation that shows the phase change and quantifies the amount of energy required
    or released in terms of ΔHv and ΔHf. The melting phase change equation is filled in as an
    example.
    Increasing Energy →




                                          Melting

                                   Solid + ΔHf → Liquid
Energy, Phase Changes & Temperature
When energy is added to a substance it will either raise the temperature of the substance
(increase kinetic energy) or change the substance to a higher energy phase (increasing potential
energy). Both effects won’t happen at the same time.
19. For each situation below indicate what happens to temperature (T ↑, T ↓ or T ↔ ) kinetic
    energy (KE ↑, KE ↓ or KE ↔ ), and potential energy (PE ↑, PE ↓ or PE ↔ ) when heat is added or
    removed.

 Phase(s) of Substance         Heating (adding energy)            Cooling (removing energy)

 Solid


 Solid + Liquid


 Liquid


 Liquid + Gas


 Gas


20. Compare and contrast vaporization and condensation.




21. How do the boiling point and condensation point of a substance compare to each other?


22. Compare and contrast freezing and melting.
23. How do the freezing point and melting point of a substance compare to each other?




24. Compare and contrast sublimation and deposition.




25. How do the sublimation point and deposition point of a substance compare to each other?



Phase Diagrams: The Effects of Temperature & Pressure on Phase Changes
Temperature - We have already discussed how increasing temperature will change a substance to
a higher energy phase and decreasing temperature will change it to a lower energy phase.
Pressure – The kinetic molecular theory predicts that increasing pressure will hold molecules
closer together and it will take more kinetic energy to get the molecules to change to the next
higher energy phase. Lowering pressure has the reverse effect.
26. Would you expect the boiling point of water in Denver, CO (elevation ~1,600 meters) to be the
    same, higher or lower than the boiling point of water in Annville, PA (elevation ~ 150 meters)?
    Why?




27. Would you expect the boiling point of water at the bottom of the ocean to be the same, higher
    or lower than the boiling point of water at the surface of the ocean? Why?
Phase Diagrams – are used to show the phase of a substance for any given combination of
temperature and pressure. Each substance has its own unique phase diagram.




                  Phase Diagram for Water                         Phase Diagram for Carbon Dioxide

The melting/freezing and boiling points of any substance depend on the pressure. When we say
normal melting point or normal boiling point we mean the temperature at 1 atmosphere of
pressure. The triple point is the temperature and pressure where all three phases of a substance
are in equilibrium with each other. The critical point is the temperature above which a substance
cannot exist as a liquid no matter how high the pressure is. Notice the solid-liquid line for water
has a negative slope and the solid-liquid line for carbon dioxide has a positive slope. The solid-
liquid line has a positive slope for most substances. Water is the exception! The negative slope for
water means that liquid water is denser than ice and that increasing pressure lowers the melting
point instead of raising the melting point.
28. Water is in what phase at 100    and 1.2 atmospheres?



29. Carbon dioxide is in what phase at 300 K and 10 atmospheres?
Energy Calculations for Phase Changes
We can use the enthalpies of vaporization (ΔHv) and fusion (ΔHf) to calculate the amount of energy
that is required or released to change the phase of a known amount of water; or we can calculate
the amount of water that changed phases if we know how much energy was absorbed or released.
The phase change energy equations are:
            for melting or freezing

            for vaporization or condensation
            for sublimation or deposition


For water at 1 atm     ΔHf = 333.4 J/g         ΔHv = 2266 J/g
Sample Problem #1 – How much energy is required to completely vaporize 100.0 g of water?




Sample Problem #2 – It took 100.0 kJ of energy to melt a piece of ice. What was its mass?




30. How much energy is given off when 250.0 g of water is frozen?




31. How many grams of water can 479.0 kJ of energy vaporize?
Combined Energy Calculations
When a substance is heated or cooled it may change temperature as well as phase. We can use
our combined energy equation to calculate the amount of energy required to change the
temperature and phase of a substance.


                                         temp change + phase change

The specific heat of a substance is different for different phases. Be sure you are using the correct
specific heat. For example, the specific heats for ice, water and water vapor are different.
If the temperature change is going through several phase changes we must expand the equation:


      solid temp change       solid-liquid       temp change            liquid-gas    gas temp change
                             phase change                             phase change



Sample Problem
How much energy is required or released to
heat 50.0 g of ice from -20.0°C to 120.0°C
steam?
Cice = 2.03 J/°C•g
Cwater = 4.184 J/°C•g
Csteam = 1.89 J/°C•g
ΔHf = 333.4 J/g
ΔHv = 2266 J/g



  Energy to heat ice             50.0 g x 2.03 J/°C•g x (0°C – (-20.0°C))      = 2030 J
+ energy to melt ice           + 50.0 g x 333.4 J/g                            = 16700 J
+ energy to heat water         + 50.0 g x 4.184 J/°C•g x (100.0°C – 0°C)       = 20900 J
+ energy to vaporize water     + 50.0 g x 2266 J/g                             = 113000 J
+ energy to heat steam         + 50.0 g x 4.184 J/°C•g x (120.0°C – 100°C)     = 4180 J
    Total Energy                                                                157000 J required

32. How much energy is required or released to cool 25.0 g of steam from 105°C to ice at -5.0°C?

				
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